Structural Biochemistry/Volume 1

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Introduction

Physics is the scientific study of physical phenomena and the interaction between matter and energy. Generally speaking, it is the examination and inquiry of the behavior of nature. As one of the oldest branches of academia, physics is intertwined with and helps explain the fundamental nature of the living and nonliving universe.

Thermodynamics

First law

The "first law" of thermodynamics is simply that energy is a conserved quantity (i.e. energy is neither created nor destroyed but changes from one form to another). Although there are many different, but equivalent statements of the first law, the most basic is:

$dU=dQ+dW$

dU = infinitesimal change in internal energy,

dQ = infinitesimal heat flow and

dW = infinitesimal work.

In words, the first law states that

The heat supplied is equal to the increase in internal energy of the system plus the work done by the system. Energy is conserved if heat is taken into account.

Entropy

Thermodynamics allows us to predict the initial and final states of a system. In other words, it’s an extremely useful tool in predicting the equilibrium position in chemical systems. The first law of thermodynamics states that energy is always conserved between the initial and final states. However, the first law does not provide information about the extent of a chemical reaction, or the equilibrium concentrations. The second law of thermodynamics introduces the concept of entropy S, a value that is a function of a system’s state. Entropy helps predict the equilibrium because the equilibrium concentrations of a system correspond to the maximum entropy of a system; entropy can act as a driving force of a reaction. Therefore, the second law of thermodynamics states that:

"Thermodynamic equilibrium in an isolated system is reached when the system's entropy is maximized"

Mathematically, entropy is expressed as

$dS\equiv {\frac {dq}{T}}$

for a reversible process in a closed system, where dq is heat energy and T is temperature. A reversible process is where the system is always very close to the equilibrium. Perturbations to the system must be small enough that the system and the surroundings can return to the initial state. The change in entropy, a more useful value, is defined as:

$\Delta S=S_{2}-S_{1}=\int _{1}^{2}{\frac {\mathrm {d} q_{rev}}{T}}$

An example of increasing entropy can be seen by watching ice melt in a closed container held at 25^oC. Since the temperature is constant, it can be seen that the total thermal energy of molecules in liquid water is greater than the thermal energy of molecules in ice. Therefore, ΔS is positive, which implies that entropy has increased in order to obtain equilibrium.

Entropy is conceptually more difficult to understand than other state functions such as temperature or energy. Furthermore, entropy is a macroscopic property; one molecule does not exhibit entropy. Consider a closed container with a partition down the center of the container where one side of the partition is composed of gas A and the other side of the partition is empty. If the partition is removed, there is a very small probability that gas A will stay on one side of the container. The most probable distribution of the A molecules is that gas A will be evenly distributed throughout the container. This distribution can be viewed as the equilibrium thermodynamic state, which happens to be the most probable state and the state with the most disorder.

Enthalpy

Enthalpy is a measurement of heat transfer done at a constant pressure and is often denoted as $H$. Mathematically, enthalpy is expressed as

$H=E+PV$

where E is the internal energy of the system, P is internal pressure, and V is the volume. PV accounts for the energy used in expansion work.

The change in enthalpy is then expressed as

$\Delta H=\Delta E+P\Delta V$

where pressure remains constant. $\Delta H$ is then, essentially, the internal energy corrected for work.

Reaction enthalpies are assigned to chemical reactions to denote the amount of energy that is transferred into or out of the system in exchange with the surroundings. A reaction that releases heat is defined as exothermic, where $\Delta H$ is negative. A reaction that requires an input of heat is defined as endothermic, where $\Delta H$ is positive.

Gibbs Free Energy

Free energy is a measurement of the tendency of processes to occur spontaneously. It depends upon 3 different quantities: change in entropy, change in enthalpy, and temperature. The Gibbs Free Energy is the enthalpy of the system subtracted by the temperature and the entropy of the system, G = H - TS. A positive change in free energy (such as an endergonic reaction) is thermodynamically unfavorable, whereas a negative change in free energy (such as an exergonic reaction) is thermodynamically favorable. In biological systems, many reactions that have positive free energy are coupled simultaneously to reactions that have negative free energy. Example is the synthesis of glutamine from the expense of hydrolysis of ATP.

Endergonic Reactions

Endergonic reactions are reactions that have a positive free Gibbs energy. These reactions are not thermodynamically favored and the substrates are not readily going to form products. Energy must be out into the system to drive these reactions.

Exergonic Reactions

Exergonic reactions are thermodynamically favored. These types of reactions have a negative free Gibbs energy an are readily able to form products. However, although these reactions are energetically favorable, it does not mean the reaction will occur at a reasonable rate. This is due to high activation energy barriers. To reduce these barriers, the introduction of a catalysis is needed.

Charges: The Stern-Gerlach Experiment

This iconic experiment in the realm of quantum mechanics describes the nature of the charge an atom can have. This idea is pertinent to biochemistry because of the nature of how monomers, dimers, etc. bind to metal centers (for examples, iron is the metal center that hemoglobin uses as its coordination center). These protein structures bind to metal centers by coordination complexes that are in correspondence with the oxidation state of the metal ion. The oxidation state of the metal ion is influenced by the spin charge of the metal ion and it because of the Stern-Gerlach experiment that explains the nature of charge in an atom. The Stern-Gerlach experiment is an iconic quantum mechanical experiment that physically showed the spin charges of a silver atom. The electrons of the silver atoms was emitted between two vertically oriented magnets. These magnets provided the magnetic field to separate the electrons based on their apparent spin. The result of this experiment were two spins, one going up and one going down, dictated as + 1/2 and -1/2. The two beams of directionally spun electrons were then put through a second set of horizontally oriented magnets. The result of this addition were again two beams of directionally different spun electrons except in the horizontal orientation. The last part of this experiment was sending one of the beams from the horizontally oriented spun electrons through another set of magnetic oriented in the same way as the first set. The resulting beams from this last set of magnetics is the most important realization. The main idea of the implementation with the third magnet is that two beams emerge, just as experienced with the first set of magnetics. Because of the beams of electrons shown are of the same spin nature, this in turns shows that the two beams of electrons from both magnetics after being oriented differently, are eigenvalues of each other, meaning they are of the same operator. This whole experiment displays the importance of charge and how it can be used to describe the nature of quantum behavior and the nature of coordination ligand binding.

Foundation of Biochemistry

Physics is one of the foundations of biochemistry; it deals with different types of energy and forces. Types of forces:

1. Ionic force — charge-charge interaction
2. Dipole interactions — deals with electronegativity (trend from increasing order: P, H, S, C, Br, Cl, N, O, F) and consists of hydrophobic interactions.
3. Van der Waals — molecular repulsion
4. Hydrogen bonds

Further reading

1. Nelson, David (2005). Principles of Biochemistry (4th Ed. ed.). Sara Tenney. ISBN 0-7167-4339-6. {{cite book}}: |edition= has extra text (help)
2. Levine, Ira N. (2002). Physical Chemistry (5th Ed. ed.). McGraw-Hill. ISBN 0-07-231808. {{cite book}}: |edition= has extra text (help); Check |isbn= value: length (help)

Introduction

Thermodynamics is the study of energy and its interconversion. It is the branch of physics that studies temperature, heat, and other macroscopic properties.

The science of thermodynamics began in the nineteenth century with the need to describe the operation of steam engines and to set limits of what they were capable of doing. Thus, the name itself denotes the power developed from heat, which used the steam engine as the initial example. Observations of engines were then generalized, which later became the first and second law of thermodynamics.

Thermodynamics also limit a cell's growth to a narrow temperature range. So, the heat increases the molecular movement within the proteins. The species grows within a specific thermal range since its proteins have evolved to endure that range. Proteins tend to denature or function very slow for growth if it is outside the range. [Microbiology]

In chemistry, thermodynamics predicts the spontaneity of a process rather than the kinetics of the process. It takes into account only the final and the initial states and does not require knowledge of the pathway between reactants and products. It is a state function. It is the study of heat in chemical reactions and the change of physical state in correspondence to the laws of thermodynamics.

A system defines the parameters of an object or a process that is being studied and everything else in this universe is considered the "surrounding." A system is considered to be in thermodynamics in equilibrium with another system, if both the systems have the same temperature. These two things are separated by this boundary that maybe imaginative. There are four main types of boundaries: fixed, real, movable, and imaginary This provides a volume for the system in which things such as work, heat, and matter between the system and the surroundings can be studied upon. .

The principles of thermodynamics are incorporated into the three basic laws. As an old joke summarizes them, the first law says you can't win, the second law says you can't break even and the third law says you can't leave the game.

Reference

1. Zumdahl, Chemistry Seventh Edition

2.Smith, J.M. (2005). Introduction to Chemical Engineering Thermodynamics. McGraw Hill. ISBN 978-007-127055-7. {{cite book}}: Text "coauthors+ H.C. Van Ness, M.M. Abbott" ignored (help)

Definition

The Zeroth law (also called equilibrium law) states that "if objects A and B are separate in thermal equilibrium with a third object C, then A and B are in thermal equilibrium with each other" (Jewett, Serway. Physics for Scientists and Engineers - 6th ed).

The Zeroth Law of thermodynamics focuses on the thermal equilibrium of two connected bodies in the same system. This thermal equilibrium is made apparent when an object of higher temperature transfers heat to an object that is of lower temperature. Eventually, both bodies reach the same temperature where the change in heat between the two is no longer measurable. They reach a constant temperature which exists between the two starting temperatures.

To take this definition one step further, let's consider a hypothetical situation of three systems, where system A is in contact with system B only, and system B is only in contact with system C, and we assume that A,B, and C are in thermal equilibrium then all three objects have the same temperature

History

This law was named and further studied by Ralph H. Fowler, a British physician and astronomer who contributed widely to physical chemistry and statistics. The Zeroth Law was developed long after the first three thermodynamic laws had been established. Scientists did not realize the law's extreme importance long after its discovery. Because the law was basic and laid the foundation for the rest of the thermodynamic laws, it was called the Zeroth Law rather than the Fourth Thermodynamic Law.

This law allows us to make quantitative measurements about the temperature of one system by relating it to the temperature of other systems.

Introduction

The First Law of Thermodynamics is an expression of the principle of the conservation of energy. It states that the total energy of a system and its surroundings is constant, and energy can be transformed, but may not be created or destroyed.

The First Law can be applied to living organisms by thinking of them as a system. A system cannot output more energy than it contains without an external source of more energy. Once the potential energy locked in carbohydrates and other energy sources are converted into kinetic energy (energy in use or motion) by the organism, the organism will not obtain more until energy is imputed again. This is important to understand since the variety of tasks that are performed by cells, ribosomes, proteins, etc. are only possible through the intake and transformation of pre-assembled molecules into energy.

The general equation that describes this theory is:

$dU=\delta Q-\delta W\,$

where dU is the change in internal energy, dq is the infinitesimal heat exchanged, and dw is the infinitesimal work performed. Work and heat are not state variables where as internal energy is. In this equation, it is demonstrated that only heat and work can lead to a change in the internal energy of a system, which is defined as the total of all kinetic and potential energy of everything within a closed system. This equation may also be described in words as: although energy assumes many forms, the total quantity of energy is constant; and when energy disappears in one form, it appears simultaneously in other forms.

In addition to work and external potential and kinetic energy, the generalization of the law of conservation of mechanical energy was made possible by the recognition of heat and internal energy as forms of energy. As a matter of fact, examples such as surface energy, electrical energy, and magnetic energy can all serve as extensions to the generalization stated above. The validity of this generalization was supported by overwhelming amount of evidence, which has raised its stature to a law of science, known as the First Law of Thermodynamics.

It is important to look at how the internal energy of system changes under constant pressure and temperature conditions because many chemical reactions take place under these specified conditions. Using the definition of the internal energy and assuming that only expansion work is done, one may write:

$dU=\delta Q-P*\delta V\,$

Expansion of the terms gives:

ΔU = U₂ - U₁ = Q_p + W = Q_p - P*ΔV = Q_p - P*(V₂ - V₁)

Finally, rearrangement for Q_p gives:

Q = (U₂ -P*V₂) - (U₁ - P*V₁)

Since U, P, and V are all state variables, we may define Q_p, the heat transferred at constant pressure, as a new state function called the enthalpy, H. We can represent this new state function as:

H = U + PV

Where H, U, and V are molar or unit-mass values. U denotes the internal energy of a system, P denotes pressure, and V denotes volume.

When heat is added to a system, the internal energy of the system will increase, which, in turn, increases enthalpy. Work done by the surroundings on that system will produce similar (positive work) results. Conversely, heat lost to the surroundings or work done by the system are given a negative sign for enthalpy and work. In other words, any change in the energy of a system must result in a corresponding change in the surroundings. Therefore, energy is transferred from the system to the surroundings (reset of the universe), or from the surroundings to the system. And thus, energy can neither be created nor destroyed.

The sphere of influence of the process is divided into two parts when the First Law of Thermodynamics is applied to a given process. The region in which the process occurs is set apart as the system; everything with which the system interacts is its surroundings. A system may be of any size; its boundaries may be real or imaginary, rigid or flexible. A system usually consists of a single substance; however, complex systems consist multiple substances may also be found in scientific and engineering applications. In any event, the equations of thermodynamics are written with reference to a well-defined system. Attention is often focused on the particular process of interest and on the equipment and material involved in the process directly. However, the First Law of Thermodynamics applies to the system and its surroundings; not to the system alone.

Conclusion

The first law of thermodynamics is that energy can neither be created or destroyed.

General Form

• accumulation = input - output
• input = rate of energy (potential, kinetic, internal) is inputted into the system + rate of heat entering the system
• output = rate of heat leaving the system + rate at which it leaves as work
• accumulation = rate of energy in the system

References

2.Smith, J.M. (2005). Introduction to Chemical Engineering Thermodynamics. McGraw Hill. ISBN 978-007-127055-7. {{cite book}}: Text "coauthors+ H.C. Van Ness, M.M. Abbott" ignored (help)

3. Berg, Jeremy (2012). Biochemistry. Freeman. ISBN 9781429229364. {{cite book}}: Text "coauthors+ John L. Tymoczko, Lubert Stryer" ignored (help)

Overview

The first law of thermodynamics states that energy is conserved, however, it only describes the transformations observed, and it doesn’t impose any restriction on the process direction. Nevertheless, such a restriction has been observed and proved to be exited in all thermodynamic applications. The need of a law describing this phenomenon gives rise to the second law of thermodynamics.

The Second Law of Thermodynamics states that the entropy of a closed system is constantly increasing with respect to time.
It is often said jokingly that the first law states that one cannot win and that the second law states that one cannot even break even.

The second law of thermodynamics may be expressed in two related statements as follows:

Statement 1: It is impossible to operate a system in such a way that heat absorbed by the system is completely converted into work done by the system.

Statement 2: It is impossible for a process to consist solely in the transfer of heat from one temperature level to a higher one.

Statement 1 is not contradictory to the first law of thermodynamics. Statement 1 does not imply that heat cannot be converted to work done by the system; it only implies that either the system or the surroundings have to be changed or modified when such a process takes place. As a corollary, any continuous production of work from heat is proved to be impossible. To compress the gas back to its initial state, energy must be drawn from the surroundings in the form of work; heat is transferred to the surroundings to maintain constant temperature at the same time. The amount of work gained from expansion is required by the reverse process described above, thus the production of net work is impossible.

Heat Conversion to Work

A thermodynamic theorem, the Carnot’s theorem is generated based on statement 2. The Carnot’s theorem states that no engine can have a higher thermal efficiency than that of a Carnot engine. Since a Carnot engine is reversible, it is able to transfer heat form one temperature level to a higher one. Although such an engine does not exist in the real world, it is the most efficient engine based on the laws of thermodynamics.

The study of heat engines, devices that are able to convert heat to work in a cyclical fashion, often serves as the classical approach to the second law. This macroscopic viewpoint of properties was able to be conducted without any knowledge of either the structure or the behavior of molecules. Any heat engines consist the following cycles: absorption of heat into the system at a relatively high temperature, rejection of heat to the surroundings at a relatively low temperature, and production of work. The two temperature levels are often referred to as heat reservoirs; the higher temperature level as the hot reservoir, and the lower temperature level as the cold reservoir. In thermodynamic applications, the working fluid, a liquid or a gas, connects the hot and the cold reservoirs in the sense that it absorbs heat from the hot reservoir, produces work, discards heat to the cold reservoir, and returns to its initial state to get ready for a new cycle.

Reversible Carnot Cycle

Spontaneous Processes and Entropy

Some processes proceed spontaneously in one direction, but not in the other. This process is similar to a gas diffusing to fill its container, but never really collecting at one end. These processes are said to be spontaneous; in other words, they occur without outside intervention.

Entropy, denoted by the symbol S, is the thermodynamic property that describes the spontaneity of a process. It is a macroscopic property of randomness or disorder, and is also a function that describes the number of arrangements (positions and/or energy levels) that are available to a system existing in a given state.

The Second Law of Thermodynamics can be explained with a simple example. Consider throwing a deck of ordered playing cards into the air and picking them all up at random. It is very improbable to pick up the cards in their original order because the probability for this to happen is so minute that we never really observe it. Entropy can be closely associated with probability. The more ways a particular state can be achieved, the greater the likelihood (probability) of finding that state. However, this does not mean that it is impossible for the cards to be put back in their original order or for the gas to be only at one end of the container. It merely is improbable.

There is a natural tendency toward disorder in closed systems because the state that has the highest probability of existing is where the system will be at equilibrium. The equilibrium state, such as within a container of gas, is referred to as the disordered state (where gas molecules are equally dispersed). The equilibrium state is the same as the disordered state because the gas molecules will occupy the largest volume possible, meaning they are all equally spaced out.

The change in entropy of the universe can be represented as

ΔS_univ=ΔS_sys+ΔS_surr

where ΔS_sys and ΔS_surr represents the changes in entropy that occur in the system and the surrounding respectively.

If ΔS_univ is positive, the entropy of the system increases, and the process is spontaneous in the direction written. A negative value for ΔS_univ indicates that the process is spontaneous in the opposite direction. The system is at equilibrium, and the process has no tendency to occur if ΔS_univ is equal to zero.

In some cases, particularly biological systems, it is hard to see how the entropy of universe is increasing. For example, when a leaf use carbon dioxide and nutrients to produce cellulose, the randomness and consequently entropy is decreasing. However, this process does not have any contradiction with second law of thermodynamic because it accompanies with increasing heat in the environment which increases the entropy.

===Phase changes and Entropy ===

Entropy has to do with the freedom of particle motion. As a result, when entropy increases in the system, it can cause a phase change from solid state to liquid state, and even to the gas state. In the solid state, the particles' movements are restricted and they have less freedom to move around in a fixed area. However, in a liquid state, particles have more freedom to move around. Therefore, in a gas state, the particles have so much greater freedom to move around. Consequently, entropy increases as one goes solid to liquid and to gas state.

When such disorder occurs, the energy of motion becomes more dispersed. For instance,when a salt is dissolved in liquid water, there would be more ions and solvent molecules interacting with each other. As a result, the solution's energy of motion is more dispersed. The greater the freedom the particles have, the more energy of motion they would dispersed.

                                solid ---> liquid ----> gas
less freedom for particles interaction --------> much greater freedom for particles interaction
                fixed energy of motion --------> dispersed energy of motion

Therefore, the change in phase states and the freedom of motion of particles can help to determine whether a reaction is spontaneous or not.

===The number of Microstates and Entropy ===

As stated in the previous section about the freedom of motion and the dispersed energy of motion, they are the two factors that can determine the direction of the spontaneous reaction. Silberberg defines microstates as "the quantized states of the whole system of a gas molecules". Microstates is about a gas molecules' state when it reacts with other molecules in the system. Consequently, there would be an increase in the energy of motion because the molecules vibrate and rotate around one another. In addition, there are different microstates for different conditions in the system. In thermodynamic terms, microstates can be related to entropy, the state of disorder, because the number of microstates is the number of ways that the thermal energy can be dispersed in the system. The equation is

                                       S= k ln W
     k (the Boltzmann constant)= R (gas constant)/ Avogadro's number= 1.38 x 10^ -23 J/K
     W= the number of microstates
     S is entropy

Therefore, entropy depends on the number of microstates.

           small number of microstates -----> much greater number of microstates
                           low entropy -----> high entropy

The Effect of Temperature on Spontaneity

Entropy changes in the surroundings ΔS_surr are primarily determined by heat flow.

The sign of ΔS_surr depends on the direction of the heat flow. In an exothermic process, the resulting energy flow increases the random motions in the surroundings, increasing the entropy of the surroundings (ΔS_surr is positive). Similarly, the tendency for systems to undergo changes that lower its energy can be explained by the fact that when a system at constant temperature moves to a lower energy state, the energy it gives up is transferred to the surroundings, leading to an increase in entropy there.

The magnitude of ΔS_surr depends on the temperature. At high temperatures, atoms in the surroundings are in rapid motion. A given quantity of energy transferred to the surroundings do not make a large percent change in their motions. Thus, the impact of the transfer of a given quantity of energy as heat to or from the surroundings is greater at lower temperatures, where the randomness of the surroundings experience a greater percent change. In other words, ΔS_surr depends directly on the quantity of heat transferred and inversely on temperature.

References

4.Smith, J.M. (2005). Introduction to Chemical Engineering Thermodynamics. McGraw Hill. ISBN 978-007-127055-7. {{cite book}}: Text "coauthors+ H.C. Van Ness, M.M. Abbott" ignored (help)

5. Silberberg, Martin S.(2010). Principles of General Chemistry (2nd Edition).McGraw Hill Publishing Company. ISBN978-0-07-351108-05 Structural Biochemistry/Carnot Cycle/

Introduction

The Third Law of Thermodynamics is a physical law regarding the role of entropy in nature. This law, along with the first two Laws of Thermodynamics are absolute, in that everything in the observable universe obeys these laws and, like time and gravity, nothing is exempt from them. The Third Law of Thermodynamics states that:

"As a system approaches absolute zero, all processes cease and the entropy of the system approaches a minimum value."

This minimum value, however, is not necessarily zero, although it is almost always zero in a perfect, pure crystalline substance. A perfect, pure crystal is one where the atomic and molecular arrangement is perfectly symmetrical and evenly distributed throughout the substance, and where each molecule is identical to one another. On the Kelvin scale, zero degrees Kelvin is the lowest mathematically possible temperature in the universe, corresponding to about -273.15° Celsius or -459.7 Fahrenheit.

In actuality, it is impossible for any real system to reach absolute zero, in part due to the Second Law of Thermodynamics, which says that heat transfer cannot occur spontaneously from a colder body to a hotter body. Thus, any system must draw energy from nearby systems as the system approaches absolute zero. Because it must draw energy, obtaining absolute zero is physically impossible and is a mathematical limit of the universe.

Entropy

The melting of ice is a common example of increasing entropy.

As a result of the Third Law of Thermodynamics, the concept that the entropy of a system, ΔS, reaches a constant (or 0 Kelvin for a perfect crystal) at a temperature of absolute zero is introduced. This is important because it provides a bottom foundation from which entropy can be measured. For any isothermal process that involves only substances in internal equilibrium, the entropy change goes to zero as the temperature approaches zero. Entropy quantifies the disorder of a system and is used to predict how a system will spontaneously change. Since entropy is proportional to the temperature of a system, once the temperature reaches its lowest point (absolute zero), the entropy, too, approaches zero. This idea can be visualized by drawing an example from water. Water at its most physically free state, water vapor, is fairly unaffected by intermolecular forces, and thus the molecules move and disperse freely and randomly throughout the atmosphere. Water in this form has very high entropy (randomness). As the temperature drops below 100° C and the gas condenses into liquid, the intermolecular forces come into play to a larger degree and thus the molecules move less freely. Water in this form has lost some entropy. As the liquid water approaches 0° C and freezes into solid ice, the intermolecular forces become extremely strong and the molecules can no longer move freely, but can only vibrate within the ice crystals. The entropy in this form is extremely low. As the water is cooled more and the temperature becomes closer and closer to absolute zero, the molecules would completely cease motion and the entropy at this state would have zero entropy.

Overview

In Thermodynamics, the total energy of a specific system is called the Internal Energy. It is the total amount of energy within the system, excluding the energy in the surroundings. Internal Energy can be divided into two parts: Kinetic Energy and Potential Energy. We say that the change in Internal Energy equals to ∆U, where the final Internal Energy is subtracted from the initial Internal Energy to obtain the change in Internal Energy, ∆U.

Internal Energy is a state function because of the fact that its value depends on the present state of the system, as it is independent of the path. In calculating the change in Internal Energy, by whatever process, as long as the initial and final states are the same, the Internal Energy does not differ. Another way to look at the Internal Energy of a system is to take the perspective of work in relation to Internal Energy. Either mechanical work from the change in pressure or volume can result in the change in Internal Energy. Overall, the internal energy of the system does increases as mass is added into the system, hence making Internal Energy an extensive property as it directly proportional to the amount of material in the system at that time.

Internal energy is defined as the energy associated with the random, disordered motion of molecules. It is separated in scale from the macroscopic ordered energy associated with moving objects. It refers to the invisible microscopic energy on the atomic and molecular scale. For example, a room temperature glass of water sitting on a table has no apparent energy, either potential or kinetic . But on the microscopic scale it is a seething mass of high speed molecules traveling at hundreds of meters per second. If the water were tossed across the room, this microscopic energy would not necessarily be changed when we superimpose an ordered large scale motion on the water as a whole.

Internal energy involves energy on the microscopic scale. For an ideal monoatomic gas, this is just the translational kinetic energy of the linear motion of the "hard sphere" type atoms , and the behavior of the system is well described by kinetic theory. However, for polyatomic gases there is rotational and vibrational kinetic energy as well. Then in liquids and solids, there is potential energy associated with the intermolecular attractive forces. A simplified visualization of the contributions to internal energy can be helpful in understanding phase transitions and other phenomena which involve internal energy.

More generally, while external energy is energy due to macroscopic motion (of the system as a whole) or to external fields, internal energy is all other forms of energy, including random motion (relative motion of molecules within the system) and dipole moments and stress.

Molecular Interpretation

Each molecule has a specific number of degrees of freedom, which include translations, rotation, or vibration. Yet, as the equipartition theorem states, the increase in thermal energy is even spread between degrees of freedom of the molecule in question. According to the equipartition theorem, the average energy of each contribution to the total energy is ½*kT so for a monatomic gas that only contains translation the equation looks like this: U_m (T)=U_m (0)+3/2 RT . With the inclusion of rotational energy for linear molecules the equation then becomes: U_m (T)=U_m (0)+5/2 RT . Now, with a non-linear molecules, including both transitional and rotational energies, the equation then becomes: U_m (T)=U_m (0)+3RT .

Internal Energy for an Ideal Gas

With the assumption that the gas being studied is ideal, the definition of Internal Energy can change. Since there are absolutely no intermolecular interactions in a perfect gas, so long as the distance between the molecules take no effect on energy. Then the assumption can be made that the Internal Energy of a gas to be independent of its volume.

The internal energy of an ideal gas is a function of the temperature only.

An ideal gas is defined as one in which all collisions between atoms or molecules are perfectly elastic and in which there are no intermolecular attractive forces. One can visualize it as a collection of perfectly hard spheres which collide but which otherwise do not interact with each other. In such a gas, all the internal energy is in the form of kinetic energy and any change in internal energy is accompanied by a change in temperature.

An ideal gas can be characterized by three state variables: absolute pressure (P), volume (V), and absolute temperature (T). The relationship between them may be deduced from kinetic theory and is called the Ideal Gas Law: PV=nRT

The ideal gas law can be viewed as arising from the kinetic pressure of gas molecules colliding with the walls of a container in accordance with Newton's laws. But there is also a statistical element in the determination of the average kinetic energy of those molecules. The temperature is taken to be proportional to this average kinetic energy; this invokes the idea of kinetic temperature.

Reference

Physical Chemistry by Atkins and de Paula (9th edition) Entropy (S) is the thermodynamic measure of randomness throughout a system (also simplified as “disorder”). Entropy can also be described as thermal energy not able to do work since energy becomes more evenly distributed as the system becomes more disordered. Entropy is particularly important when describing how energy is used and transferred within a system. As an exact value of entropy is impossible to measure; however, through relationships derived by both Josiah Willard Gibbs and James Clerk Maxwell the change in energy between one state and another can be calculated based on measurable functions, like temperature and pressure. That value in turn gives insight into how chemical reactions are favored and, most importantly, allows for the calculation of Gibbs Free Energy (ΔG = ΔH-TΔS).

Using statistical mechanics of the gas phase, entropy can be estimated by using Boltzmann’s formula. According to the formula, S = k ln W where k, the Boltzmann's constant, equals 1.381 x 10⁻²³ (in J/K). The Boltzmann's constant was calculated by relating to the gas constant R = kNA. W stands for the number of ways that the atoms or molecules in the sample can be arranged while still containing the same total energy.

It is important to note that the change in entropy, like temperature and volume, is a state function: the value is independent of the path used to get from the original state to the final state. Additionally the overall change in entropy of the universe is positive, meaning that the universe is continuously moving to a state of higher disorder.

A simple example where entropy is increased is when ice melts to water. The structure of ice is a well-ordered, crystalline system. When energy is put into the system in the form of heat, molecules begin to move more rapidly and no longer have the neatly ordered structure of ice. Thus, there distribution throughout space is more “random”. Another example where entropy is increased is when a reaction produces more moles of products than the reactants in the same phase.

The favorability of intramolecular reactions over intermolecular reactions is explained entropically. In an intermolecular coupling, two molecules come together to form one thus increasing the order in the system and decreasing the entropy. In an intramolecular reaction there is one molecule to start and one at the end which does not change the entropy of the system in an unfavorable way as is seen in intermolecular reactions.

Entropy can further be divided into thermal disorder, in which the entropy increases as heat is added to the system, and positional disorder, which related to the increase in entropy as the volume of the system is increased.

Entropy is also of particular interest in biochemistry as one of the unofficial definitions of life is an aggregate of molecules that work to decrease entropy in a certain localized area or volume. Additionally, it helps describe many phenomenons found in biochemical systems, which are described next.

Entropy in Biochemical Interactions

Entropy is a measure of the unavailable energy in a closed thermodynamic system that is also usually considered to be a measure of the system's disorder, that is a property of the system's state. It is varies directly with any reversible change in heat in the system and inverserly with the temperature of the system.

Entropy can be a strong driving force in nature. For example, it plays a very large part in the behavior of hydrophobic substances in water. A very common example of entropy at work would be lipids in solution. The lack of polarity in longer hydrocarbon chains tends to "force" water molecules to align themselves in an orderly pattern around the saturated part of the molecule. This orderly pattern decreases entropy as it prevents water from freely associating itself with other water molecules via hydrogen bonding. An increase in entropy would lead to a more negative Gibbs Free Energy, and a spontaneous reaction. The saturating effect of decreasing the change in entropy serves as the driving force for lipids to associate with one another instead of with water. Lipids coalesce to reduce the amount of water surrounding its molecules, and thereby increasing entropy. This phenomenon results in the formation of essential evolutionary components of life, such as lipid bilayer structures such as the lipid bilayer membranein eukaryotic cells.

Another place that the entropic favoring of hydrophobic molecules to dissociate with water can be found is in the active sites of enzymes. Many enzymes have a high concentration of hydrophobic residues in their active sites. The binding of an enzyme to its substrate alleviates the lack of entropy by driving water molecules out of the active site.

The value of understanding entropy’s role in chemistry can be utilized in the lab. For example, ammonium sulfate can be added in high concentrations to an aqueous solution containing at least one or more proteins. Proteins are much larger than ammonium sulfate ions. Thus the condensed charge of the newly dissolved salt attracts the water molecules in solution to form hydrate shells around them. In order to form these shells around ammonium sulfate, water molecules from the hydration shells around proteins must be used. The hydration shells of proteins are generally more ordered since not all of the protein surface is charged nor are the charges as condensed. Thus, the decrease in water molecules around proteins reaches a limit where the proteins become insoluble in water and are precipitated out for isolation and further study. .

However, entropy can also play a negative role in biochemistry.
For example, denatured protein by heating is an example where entropy plays a role in denaturation. In a folded protein, entropy is high due to its packed structure. As the protein becomes unfolded(denatured), the hydrophobic regions in a protein are surrounded by water. Overall there is no change in entropy but the protein becomes denatured.

Oil spill in sea also follows the same argument.

Entropy during phase changes

An entropy diagram for phase changes of water

As mentioned earlier, entropy is the measure of disorder, and this is also the case when it comes to phase changes.

This could be thought of a more simplified manner. For example, in the solid form of H₂O, they are in a very rigid and ordered crystal structure. As the temperature is raised, the rigid crystal structures begin to loosen up from its tight grasp of one another(via hydrogen bonds), and eventually, melting occurs. The ice turns into liquid. There was an increase in the disorder of this system, it went from a solid rigid crystal structure, to a bunch of freely moving molecules. This amount of disorder for this phase change to occur is entropy, specifically in the case of melting, the entropy of fusion.

History of Development

The enthalpy of reaction at specific values of T and P is defined as the heat exchanged between the system and the surroundings as the reactants are transformed into products at conditions of constant T and P. The heat flowing into the system is given a positive sign. Also, the standard enthalpy of formation refers to one mole of the specified reaction at a pressure of 1 bar. The term "enthalpy" and its accurate definition did not come about until the late 1840s. Before then, scientists and engineers knew from experience that a fluid called heat transfers from a hotter object to a cooler object when they are brought into contact. The result is that the cold object becomes warmer, and the hot object becomes cooler as heat transfers between them. Enthalpy is also the heat energy that absorbed or released as the reactants become the products at a constant pressure. When enthalpy is negative, the heat energy is released. However, when enthalpy is positive, the reactants absorb the heat from their surroundings.

A reasonable view is that there is something transferred from the hot object to the cold one. This “something” was referred as heat. Thus, it was convenient to say that heat always flows from higher temperature to lower one. This conclusion established the concept of temperature gradient, which may be viewed as the driving force for the transfer of energy as heat. Later, through conducting more precise measurements and experiments, scientists concluded that the rate of heat transfer from one object to another is proportional to the temperature difference between the two objects. Thus, the heat transfer is negligible when the temperature differences between the two objects are negligible.

A Thermodynamic View of Heat

Heat, or energy transferred between a hot and a cold object, is never regarded as being stored with an object from the thermodynamic point of view. Form this view point, heat is defined to exist only as energy is in transit from one object to another; or in the thermodynamic terminology, from the system to the surroundings. When heat is added or subtracted from a system or its surroundings, it is thought that energy is not stored in or being pulled away from the point of contact. Rather, heat is thought to act as kinetic and potential energy. The addition and subtraction processes are considered to be the transfer from potential energy to kinetic energy, and vice versa.

Relationship between Heat and Enthalpy

In spite of the transient nature of heat, the definition of heat is always viewed as a unique property of the object that is being measured. As mentioned above, temperature changes were used as the primary units to measure heat until the late 1840s. In fact, the definition of the unit of heat was based on the temperature change of a unit mass of water, as seen in the British thermal unit. The British thermal unit is commonly known as the Btu.

The British thermal unit is defined as 1/180th quantity of heat which when transferred to one kilogram mass of water raised its temperature from 0 to 100 degree Celsius at standard atmospheric pressure. This primitive definition of heat failed to provide an accurate measurement. The main reason accounts for this inaccuracy lie in the reference object, which in this case is water. The measure temperature change is partially determined by the purity of the water. The purer the water is, the more amount of heat is need for the temperature to increase by one unit.

In the twentieth century, more and more scientists realized that heat is just another form of energy that is being transferred by making contact of two objects. Thus it is convenient and practical to measure heat by using an energy unit.

The Modern Concept of Enthalpy

The modern concept of enthalpy or heat was developed largely based on the series of experiments carried out by James P. Joule (1818–1889). Joule’s experiment involves carefully measured amounts of water, oil, and mercury in an insulated container. He placed known amounts of these afore mentioned substances in the container, and agitated the fluids with a rotating stirrer. The mechanical work needed to carry out the rotation was carefully calculated, and the temperature changes of the fluid temperatures were accurately measured. In his result, Joule stated that for each fluid that a fixed amount of work was required per unit mass for every degree of temperature rise caused by the mechanical work provided. Furthermore, the original temperature of the fluids could be restored by the transfer of heat through contact as simple as to cooler objects. The quantitative relationship between heat and work was thus firmly established by Joule’s experiment, which proved definitively that heat is a form of energy.

In order to give heat a more concrete definition and moreover, to recognize a common basis for all energy units, international steam table calorie is defined in relation to joule, the SI unit of energy. Joule is the mechanical work done when a force of one Newton acts through a distance of one meter. Upon the establishment of the SI unit, the international steam table calorie was defined as equivalent to 4.1868 Joule (exact, by definition) and thermochemical calorie is equivalent to 4.184 Joule (exact, by definition). As a reference, one international steam table Btu is equivalent to 1055.056 Joule, and one thermochemical Btu is equivalent to 1054.35 Joule.

Mathematical Expression of Enthalpy

Enthalpy(H) is a measure of the heat energy of a reaction. The Enthalpy H of a thermodynamic system whose internal energy, pressure, and volume are U, P, and V is defined as H =U+PV. Since U, P, and V are state functions, H is a state function. note from dw=-PdV that the product of P and V has the dimensions of work and hence of energy. Therefore it is legitimate to add U and PV. Naturally, H has units of energy.

For example, a reaction that is exothermic will have a negative change in enthalpy. This is because the enthalpy of the products is less that that of the reactants.

The enthalpy of water formation.

In mathematical terms enthalpy of a reaction can be known as the following:

ΔH = (sums of bonds broken) - (sums of bonds formed)

ΔH = ΣnH_products - ΣnH_reactants

When ΔH is negative then the reaction is exothermic and more bonds are formed than broken.

If ΔH is positive then the reaction is endothermic. Chemical bonds tend to form spontaneously that the negative value of ΔH represents exothermic reaction.

The enthalpy for the reverse reaction is equal in magnitude, but opposite in mathematical sign.

Enthalpy is a state function and thus according to Hess's Law, the overall enthalpy of the reaction is equal to the sum of the enthalpies of the individual reactions or steps for which the overall reaction can be divided.

Since enthalpy is a function defined for the sake of calculations, it is difficult to measure the actual enthalpy of a substance. The change in enthalpy however, is easily measured and is an important quantity in many calculations.

When enthalpy changes, it signifies there is a change of state happening in the system. But enthalpies are reversible in the sense that the physical state changes are usually reversible. Therefore, since the enthalpy for formation of everything is always given, then at any given change, the new enthalpy can be calculated.

Reference

Slonczewski, Joan L.. Microbiology "An Evolving Science." Second Edition.

Smith, J.M. (2005). Introduction to Chemical Engineering Thermodynamics. McGraw Hill. ISBN 978-007-127055-7. {{cite book}}: Text "coauthors+ H.C. Van Ness, M.M. Abbott" ignored (help)

Engel, Thomas. Physical Chemistry. Third Edition.

Introduction

Heat is the energy transfer in body. The smaller the temperature change, the greater its capacity gets. The equation of heat capacity is as follow: $C=dQ/dT$

The heat capacity of an object is the amount of energy needed to raise the temperature of a substance 1 degree. The units are J/oC. Heat capacity is an extensive property. This means that a larger object has a larger heat capacity than a smaller object made from the same material. Heat Capacity = heat supplied / temperature rise The heat capacity of an object depends on both the quantity as well as the types of matter in the object. In order to compare heat capacities of different substances, we must relate heat capacity to the amount of material. One way to do this is to refer to a mole of substance. Then, the heat capacity will become the molar heat capacity. A more useful procedure is to compare heat capacities for one gram of material. This is called the specific heat capacity or simply specific heat. Specific heat is the quantity of heat required to increase the temperature of one gram of material one degree Celsius (or one kelvin). When divide the heat capacity of a material by its mass, we will have a specific heat. Specific heat = heat capacity / mass = C / m

To find the heat q required to raise the temperature of a sample by a certain amount, we multiply the specific heat of the substance, s, by the mass in grams, m, and the change in temperature,t. q = s x m x t

Furthermore, the heat capacity is given by the derivative of the internal energy with respect to temperature for a given energetic degree of freedom. There are many types of heat capacities: translational heat capacity, rotational heat capacity, vibrational heat capacity, and electronic heat capacity. [Physical Chemistry]. In the translational heat capacity, the translational energy-level spacing are extremely small. This makes the high temperature approximation is valid. In the rotational heat capacity, at the lowest temperaturs, there is insufficient thermal energy to provide for population of excited rotational energy levels. In contrast, as the heat capacity increases until the high temperature limit is reached. In the vibrational heat capacity, the high temperature limit is not applied to the vibrational degrees of freedom. The last heat capacity is the electronic heat capacity. In this heat capacity, there is no contribution to that constant volume heat capacity from the degrees of freedom since the partition function for the energetic degree of freedom is equal to the ground-state degeneracy. Also, the average energy is zero as well.

Heat Capacity at Constant Volume

$C_{V}=(dU/dT)_{V}$

It is usually called molar or specific heat capacity. U is the molar or specific internal energy, and T is the temperature. $U,C_{V},$ and $T$ are all state function.

It also can be written:

$dU=C_{V}dT$

During this process, the volume should always be constant. If the volume changes during the process, even the initial value and final value are the same; it is not a constant volume. However, because $U,C_{V}$, and$T$ are all state function, the equation applies to any process for initial and final values are the same.

Heat Capacity at Constant Pressure

$C_{P}=(dH/dT)_{P}$

$C_{P}$ is molar and specific heat capacities. H is the molar or specific enthalpy, and T is the temperature. This process is a closed system process.

It also can be written:

$dH=C_{P}dT$

The equation applies to any process for initial and values are the same, which means the pressure would not necessary be constant during the whole process because H, $C_{P}$, and T are all state function.

Importance of Heat Capacity

The heat capacity, otherwise known as the specific heat, of various molecules is extremely important for many different reasons. It is because of the specific heat of the water molecule that allows for life to exist on planet earth. For example, because of water's unique specific heat of 1 cal to 1 gram of water, a large body of water , like a lake, can absorb and store a huge amount of heat from the sun in the daytime and during summer while warming up only a few degrees due to allocation, diffusion and spreading of heat throughout the water system. This coupled with waters high specific heat only changes the overall temperature by such a small amount. During nights and winter season, the gradually cooling water can warm the air. This is the reason that contributes to coastal areas having milder climates than inland regions. The high specific heat of water also tends to stabilize ocean temperatures, creating a favorable environment for marine life. Thus because of its high specific heat, the water that covers most of Earth keeps temperature fluctuations on land and in water within limits that permit life. Also, because organisms are primarily made of water, they are more able to resist changes in their own temperature than if they were made of a liquid with a lower specific heat.

Reference

Engel, Thomas and Reid, Philip. Physical Chemistry. Pearson Education. Inc. 2006. Third Edition.

Smith, J. M., and Ness H. C. Van. Introduction to Chemical Engineering Thermodynamics. New York: McGraw-Hill, 1987. Print.

General Information

American scientist Josiah Willard Gibbs (1839-1903) created the theory of available energy, known as Gibbs Free Energy, in 1873. The theory relates the energy changes within the chemical reaction and how they depend upon the following quantities: enthalpy, temperature, reagents concentration and entropy of the system. In other words, these quantities will determine whether the reaction is favorable (exergonic) or not (endergonic).

The free energy change of a reaction (delta G) can tell us whether or not a reaction occurs spontaneously. Reactions that occur spontaneously have a negative delta G value, and such reactions are called exergonic. When delta G is positive, the reaction does not occur spontaneously, and the input of free energy is required for the reaction to proceed, thus it is called an endergonic reaction. When a system is at equilibrium where no net change occurs, then delta G is zero. The delta G of a reaction is the free energy of the final state minus the free energy of the initial state, making it is independent of the reaction pathway. However, the value of delta G provides no information on the rate of a reaction.

Gibbs Free Energy Equation

This is a Gibbs free energy graph by Josiah Willard Gibbs. it shows a plane perpendicular to the axis of v (volume) and passing through point A - represents the initial state of the body. MN is the section of the surface of dissipated energy. Qε and Qη are sections of the planes η = 0 and ε = 0, and therefore parallel to the axes of ε (internal energy) and η (entropy), respectively. AD and AE are the energy and entropy of the body in its initial state, AB and AC its Gibbs free energy and its capacity for entropy (the amount by which the entropy of the body can be increased without changing the energy of the body or increasing its volume) respectively.

We often focus on the use of the Gibbs free energy equation instead of its derivation. The most commonly-used equations for calculations are:

$\Delta G=\Delta H-T\Delta S\,$ (for constant temperature) - equation(1)

$\Delta G=-RT\ln K\,$ (for equilibrium constant that depends on temperature) - equation(2)

Where ΔH is change in enthalpy, T is the temperature of the system (in kelvin (K)), ΔS is change in entropy of the system, R is gas constant, K is equilibrium constant.

Numerical Meaning of ΔG

If ΔG < 0 (negative), then the reaction will proceed spontaneously, meaning the reaction is favorable (exergonic).

If ΔG > 0 (positive), then the reaction will not proceed spontaneously, meaning the reaction is unfavorable (endergonic).

If ΔG = 0 (equal to zero), then the reaction is at equilibrium.

In general, every system wants to achieve a minimum of free energy. Therefore, the more negative the Gibbs free energy, the more favorable the reaction.

Meaning of work to free energy

The sign of entropy (S) can determine whether a reaction is spontaneous or not. However, the sign of work (Delta H) cannot determine the spontaneous process. For instance, the exothermic reaction become spontaneous under certain conditions. And the endothermic reaction can also become spontaneous under different conditions. Silberberg used water as an example to explain such conditions.

"H₂O (l) ---> H₂O (s) Delta H of the reaction = -6.02 KJ (an exothermic reaction; spontaneous when T<0⁰C)

H₂O (s) ---> H₂O (l) Delta H of the reaction = +6.02 KJ (an endothermic reaction; spontaneous when T>0⁰C)"

In both reactions,the sign of enthalpy has no effect on the spontaneous change. Therefore, one cannot use enthalpy as a factor to determine the direction of a spontaneous reaction.

Standard Gibbs Free Energy of Formation

When we have to consider the relationship between Gibbs free energy and the standard-state free energy of a reaction, we use this equation:

$\Delta G=\Delta G^{\circ }+RT\ln Q\,$

to calculate Gibbs free energy at that of time under a specific circumstances. Where ΔG^o is the standard-states - reactants (or components) at 25^oC (degrees Celsius) and 1 atm (atmospheric pressure, 1 atm same as 100 kilopascals), Q is the reaction quotient. The motivation behind it is that these elements, reactants, and substances, are thermodynamically stable at such atmosphere.

Chemical potential

From elementary thermodynamics, Gibbs free energy, G, is defined as,

$dG=-SdT+VdP+\sum _{i=1}^{C}u_{i}dN_{i}$

$u_{i}$ is the partial molar Gibbs free energy of species i.

The chemical potential is not favorable for phase-equilibria calculation when the pressure approaches zero. Then, fugacity is used instead:

$f_{i}=C\exp(u_{i}/RT)$

$f_{i}$ is the partial fugacity of species i. C is a temperature-dependent constant.

Because fugacity has relationship with pressure, then fugacity coefficient of a pure species:

$\phi _{i}=f_{i}/P$

fugacity coefficient of a species in a mixture:

$\phi _{iV}=f_{iV}/(y_{i}P)$

Free Energy of Enzymes

Free energy determines whether a conversion of reactants to products will occur spontaneously. In the case of an enzyme, ΔG determines the rate of a reaction. Enzymes cannot affect thermodynamics of a reaction, and hence do not affect the equilibrium; Additionally, enzymes accelerate the attainment of equilibria but do not shift their positions. The equilibrium position is a function only of the free-energy difference between reactants and products. They are however, able to reach the equilibrium point at a far faster rate than without the presence of an enzyme.

For instance, in the presence of an enzyme, products could form within a second. On the other hand, products could take a as long as days to form without the presence of the catalyst. In both cases, concentration and amount of product formed remains entirely the same- it's equilibrium state. The amount of products it has formed has balanced with the amount of substrate.

Enzymes decrease only the free energy of activation- otherwise known as the activation energy. The Transition state between a substrate and the product is the point between a reaction where the substrates and products "meet in the middle". At this point, the highest free energy exists for the reaction. The activation energy is the energy it takes for a substrate to reach this transition state.

There are many competing theories of how enzymes actually bind their substrates, and each theory has a different graphic representation of the affect of the enzyme on the free energy of the reaction. In the lock and key mechanism theory, an enzyme has the pre-existing conformation to bind to a unique substrate. After binding and catalyzing the reaction, the enzyme will release the final products.

In the induced-fit mechanism theory, a similar approach is hypothesized. The only difference is that the pre-existing, unbound enzyme does not originally assume the exact conformation to bind the substrate; but rather assumes a slightly different structure prior to binding. Then, as the substrate binds to the enzyme, the structure of the active site conforms around the structure of the substrate to fit properly. Both of these mechanisms can be represented similarly in relation to their effect on the free energy of the reaction. Without really changing the pathway of the energy curve, these models serve to decrease the activation energy of a reaction, thereby increasing the rate of the reaction.

Another model has been suggested however, that appears slightly different on the free energy graph. This is the proposed transition-state model. This model suggests that an enzyme is not structurally adept to bind to the substrate itself, but that it is actually optimized to bind to the transition state of the reaction pathway. This produces a small stabilization of the transition state decreasing the overall activation energy as is characteristic of enzymes. The first increase in energy is due to the binding of the enzyme to the original substrate. The return to original free energy state is stabilization of the enzyme-substrate complex before reaction occurs. The next increase in energy comes from achieving the transition state, and the subsequent fall is the creation of the products. This theory is currently accepted as an alternative because the enzyme-substrate complex of the other theories acquires a very low free energy level due to stabilization. To achieve the transition state after this relatively low level of free energy is much more difficult than achieving the transition state from the relatively more energetically free enzyme-substrate complex suggested in this transition-state model. [6]

Formation of Double Helix

Formation of Double Helix

Double-stranded molecules of nucleic acids form the double helix structure such as DNA and RNA. Formation of double helix is one of biological process that the principles of thermodynamics are applied to it. In a solution containing single strands, all stands can easily move around, rotate, and disperse in the solution. In addition, forming conformation is easy in the single strands solution. However, when the double helix forms, it cannot displace as easy as two single strands could before. Moreover, it has less possible conformations. Thus, by forming double helix the randomness and entropy decrease.

Due to the Second Law of thermodynamics, significant heat has to be released to the surroundings for the process to be consistent with increasing the entropy of universe. Measuring the changing of temperature of a solution before and after formation of double helix reveals that approximately 250 kJ/mol (60 kcal/mol)heat is released. This large released energy is sufficient to overcome the effects of formation double helix - increase of order- and make universe more disorder.

Electrostatic freeenergy of the DNA doublehelix in terms of the counterion condensation theory: The polyelectrolyte theory reveals the formation of the double helix. The secondary structure of DNA is similar to the secondary structure of proteins. The number of condensed counterions is the same as for a line charge with charge density equal to the axial charge density of the helix. The logarithmic salt dependence of the electrostatic free energy is equivalent in range of lower salt concentration, thus the limiting laws stays constant. The helical parameters have a large influence on the overall electrostatic free energy and on the internal free energy of the condensed layer of counterions. The free energies of the single and double helix are negative at a higher salt level. Being negative indicates of the stabilization of the helical charge lattices electrostatic, because of mixing entropy of the condensed counterions. On the other hand, when the salt level is low, the free energy of a single helix is higher than the free energy of a double helix. With B-DNA parameters imagined as single helixes, the salt dependence of the free energy of transition from double strand to single is greatest at about 0.2 M salt, which is very similar to the area of the feature of separation of the DNA strand.The electrostatic freeenergy for the transition of the DNA doublehelix from the B to the A conformation can also be calculated. The Bform is the most electrostatically stable over the salt range. The electrostatic freeenergy values are close to the experimental values of the overall (electrostatic plus non-electrostatic) transition freeenergies for A-philic base pair sequences. B-to-A transition for A-philic sequences around concentration of 1 M is watched over by the polyelectrolyte properties of these two orientations of the DNA double helix. On the other hand the effect of ethanol cannot be tied to the lowering of the dielectric constant.

Bond Energies

How is energy being used? Is energy being consumed or absorbed in a reaction?

1) Bonds formed = Energy is released because it forms a more stable state. ΔH < 0 heat is released.

2) Bonds broken = Energy is absorbed because breaking a stable state and moving towards a less stable state. ΔH > 0 heat is absorbed.

Bond Energy products > Bond Energy reactants : spontaneous

Bond Energy products < Bond Energy reactants : non-spontaneous

Bond Energy Diagrams

ΔG = G products - G reactants

Note: You cannot switch the equation to be G reactants - G products.

The key is to understand if energy is being overall released or absorbed in a reaction. This will give you the correct sign for your ΔG.

Example of Gibbs Free Energy

References

Reece, Jane (2011). Biology. Pearson. ISBN 978-0-321-55823-7. {{cite book}}: Text "coauthors+ Lisa A. Urry, Michael L. Cain, Steven A. Wasserman, Peter V. Minorsky, Robert B. Jackson" ignored (help)

Seader, J. D., and Ernest J. Henley. Separation Process Principles. Hoboken, NJ: Wiley, 2006. Print.

Introduction

The definition of Material Equilibrium is that in each phase of the closed system, the numbers of moles of each substance in that phase remains constant in time. Material equilibrium can be subdivided into Reaction equilibrium and Phase Equilibrium. Reaction Equilibrium is the equilibrium where the conversion of quantity stops between two sets of chemicals. Phase Equilibrium is where the transport of matter reaches a balance point without conversion of one species to another.

Entropy and Equilibrium

In an isolated system which is not at material equilibrium. The spontaneous chemical reactions occur between difference phases in this system are irreversible processes that increase the entropy.Stot>0

Introduction of phase diagrams

This is an example of a phase diagram. In this case, the substance represented on the PT diagram is for water. The dotted region above the critical point is the supercritical fluid.

A phase diagram is a chart that helps define the conditions at which a substance can be in its solid, liquid, and vapor states. The three boundary lines of the chart signify the equilibrium relationships between the phases: fusion curve, vaporization curve, and sublimation curve. There are also two distinct points on the graph: the triple point and the critical point. To better define these terms the picture diagram can be examined. The red line is the transition between solid to gas and is also called the sublimation curve. The green line is from solid to liquid and is also called the fusion curve. The blue is liquid to gas and is called the vaporization curve. The point where all three lines intersect is the triple point. At this temperature and pressure the three phases: solid, liquid, and vapor can coexist. The critical point is the condition of the highest pressure and temperature that the substance can be observed in vapor/liquid equilibrium.

Supercritical

Any condition that exceeds that highest pressure and temperature becomes a region that has no definite phase boundaries. A phase is usually considered a liquid when vaporization comes from pressure reduction at constant T, the same goes for a gas if condensation results from temperature reduction at constant T. Since this process cannot be applied to either situation the region beyond the highest P and T can be regarded as supercritical and as the fluid region.

A supercritical fluid is defined as a substance above its critical temperature (Tc) and critical pressure (Pc). The critical point represents the highest temperature and pressure at which the substance can exist as a vapour and liquid in equilibrium. Supercritical fluids are highly compressed gases which combine properties of gases and liquids in an intriguing manner. Fluids such as supercritical xenon, ethane and carbon dioxide offer a range of unusual chemical possibilities in both synthetic and analytical chemistry. Supercritical fluids have solvent power similar to a light hydrocarbon for most solutes. However, fluorinated compounds are often more soluble in scCO2 than in hydrocarbons. This increased solubility is important for polymerisation.

Phase Diagram for proteins

This is a phase diagram used in protein crystallization. It displays the necessary protein and precipitant concentrations for crystallization.

Though phase diagrams are generally viewed as a map that shows the state of a substance at certain conditions, this simple tool can be used to analyze other complicated materials. An example is that the information derived from a phase diagram can be used for the discussion of protein crystallography. This diagram is a representation of what protein and precipitant concentrations are needed for protein crystallography. The red line is called the solubility line and it distinctly separates undersaturated conditions from the supersaturated, which is the condition desired for protein crystallization. The supersaturated part of the diagram consists of three parts: metastable, labile, and precipitation.

Supersaturated

Metastable-Crystals here can grow larger from seeds but they cannot nucleate. There is a lower protein and precipitant concentration with less aggregation events. This zone signifies the concentrations that are usually too slow for crystallization to occur.

Labile-Crystals here can also grow larger from seeds but they can nucleate. There is a higher protein and precipitant concentration with more aggregation events. This zone is also known as the nucleation zone or the crystallization zone and is more favorable for the formation of crystals.

Precipitation-Proteins usually surface as an aggregate or precipitate and are not used for crystallography. This section of the diagram displays less favorable conditions because aggregation and precipitation occur much faster than crystallization.

Nucleation: process of creating a nucleus.

Aggregation: the clumping or clustering of proteins.

Use of phase diagrams for protein crystallization

The connection between phase diagrams and protein crystallization is the placement of the solubility curve (the red line). As described, crystals can only form in supersaturated solutions and having information on the location of the solubility curve allows for the growth of crystals for X-ray crystallography. Though crystals are needed to determine the solubility curve, it is through various trials that the suitable concentrations for crystal formation are found. By this accumulation of data the crystals can be used as a guide for refining conditions for protein crystallization.

This phase diagram is used also because it unveils the interactions between the many components of the solutions. Examples of such interactions are the liquid-liquid phase separation (LLPS) that define the attraction between proteins, a necessary characteristic for crystallization of proteins. This attraction is thoroughly defined in LLPS. Other data that can be extracted from the phase diagram include enthalpies (from LLPS) and entropies (from solubility curve) of the protein in liquid and solid states. Also, additives can change the curves of the phase diagram giving data on the effect of the additive and how it interacts with the protein analyte. It can also be utilized for predictions of the conditions that are needed for a variety of protein crystallization which would be ideal for the crystallization process, less trial and error, more crystallizing.

Liquid-liquid phase separation (LLPS)

A result of adding precipitant to a protein solution is the formation of liquid drops. These liquid drops are sometimes referred to as oils or coacervates and can be observed at changing temperatures, pHs, and other solution condition changes. The drops usually contain high protein concentrations and when influenced by a higher force, gravity, the drops separate from the rest of the solution. The result of this is two liquid phases from the solution. This entire process is called liquid-liquid phase separation. LLPS is useful because the high concentrations of proteins help increase saturation to a supersaturated state which, as a result, increases crystallization.

The left picture shows LLPS of bovine crystallin in sodium phosphate. The picture on the right shows LLPS of thaumatin.

Introduction

More than one phases coexist at the same temperature and pressure.

K-Values

K-values is a phase-equilibrium ratio, for vapor-liquid systems,

$K_{i}=y_{i}/x_{i}$

$y_{i}$ is the mole fraction of a species in vapor phase, and $x_{i}$ is the mole fraction of a species in liquid phase.

For ideal solutions,

$K_{i}=P_{i}^{S}/P$

K-values can be used for calculating bubble and dew points.

Reference

Seader, J. D., and Ernest J. Henley. Separation Process Principles. Hoboken, NJ: Wiley, 2006. Print.

General information

endergonic reaction

An endergonic reaction refers to a chemical reaction in which energy is being used in the overall reaction, making the reaction non-spontaneous and thermodynamiacally unfavorable. Energy is being absorbed as the reaction proceeds, and there is a net loss of energy in the surrounding system. Due to this consumption of energy, standard change in Gibbs free energy (ΔG) is a positive value under constant pressure and temperature: ΔG° > 0.

The magnitude of ΔG also represents the quantity of energy required to drive the reaction. If a chemical process is exergonic in one direction, then the reverse process must be endergonic. Plants get the required energy to make sugar from the environment by capturing light and converting it into chemical energy that can be used for other processes.

Some examples of endergonic reactions are muscle contractions and protein synthesis.

The equilibrium constant of endergonic reaction where ΔG° > 0 is less than 1: K < 1.

Endergonic reactions require an input of energy, usually larger than those of non spontaneous exergonic reactions, from an outside source to disturb the chemical equilibrium to cause changes, such as bond formation. This input of energy is called the activation energy. In certain reactions, a catalyst is available to speed up endergonic reactions. A catalyst can lower the activation energy barrier for the reaction. Thus, it speeds up the reaction process. The energy for an endergonic reaction is obtained by coupling the reaction with an exergonic reaction.

A familiar example of coupling exergonic reactions to endergonic reactions to promote spontaneity comes from ATP. ATP powers cellular work by coupling exergonic reactions to endergonic reactions. It is responsible for mediating most energy coupling in cells, and in most cases, acts as the immediate source of energy that powers cellular work. The synthesis of the amino acid glutamine from glutamic acid and ammonia is naturally endergonic and non spontaneous with a ΔG value of +3.4 kcal/mol but coupling this reaction with the exergonic process of ATP hydrolysis, -7.3 kcal/mol, will drive the reaction forward, making it spontaneous. In the same sense, a exergonic reaction must be coupled with the formation of ATP from ADP in order to make the reaction spontaneous and in most cases, cellular respiration provides the energy for the endergonic process of making ATP and plants use light energy, instead, to produce ATP.

An endergonic reaction can simply be understood by studying the following situation. In a chemical reaction, the reactants make products and an equilibrium is reached. An endergonic reaction is one where more products are made from the equilibrium amount by disturbing the equilibrium with a form of energy. For example, heat will be absorbed into the system and the equilibrium will shift to the right (towards product side). Consequently, more products will be formed.

Summary

An endergonic reaction:

• The free energy of initial state < free energy of final state
• Energy needs to be put into the system in order to go from the initial start to the final state
• +ΔG

References

Zumdahl, Chemistry Seventh Edition

Neil A. Campbell, Jan B Reece. Biology Seventh Edition, 2005 Pearson Education, Inc.

http://academic.brooklyn.cuny.edu/biology/bio4fv/page/exergon.htm

General information

exergonic reaction

An exergonic reaction refers to a reaction where energy is released. Because the reactants lose energy (G decreases), Gibbs free energy (ΔG) is negative under constant temperature and pressure. These reactions usually do not require energy to proceed, and therefore occur spontaneously. In a chemical reaction, breaking and forming bonds between atoms is a form of energy. Since chemical reactions mainly consist of forming and/or breaking chemical bonds, exergonic reactions release energy by breaking less stable chemical bonds and forming more stable bonds. The example of exergonic reactions occur in our body is cellular respiration: C6H12O6 (glucose) + 6 O2 -> 6 CO2 + 6 H2O this reaction release energy which is used for cell activities.

However, some exergonic reactions do not occur spontaneously and require a small input of energy to start the reaction. This input of energy is called activation energy. Once the activation energy requirement is fulfilled by an outside source, the reaction proceeds to break bonds and form new bonds and energy is released as the reaction takes place. This results in a net gain in energy in the surrounding system, and a net loss in energy from the reaction system; hence, the change in Gibbs free energy (ΔG)] is negative (ΔG < 0). A negative ΔG denotes that the reaction is spontaneous and thermodynamically favorable.

Here is an endergonic reactions of ATP give energy. First Breaking down the ATP formed ADP and Pi is an exergonic reaction, where Delta G is less than 0. However, by combining the reaction glucose+Pi ---> glucose 6-phosphate, a thermodynamically unfavorable reaction. The cell can drive an endergonic reaction.

Conclusion

File:Exergonic.png

Exergonic reaction is

• Free energy of final state < free energy of initial state
  
• Free energy is free during the reaction
  
• The energy can be used to do biological work
  
• Able to work in spontaneous reaction
  
• doesn’t have to occur in great rate which is where the enzyme comes in
  
• -ΔG (negative number)
  

Reference

http://academic.brooklyn.cuny.edu/biology/bio4fv/page/exergon.htm Different from classical thermodynamics, essentially a deductive science, molecular thermodynamics focuses on the properties of individual chemical species and their mixtures at molecular level.

History of Development

By the end of the nineteenth century, most laws and postulates of classical thermodynamics were well-established. However, the rapid pace of development in scientific fields relating to chemistry, physics, and chemical engineering urges scientists and engineers to have a more holistic view on the subject of thermodynamics, which has given rise to the birth of molecular thermodynamics.

More often than not, the source of property values is experiment. For instance, the ideal gas equation was evolved as a statement of observed physical behavior of gases and their interconnecting relationships between volume, pressure, temperature, and number of moles of gas present.

At the turn of the century, physicists and chemists who worked with principles of classical thermodynamics increasingly realized that experiments at macroscopic level often failed to provide any insight into why substances exhibit their observed properties. By conducting further experiments, they found that the basis for insight should rather be established on a microscopic view of matter.

Intermolecular Forces

Intermolecular forces are relatively weak forces between molecules in random motion. The energy resulted from this random motion is referred as internal energy.

The ideal gas model was the very first model introduced when internal energy between molecules was studied. An ideal gas is characterized by the absence of molecular interactions. However, ideal gas also possesses internal energy. Real gases, on the other hand, are composed of molecules that have not only the energy of individual molecules, but also energy shared between molecules due to their interactions. The intermolecular potential energy is associated with collections of molecules, and the intermolecular forces are reflected by the existence of energy in this form.

Pair-Potential Function

Two molecules attract each other when they are far apart and repel each another when they are close together. This fact was also established through the study of molecular thermodynamics. A sketch of intermolecular potential energy may reveal that the potential energy for an isolated pair of spherically symmetric neutral molecules is solely dependent on the distance which separates them.

If we let U denote intermolecular potential energy, F denote intermolecular force, and r denote the distance separating the two spherically symmetric neutral molecules, the intermolecular force may be expressed as a function of intermolecular potential energy and distance separating the two spherically symmetric neutral molecules as:

F(r) = - dU(r) / dr

The negative sign shown in the above equation signifies an intermolecular attraction, whereas a positive sign indicates an intermolecular repulsion.

The above differential equation is also referred as the pair-potential function. Specific values of U and r in this form may appear as species dependent parameters in a pair-potential function.

Internal Energy from A Microscopic View

Kinetic theory and statistical mechanics are the two theories that relate the behavior of molecules from microscopic level to macroscopic level. Many thermodynamic properties, such as internal energy, enthalpy, and entropy were able to be explained after the development of the two theories mentioned above. Together these two theories represent nearly all the knowledge we possess about molecular thermodynamics.

Before further exploring molecular thermodynamics from a molecular level, it is crucial to understand how the energy associated with each individual molecules of an ideal gas relates to the macroscopic internal energy of a system defined.

Energy, no matter external or internal, is quantized from a quantum mechanics point of view. In other words, the total amount of internal energy of a system may be treated and analyzed as tiny measurable units that carry discrete amount of energy. One of such an energy unit is often referred as quanta. Thus there are enormous numbers of quanta contained in a system, and the sum of the quanta determines the energy level of the system. The set of energy levels allowed to exist to a closed system, as specified by quantum theory, is determined by its volume. Each energy level of a system has a quantum states associated with its energy level, which is also known as the degeneracy of the level.

Statistical Mechanics' Contribution

From a molecular thermodynamic point of view, the state of a system is firmly established if and only if the temperature and volume of the system are defined. Nevertheless, a fixed temperature and a fixed volume do not guarantee an equilibrium is reached within the system. In the case of ideal gas, the random motion and collision of the gas molecules with each other and the wall of the container result in exchanges of energy with the surroundings. Momentary fluctuations caused by these random motions and collisions shift the energy level back and forth within the system. Therefore, it makes sense to define an average value over the discrete set of energy levels of the allowed quantum states. Moreover, statistical mechanics is sufficient in providing the means for arriving at the proper average value.

One of the fundamental postulates of statistical mechanics for a system with defined volume and temperature is that the probability of a quantum state depends only on its energy. The importance of this postulate is that it relates the energy level of a system with its probability. In other words, all quantum states with the same energy have the same probability. Following the same logic as stated above, a value for the thermodynamic internal energy may be obtained as the average of the energies of the quantum states, which is equivalent to its probability.

Reference

Smith, J.M. (2005). Introduction to Chemical Engineering Thermodynamics. McGraw Hill. ISBN 978-007-127055-7. {{cite book}}: Text "coauthors+ H.C. Van Ness, M.M. Abbott" ignored (help) Equilibrium may be defined as lack of change, or a static condition. In thermodynamics, equilibrium not only implies the lack of change, it also refers to the absence of any tendency toward change on a macroscopic scale.

==Relationship between Equilibrium, Driving Force, and Resistance in a System==

A system of interest at equilibrium exists if and only if there is no net change in the surrounding conditions. In this case, the driving force can be viewed as negligible also since the absence of any tendency toward change also indicates the absence of any driving force. Thus, all forces are said to be in balance when equilibrium exists in a system. One thing that needs to be pointed out here is that whether a change actually occurs in a system not at equilibrium depends on two factors: the resistance as well as the driving force of the system.

Different Driving Forces Produce Different Kinds of Changes

Many kinds of changes might be encountered in a system; these different changes are caused by different driving forces. Driving forces may be classified by their physical properties, which include pressure, temperature, concentration, gradients in chemical potential, and so on.

For example, imbalance of mechanical forces, such as the pressure gradient on a piston, will tend to cause energy transfer in the form of work. Gradients in chemical potentials tend to cause substances to be transferred from one phase to another; temperature differences tend to cause the flow of heat in or out of the system. Nevertheless, the final results of the above listed tendencies are the same--equilibrium will be achieved in a system.

Equilibrium, Free Energy, and Reaction Direction

Thermodynamic equilibrium refes to a condition in which equilibrium exists with respect to P, T, and concentration. Equilibrium is established with respect to a given variable only if that variable does not change with time, and if it has the same value in all parts of the system and surroundings. The equilibrium with respect to concentration exists only if transport of all species across the boundary in both directions is possible. If the boundary is a movable wall that is not permeable to all species. The sign of free Energy (ΔG) can determine the direction of a spontaneous reaction. In addition, one can determine whether the reaction will proceed to the right or the left based on the reaction quotient (Q) and the equilibrium constant (K). As a result, Silberberg states that

If Q < K or Q/K <1 , then the reaction will head to the right

If Q > K or Q/K >1 , then the reaction will head to the left

If Q = K or Q/K =1 , then the reaction will be at equilibrium.

As a result, the equation that combines the equilibrium constant, free energy, and reaction quotient is ΔG= RT ln (Q/K)

Different Driving Forces Produce Different Kinds of Changes

Many kinds of changes might be encountered in a system; these different changes are caused by different driving forces. Driving forces may be classified by their physical properties, which include pressure, temperature, concentration, gradients in chemical potential, and so on.

For example, imbalance of mechanical forces, such as the pressure gradient on a piston, will tend to cause energy transfer in the form of work. Gradients in chemical potentials tend to cause substances to be transferred from one phase to another; temperature differences tend to cause the flow of heat in or out of the system. Nevertheless, the final results of the above listed tendencies are the same--equilibrium will be achieved in a system.

Effects of Chemical Reactions on Thermodynamic Equilibrium

In most of the applications of thermodynamics, the effects of chemical reactions are usually ignored. Since a system may stay in a long-term equilibrium if a chemical reaction is not initiated, it is convenient to not consider the effects that a chemical reaction might bring about to the system. Furthermore, equilibrium is mostly likely to be reached after a chemical reaction is carried out in most thermodynamic cases. Therefore, a purely physical process may be analyzed without regard to possible chemical reactions.

References

Smith, J.M. (2005). Introduction to Chemical Engineering Thermodynamics. McGraw Hill. ISBN 978-007-127055-7. {{cite book}}: Text "coauthors+ H.C. Van Ness, M.M. Abbott" ignored (help)

Silberberg, Martin S.(2010). Principles of General Chemistry (2nd Edition).McGraw Hill Publishing Company. ISBN978-0-07-351108-05

Equation of State

PV = nRT

Ideal gas: A hypothetical gas that exhibits linear relationships among volume, pressure, temperature, and amount (mol) at all conditions; approximately by simple gases at ordinary conditions. Although no ideal gas actually exists, most simple gases, such as N2, O2, H2, and the noble gases, show nearly ideal behavior at ordinary temperatures and pressures. Ideal gas law: An equation that expresses the relationships among volume, pressure, temperature, and amount (mol) of an ideal gas: PV=nRT.

Types

There are three basic classes of ideal gas:

• the classical or Maxwell-Boltzmann ideal gas,
• the ideal quantum Bose gas, composed of bosons, and
• the ideal quantum Fermi gas, composed of fermions.

The classical ideal gas can be separated into two types: The classical thermodynamic ideal gas and the ideal quantum Boltzmann gas. Both are essentially the same, except that the classical thermodynamic ideal gas is based on classical statistical mechanics, and certain thermodynamic parameters such as the entropy are only specified to within an undetermined additive constant. The ideal quantum Boltzmann gas overcomes this limitation by taking the limit of the quantum Bose gas and quantum Fermi gas in the limit of high temperature to specify these additive constants. The behavior of a quantum Boltzmann gas is the same as that of a classical ideal gas except for the specification of these constants. The results of the quantum Boltzmann gas are used in a number of cases including the Sackur-Tetrode equation for the entropy of an ideal gas and the Saha ionization equation for a weakly ionized plasma.

Internal Energy

a function of temperature only
U = U(T)

Internal energy: is the total of kinetic and potential energies of all the particles in a system.

Implied Property Relations for an Ideal Gas

C_v is a function of temperature only

C_(v )≡(∂U/∂T)_v= dU(T)/dT= C_v (T)
H is a function of temperature only

H ≡U+PV=U(T)+RT=H(T)
C_p is a function of temperature only

C_p= dH/dT= dU/dT+R=C_v+R
Any change of state of an ideal gas
dU = C_(v ) dT ∆U= ∫▒C_(v ) dT
dH = C_(p dT) ∆H= ∫▒C_(p ) dT

Some facts about Ideal Gas Law: The ideal gas law can be rearranged to calculate the density and molar mass of a gas. Also, in a mixture of gases, each component contributes its own partial pressure to the total pressure (this is also called the Dalton's law of partial pressures). The mole fraction of each component is the ratio of its partial pressure to the total pressure. The most important idea is that the total pressure is the sum of the gas pressure and the vapor pressure of water at the given temperature when a gas is in contact with the water. [The Molecular Nature of Matter and Change].

Reference

Silberberg, Martin S. Chemistry: The Molecular Nature of Matter and Change. 5th ed. 2009

Equation for Process Calculations for Ideal Gases

For reversible, closed-system, work is given by
dW=-PdV
For ideal gases, the first law can be written by
dQ+dW=C_v dT
From two equations above, we get
dQ=C_v dT+PdV
These three equations can be applied to four types of processes: isothermal, isobaric, isochoric and adiabatic.

Isothermal Process

File:Isothermal.png

From left to right the lines signify: isochoric, adiabatic, isothermal, and isobaric.

Isothermal process deals with closed-system that has constant temperature. So ΔT=0:
ΔU=ΔH=0
Q=RTln V_2/V_1 =-RTln P_2/P_1
W=-RTln V_2/V_1 =RTln P_2/P_1
Q=-W (constant T)

Therefore,
Q=-W=RT ln V_2/V_1 = -RT ln P_2/P_1 (constant T)

Isobaric Process

File:Isobaric.png

From left to right the lines signify: isochoric, adiabatic, isothermal, and isobaric.

Isobaric process deals with closed-system that has constant pressure. So ΔP=0.
ΔU=∫▒〖C_v dT〗 and ΔH=∫▒〖C_p dT〗
Q=∫▒〖C_p dT〗 and W=-R(T_2-T_1)
Therefore,
Q=ΔH=∫▒〖C_p dT〗 (constant P)

Isochoric Process

File:Isochoric.png

From left to right the lines signify: isochoric, adiabatic, isothermal, and isobaric.

Isochoric process deals with closed-system that has constant volume. So ΔV=0.
ΔU=∫▒〖C_v dT〗 and ΔH=∫▒〖C_p dT〗
Q=∫▒〖C_v dT〗 and W=-∫▒PdV=0
Therefore,
Q=ΔU=∫▒〖C_v dT〗 (constant V)

Adiabatic Process

File:Adiabatic1.png

From left to right the lines signify: isochoric, adiabatic, isothermal, and isobaric.

Adiabatic process deals with closed-system that has no heat transfer between the system and the surroundings. So ΔQ=0.

dT/T= -R/C_v dV/V

T_2/T_1 =(V_1/V_2 )^(R⁄C_v )

T_2/T_1 =(P_2/P_1 )^(R⁄C_p ) and P_2/P_1 =(V_1/V_2 )^(C_p⁄C_v )

The following equations apply to ideal gases with constant heat capacities that undergo mechanically reversible adiabatic expansion or compression.

〖TV〗^(γ-1)=constant 〖TP〗^(((1-γ))⁄γ)=constant 〖PV〗^γ=constant

γ≡ C_p/C_v

For any adiabatic closed-system,

dW=dU= C_v dT

W= △U= C_v△T

γ ≡ C_p/C_v = (C_v+R)/C_v =1+ R/C_v or C_v= R/(γ-1)

W= C_v△T= (R△T)/(γ-1)

W= (〖RT〗_2-〖RT〗_1)/(γ-1)= (P_2 V_2-P_1 V_1)/(γ-1)

For mechanically reversible process,

W= (P_1 V_1)/(γ-1) [(P_2/P_1 )^(γ-1)-1]= (RT_1)/(γ-1) [(P_2/P_1 )^(((γ-1))⁄γ)-1]

Diabatic Process

Opposite of adiabatic process
There is heat transfer

Polytropic Process

Polytropic process deals with a model of some versatility. So δ=constant.
〖PV〗^δ=constant

〖TV〗^(δ-1)=constant 〖TP〗^(((1-δ))⁄δ)=constant

W= (RT_1)/(δ-1) [(P_2/P_1 )^(((δ-1))⁄δ)-1]

Q= ((δ-γ)RT_1)/((δ-1)(γ-1)) [(P_2/P_1 )^(((δ-1))⁄δ)-1]

Reference

Smith, J. M., and Ness H. C. Van. Introduction to Chemical Engineering Thermodynamics. New York: McGraw-Hill, 1987. Print.

Irreversible Process

For mechanically reversible, closed-system processes,

• every equation is derived.
  
• the equations have property changes- dU,dH,△U,△H.
  

This is applied to reversible and irreversible processes in both closed and open systems.
An irreversible process is a two-step process.

1. W is for mechanically reversible process.
   
2. Through efficiency, the results are multiplied or divided to get the actual work.
   

If work is produced, then the absolute value for the reversible process will be too big and has to be multiplied by the efficiency.
If work is required, then the reversible process value will be too small and has to be divided by the efficiency.

Reference

Smith, J. M., and Ness H. C. Van. Introduction to Chemical Engineering Thermodynamics. New York: McGraw-Hill, 1987. Print.

Application of the Virial Equations

2 forms of virial expansion are infinite series.

1. Z=1+B^' P+C^' P^2+D^('P^3 )+⋯
   
2. Z=1+B/V+C/V^2 +D/V^3 +…
   

• B^'= B/RT
  
• C^'=(C-B^2)/〖(RT)〗^2
  
• D^'= (D-3BC+2B^3)/〖(RT)〗^3
  

When we derive the relations, set Z = PV/RT and solve for P.

Reference

Smith, J. M., and Ness H. C. Van. Introduction to Chemical Engineering Thermodynamics. New York: McGraw-Hill, 1987. Print. Bioenergetics refers to the transformation of energy that occurs within living organisms. In order to fuel the chemical mechanisms within cells, organisms require an input of energy. This energy is used to drive chemical reactions and help store and process information, which is essential in propagating life. Energy may be obtained from sunlight, in which case the organisms are referred to as phototrophs, or it may be extracted from chemicals, in which case the organisms are referred to as chemotrophs. Because energy may not be available at all times to fuel these life processes, organisms have adapted mechanisms to couple chemical reactions so that exergonic reactions can provide energy for those that are endorgernic.

The chemical reactions performed by an organism make up its metabolism. Catabolic reactions involve the break down of chemical molecules, while anabolic reactions involve the synthesis of compounds.

The Laws of Thermodynamics

The energy processes in living organisms are defined by the basic laws of thermodynamics. The first law dictates that the total energy present in the universe always remains constant (Note: though the energy total is static, it often changes forms such as when an animal converts the chemical energy of food to mechanical energy as it moves). Meanwhile, the second law asserts that the total entropy present in the universe is ever increasing.

The Energy Process

For phototrophs, sunlight takes the form of potential energy, while complex molecules serve as potential energy for chemotrophs. In living organisms, energy is commonly used in the form of work. This work energy is acquired by breaking weak bonds and forming stronger bonds. Work may take the form of synthetic chemical reactions, maintaining chemical and ionic gradients (homeostasis), and transferring genetic information. This most important form of work is the polymerization of information-containing macromolecules, such as protein, DNA, and RNA. These macromolecules form the basis of life and are what ultimately drive life’s processes. Though mechanisms have been adapted to make the most efficient use of the acquired energy, some of it is inevitably released in the form of heat or metabolic waste products. These waste products can be eliminated from the body of an organism and consumed by bacteria or other organisms that are able to extract energy from them. Often, materials and compounds useless to one organism are energy sources for another and by this exchange energy is constantly recycled.

Energy Coupling

In order to increase energetic efficiency, cells often couple reactions together. Endergonic reactions are those that require an input of energy. Exergonic reactions are those that release energy. By coupling these two reactions together, the overall chemical process is made exergonic so it can occur spontaneously due to a negative free-energy change. Using this process, unfavorable chemical reactions can be made to proceed.

The basis of reaction coupling is a shared chemical intermediate. After one reaction produces one product, another can use it as a reactant to drive the production of an essential compound.

Gibbs Free Energy

Named after Josiah Willard Gibbs who developed the concept in 1878, the Gibbs free energy describes the overall favorability of a reaction to proceed. It is characterized by the following equation: ΔG = ΔH - TΔS where ΔG is the Gibbs free energy change, ΔH is the change in enthalpy (heat), T is temperature (measured in Kelvins), and ΔS is the change in entropy. A reaction, or process, will occur spontaneously if and only if ΔG < 0 for that reaction. If ΔG > 0 then that reaction must be coupled (see Energy Coupling) with a reaction which has ΔG < 0 such that the overall ΔG of the 2 reactions is <0. ATP hydrolysis is a commonly used reaction for such situations (see ATP). As the equation shows, a reaction is more favorable, the more ΔH < 0 (i.e. the more heat is given off by the reaction) and the more ΔS > 0 (i.e. the more disorder is increased) and the effect of entropy gained or lost is magnified by the temperature surrounding the reaction.

ATP

Adenosine Triphosphate (ATP)

Most coupling reactions use the break down of adenosine triphosphate (ATP) as the intermediate process to drive chemical synthesis. ATP is used as an energy-storing compound. The phosphoanhydride bond between phosphate groups found in ATP stores a significant amount of energy due to the negative charges carried by the phosphate groups. This bond stored energy that is not currently use, but available later for running reactions is called Potential energy. This energy is required to keep the negatively charged groups close to each other in the ATP molecule because they, as do all like-charged groups, repel each other. This energy can be released (Exothermic reaction) for use in the cell to do work, move things and build things by hydrolysis and breakage of the bond.

ATP -> ADP + P + Energy

When carbohydrates and other foods are consumed, they are broken down by enzymes to release the energy within them. The exothermic energy released is used to reattach a phosphate to ADP through Endothermic reactions which will regenerate ATP formation.

ADP + P + Energy -> ATP

Then, the process of bond breaking and bond forming will repeat over and over within human cells to provide energy for all the chemical reactions.

GTP

Guanosine Triphosphate (GTP)

GTP (guanosine-5’-triphosphate) can be used as a source of energy, just like ATP but not in any type of organisms. However GTP is only used in specific areas of the cell, namely protein synthesis. ATP and GTP are similar in structure; both have a purine base, and 3 phosphate groups, but ATP has adenine attached to the purine, whereas GTP has a guanine. The energy stored in GTP is released in the same way as ATP.

Feedback

In order for the body to maintain homeostasis, feedback loops are often used. Negative feedback maintains homeostasis by slowing down or stopping a mechanism once it approaches the appropriate range. For instance, the hydrolysis of ATP is an exergonic reaction, meaning it gives off energy in the form of heat. If the temperature in a cell becomes too high, the cell will die, so negative feedback will stop the reaction to maintain homeostasis in terms of temperature. In positive feedback, a process will speed up once a receptor detects the occurrence of a certain reaction. A positive feedback mechanism is used in the stomach during digestion: for example, HCl secreted by parietal cells in the stomach convert pepsinogen to pepsin, and this reaction causes the pepsin to convert all the pepsinogen to pepsin to aid in the enzymatic breakdown of proteins. Another example of a positive feedback mechanism is childbirth, once the contractions start, they begin to occur with increasing frequency and pressure.

Energy Coupling

The total ∆G is negative because of the coupling of reactions. In this diagram, ∆G(1) stands for the change in G resulting from the reaction of glucose + Pi --> Glucose 6-phosphate, ∆G(2) is the free energy resulting from the reaction of ATP --> ADP + Pi, and ∆G3 is the total change in free energy by coupling these two reactions together.

Many chemicals' reactions are not spontaneous and require energy to occur. The spontaneity of a chemical reaction is determined by its Gibbs free energy value. If negative, the reaction will proceed spontaneously; if positive, the reaction will not be spontaneous. Note that this is not equivalent to kinetics, or the speed of a reaction; Gibbs free energy only determines the spontaneity of a given reaction. How quickly the reaction proceeds is determined by other factors, such as the presence of enzymes, the amount of heat (energy) that the system has available, and the physical properties of the reacting molecules.

Spontaneous reactions occur without the need for extra energy, but they may happen slowly. In order to catalyze non-spontaneous reactions, such as the synthesis of macromolecules, enzymes and coupling are used by the cell. Almost all reactions that take place in a cell are catalyzed by enzymes that decrease the activation energy of the reaction. Essentially, this means that the enzyme opens up a more favorable "pathway" for the reaction, allowing it to initiate more easily than before, and with less energy. In addition, an unfavorable reaction can be coupled together with a favorable one to make the overall reaction favorable. For example, Glucose + P_i -> Glucose-6-phosphate has a positive $delta$ G and is therefore unfavorable. But it can be coupled with ATP -> ADP + P_i (which has a negative $delta$ G) to make the reaction favorable. The overall reaction thus becomes ATP + Glucose -> ADP + Glucose-6-phosphate and has a negative $delta$ G. Therefore, ATP is considered the energy currency of the cell. However, it should be noted that other energy carrying molecules, such as GTP, do exist and are used for certain processes.

Anabolic and catabolic processes are examples of how the cell couples reactions together to create efficient energy exchange cycles. These processes are explained in more detail in their respective sections; however, it is appropriate to mention them here as they provide a relevant example of reaction coupling. Basically, catabolic reactions are those that convert chemical fuels to molecules that the cell can use for energy, such as ATP and other high-energy compounds. Anabolic reactions are those that require some amount of energy to occur. Thus, the cell can conveniently couple anabolic reactions with catabolic ones - the products of catabolic reactions can be used to drive anabolic reactions to completion. This allows the cell to link different types of reactions together efficiently; it's almost a type of "cellular recycling", as the products of one reaction (i.e. a catabolic one) can be re-used to help another reaction reach completion (i.e. an anabolic one).

Organisms are Energy Transducers

Organisms are transducers of energy, since in the transfer of energy they are less than 100% efficient. Organisms employ the energy harnessed to grow, repair, and maintain their bodies. The energy is also use to compete with other organisms, and to produce new organisms (offspring). In the process of doing these things, organisms generate waste, chemicals and heat. Organisms create local regions of order at the expense of using up some fraction of the total supply of useful energy found in the universe.

General Information

Metabolism

Metabolism regulates life through a set of chemical reactions. Chemical reactions are often coordinated with each other and occur in sequence called metabolic pathways, each step of which is catalyzed by a specific enzyme. These pathways are categorized according to whether the reactions lead to the breakdown or synthesis of substances. Catabolic reactions result in the breakdown of molecules into smaller molecules. Such reactions are often exergonic. By comparison, anabolic reactions promote the synthesis of larger molecules from smaller molecules. This process usually is endergonic and, in living cells, must be coupled to an exergonic reaction. These processes are responsible for the growth and reproduction of organisms, maintaining their structures, and responding to changes in the environment. The involvement of enzymes is essential for metabolism because they couple the organisms, which are thermodynamically unfavorable, to other organisms which are thermodynamically favorable and drive the metabolism towards desirable reactions. Not only do enzymes drive organisms toward desirable reactions, but they also regulate metabolic pathways to respond to changes within the cell's environment as well as signals from other cells. The metabolism of an organism also establishes which substances which enter the organism are nutritious and beneficial and which are harmful. Additionally, the speed of metabolism, or the metabolic rate, of an organism affects how much food the organism consumes. Prokaryote's metabolism is diverse. In other words, prokaryotes run all the major nutrient cycles. They play a major role in the Sulfur cycle and biological process that affects oxidation states of minerals in the earths crust. Also, the cyanobacteria both invented photosynthesis and still dominate the carbon fixation on this planet. We can also harness energy from redox couples by releasing all at once or releasing gradually (chemical bonds).

In Metabolism, there are something that we need for life. For example, we need a) the energy/reducing equivalents (ATP, NADH, NADPH). Also, we need b) Carbon Skeletons (glucose, glycine, etc...). We also need c) Other minor stuff (NPK, metals, etc...).

About the various Trophs: 1) Chemotroph: is the energy from redox 2) Phototroph: is the energy from light 3) Heterotroph: is the carbon from organic molecule 4) Autotroph: is the carbon fixed from CO2

In metabolism, glycolysis and the TCA cycle can be used just to produce molecules for growth. About Anaerobic Respiration: It is the use of a terminal electron acceptor other than O2. Moreover, the process produce generally less energy than using O2. Also, it can be considered the Redox Tower.

Assimilative Vs. Dissimilative Metabolism: 1) Assimilative Reduction: Compound gets reduced and incorporated into the organism. For example: NO3- becomes amino group -NH2. 2) Dissimilative Reduction: Compound gets reduced as an electron acceptor and is discarded. For example: NO3- becomes amino group -NH2.

Nitrate Reduction and Denitrification: a) Nitrite is one of the most common alternative electron acceptors. b) Denitrification removes nitrogen from systems as gas. Agriculture is bad and sewage is good.

- Nitrate Reduction: Nitrate Reductase is the critical protein for NO3- reduction. Nitrate reductase production is repressed by O2 and activated by NO3-. There are fewer protons are pumped. - Denitrification: Utilizes 4 reductive enzymes. It can start with nitrite from nitrate reduction. Also, the process can produce more energy than nitrate reduction. - Reduction of Various Metals: a) Geobacter metallireducens is a gram-proteobacterium. It can reduce a number of metals by oxidation of organic compounds. For examples, for iron: Fe3+ to Fe2+; for manganese: Mn4+ to Mn2+; for uranium: U6+ to U4+. Fe3+ compounds are much less soluble than Fe2+ iron compounds. U4+ compounds are much less soluble than U6+ compounds.

Chemoautotrophs: - Generate energy (ATP) from oxidation of inorganic compounds. - Carbon fixation of CO2 is by Calvin Cycle (same as plants/cyanobacteria). - Using the Electron Tower, we can predict metabolisms.

Hydrogen Oxidation: - Many bacteria produce H2 as a metabolic byproduct. - The critical enzymes are hydrogenases.

The Phototrophs: a) Photoautotrophs: Use light for energy and CO2 as carbon source. 1) Oxygenic: Oxidize H2O for electrons generating O2. 2) Anoxygenic: Oxidize other compounds for electrons (ie: H2S). b) Photoheterotrophs: Use light for energy and organic carbon as carbon source.

Photosynthetic Pigments: - All phototrophs contain chlorophyll or bacteriochlorophyll. - Each chlorophyll variant has different absorption spectra. - Different pigments allow bacteria to co-exist in one environment. - All chlorophylls have a porphyrin ring with different substitutions.

Accessory Pigments: a) Function: expand light spectrum and photoprotection. b) Main Pigments: Carotenoids and Phycobilins.

Photosynthetic Structures: - The purpose: a) increase surface area for light absorption. b) cluster proteins/pigments for electron transfer. - Bacteria do not have chloroplasts...they are the chloroplast. - Membrane Folds: Highly folded membrane invaginations. - Chlorosomes: Found in green sulfur and non-sulfur bacteria. - Phycobilisomes: Specific to cyanobacteria.

Reference: Slonczewski, Joan L. Microbiology. 2nd ed. New York, 2009.

Hydrothermal Vents: - Chemoautotrophs are the producers of these communities. - Do not require light for survival. - Can require no input from outside organisms.

Acid Mine Drainage (Fe2+ Oxidation): - Fe2+ is stable in acidic water (won't spontaneously oxidize). - Fe2+ gives e- for pumping out cytoplasmic H+. - e- is also used in reverse electron flow to produce NADH.

Anoxygenic photosynthesis: a) Represented Phyla: - Proteobacteria (purple bacteria) - Green sulfur bacteria - Green non-sulfur bacteria - Heliobacteria - Acidobacteria b) Only a single light reaction occurs. c) Electrons are derived from non H2O.

Purple Sulfur Bacteria: - They are gram negative and they are proteobacteria. They use bacteriochlorophyll a and b. They produce visible sulfur granules. They are generally found in anoxic environments. They thrive in H2S rich environments (hot spring).

Oxygenic Photosynthesis: a) Represented Phyla: - Cyanobacteria - Plants b) Electrons flow through two light reactions. c) Electrons are derived from H2O.

Cyanobacteria: a) Require oxygen for growth b) Primarily use phycobili-proteins and chlorophyll-a c) Grow in a wide range of environments (some extreme) d) Many species can fix nitrogen

Nitrogen Fixation: a) Groups: - Cyanobacteria - Rhizobia - Green sulfur - Azotobacteria b) Very energy intensive process (16-24 ATP) c) Nitrogenase enzyme is O2 sensitive d) Heterocyst Specific Properties: - No active photosynthesis - Transfer N via Glutamate - Specialized extracellular matrix of glycolipids and polysaccharides - Contain polar bodies to inhibit gas exchange

Storage of Nutrients: a) Bacteria Store Nutrients as Polymers - Low solute concentration - Inert b) Nitrogen Storage - Cyanophycin c) Carbon Storage - Glycogen (starch) - PHB - Oils

Glycogen Metabolism in Cyanobacteria: a) Glycogen functions as a major carbon storage polymer in cyanobacteria. b) The genes glgA and glgP show evidence of transcriptional regulation while GlgC is allosterically regulated. c) Glycogen synthesis and degradation is diurnal in cyanobacteria.

Cyanobacteria as Biofuel Producers: - Nitrogen Fixation - Hydrogen Evolution - Flotation/Motility - Dense Mat Growth - Extremophiles (Temp, ph) - Diverse Secondary Metabolites - Naturally Transformable

Fatty Acid Secretion: - Cyanobacteria do not have thioesterases, they come from plants. - Glycogen synthesis competes with pathways for lipid production.

Reference: Slonczewski, Joan L. Microbiology. 2nd ed. New York, 2009.

• ATP

Structure of ATP

During catabolism, useful energy is temporarily conserved as ATP - adenosine triphosphate. ATP is the universal standard of energy exchange in biological systems as the energy is always transformed and conserved as ATP .

• NAD

Structure of NAD

NAD (Nicotinamide Adenine Dinucleotide)also involves in metabolic pathways. The chemical nature of metabolic pathways usually involves oxidation/reduction reactions. In order for a biochemical to be oxidized, its electrons must be removed by an oxidizing agent. The oxidizing agent is an electron acceptor that gets reduced in the reaction. The molecule that usually functions as the electron carrier in the biochemical oxidation-reduction reactions is NAD and its phosphorylated derivative, NADP. NAD or NADP can become alternately oxidized or reduced by the loss or gain of two electrons. The oxidized form of NAD is symbolized NAD; the reduced form is symbolized as NADH₂.

Metabolic pathways are regulated in three general ways:

1.Gene Regulation Because enzymes in every metabolic pathway are encoded by genes, cells can control chemical reactions via gene regulations. For example, if a bacterial cell is not exposed to a particular sugar in its environment, it will turn off the genes that encode the enzymes that are needed to break down the sugar. Alternatively, if the sugar becomes available, the genes are switched on.

2.Cellular Regulation Metabolism is also coordinated at the cellular level. Cells integrate signals from their environment and adjust their chemical reactions to adapt to those signals. Cell-signaling pathways often lead to the activation of protein kinases that covalently attach phosphate groups to target proteins. For example, when people are frightened, they secrete a hormone called epinephrine into their bloodstream. This hormone binds to the surface of muscle cells and stimulates an intracellular pathway that leads to the phosphorylation of intracellular proteins, including the enzymes involved in carbohydrate metabolism. These activated enzymes promote the supply of energy to the frightened individual. Epinephrine is sometimes called the “fight or flight” hormone because the added energy prepares an individual to either stay and fight or run away. After a person is no longer frightened, hormone levels drop and other enzymes called phosphates remove the phosphate groups from enzymes, thereby restoring the original level of carbohydrate metabolism.

3.Biochemical regulation Metabolic reactions can also be controlled by reactions at the biochemical level. In this case, the binding of a molecule to an enzyme directly regulates its function. Biochemical regulation is typically categorized according to the site where the regulatory molecule binds.

Reference: Biology. Brooker. Widmaier. Graham. Stiling. Chapter Seven, Enzymes and cellular respiration.

Example(s) of Metabolic Pathways

Gluconeogenesis

This metabolic pathway involves the generation of glucose from non-carbohydrate carbon substrates (i.e. lactate, glycerol, and gucogenic amino acids). This is one of two main mechanisms (the other being glycolysis) that the human body uses to keep blood glucose levels from dropping to a dangerously low level, a condition called hypoglycemia. This mechanism isn’t exclusive to humans, and is also present in plants, animals, fungi and other microorganisms, with some variation in the locations in which glucogenesis takes place. This mechanism kicks in during periods of fasting, starvation, or intense exercise and is endergonic. It is also associated with ketosis and has been a target of therapy for Type II Diabetes to inhibit glucose formation and stimulate glucose uptake by cells.

The pathway itself consists of eleven enzyme-catalyzed reactions, which can begin in the mitochondria or cytoplasm (depends on the substrate being used). Many of these steps are the reversible reactions of those found in glycolysis.

•It begins in the mitochondria with the formation of oxaloacetate through carboxylation of pyruvate. This part requires ATP and catalytic help by pyruvate carboxylase, which is stimulated by high levels of acetyl-CoA and inhibited by high levels of ADP.

•Oxaloacetate is then reduced to malate using NADH, which will prepare it for mitochondrial exit. Afterwards, it is oxidised in the cytoplasm again to oxaloacetate using NAD+, in which the remaining steps of glucogenesis will take place.

•The next step is the decarboxylation and phosphorylation of oxaloacetate to produce phosphoenolpyruvate, which is catalyzed by phosphoenolpyruvate carboxykinase (PEP carboxykinase). This step also hydrolyses one molecule of GTP to GDP.

•The next steps of the reaction are essentially the same as those involved in reversed glycolysis, with the only difference being that fructose-1, 6-bisphosphatase converts fructose-1,60bisphosphate to fructose-6-phosphate. Note that this conversion is the rate limiting step of the whole process of glucogenesis.

•Next, glucose-6-phosphate is formed from fructose 6-phosphate with the help of phophoglucoisomerase. This product can be used in other metabolic pathways or can be further dephosphorylated to make free glucose. Cell control of intracellular glucose levels is attained by the fact that free glucose can diffuse in and out of the cell, whereas the phosphorylated form is locked in the cell. Glucose formation happens in the lumen of the endoplasmic reticulum. Here, glucose-6-phosphate is hydrolyzed by glucose-6-phosphatase to produce glucose, which is then shuttled into the cytosol by glucose transporters located in the membrane of the endoplasmic reticulum.

Relation to Obesity

As the availability of energy-rich food increased, people started to gain more weight by converting excessive energy into body fat. While this kind of environmental factors play a significant role in increasing rate of obesity, lipid based metabolism in the body is also partly in charge of phenomenon.

Diabetes is one of the common metabolic diseases in relation to obesity. There are two types of diabetes:

Type 1 diabetes: Type 1 diabetes is an autoimmune disease which usually starts before 20 years of age. It is caused by destroying insulin-secreting beta cells in the pancreas. Thus, the person with Type 1 needs insulin to stay alive.

Type 2 diabetes: Most people have type 2 diabetes, in which they have a higher level of insulin in their blood (unlike Type 1 diabetics); however, they are unresponsive to a hormone, insulin resistance. Type 2 diabetes is the most common metabolic disease currently. Also, obesity is one of the main factor for developing type 2 diabetes.

Metabolic syndrome

Obesity is one of the main factors to the development of insulin resistance, which leads to type 2 diabetes. The clustering of insulin resistance, hyperglycemia, dyslipidemia is called metabolic syndrome and is presumed to be a precursor of type 2 diabetes.

One of the reasons for obesity is the amount of triacylglycerides one consumes will exceed the adipose tissue's capacity. Thus, other tissues will begin to store the excess fat (usually the liver and muscle)/

The extra fatty acids in the muscles alter metabolism

The mitochondria is not able to process all of the fatty acids by Beta oxidation. Thus, the extra fatty acids accumulate into the mitochondria and eventually go into the cytoplasm. The inability of the mitochondria to process these fatty acids leads to the fatty acids forming into triacylglyverols and then amount of fat increases in the cytoplasm.

Metabolic Linkage Between Diabetes and Cancer

The HBP binds various metabolic inputs to optimally deliver the synthesis of UDPCLcNAc, which is the donor substrate for OGT. Glucose is funneled into the HBP where it will get phosphorylated by the hexokinase to produce glucose 6-phosphate (Glc-6-P) which is then transformed to fructose 6-phosphate (Fruc-6-P) through the presence of phosphoglucose isomerase. During the rate-limiting procedure of the pathway, glutamine:fructose-6-phosphate (GFAT) converts Fruc-6-P into glucosamine 6-phosphate (GlcN-6-O). Enzymatic steps will guide the production of UDP-GLcNAc, which is termed as a negative feedback inhibitor of GFAT. Flux through the HBP and the generation of UDP-GlcNAc and O-GlcNAcylation are greatly affected by the disease states, which include diabetes (indicated by the purple arrows) and cancer (indicated by the teal arrows) through metabolism (glucose= green, amino acid =red, fatty acid = orange, and nucleotide = blue). Resistance of insulin leads to an increase in glucose levels that marks a specific diabetic condition. As a consequence, the increase in glucose will cause HBP to increase HBP flux through the rise of O-GlcNAcylation. Cancer cells require energy, and metabolites are often in excess which can be funneled into the HBP stimulating flux also, precisely glucose, glutamine, and UTP.

O-linked β-N-acetylgluosamine (O-GlcNAc) is a metabolic signaling sugar molecule. More specifically, O-GlcNAc is a post-translational protein modification that is made up of a single N-acetylglucosamine piece, which is attached to an O-β-glycosidic linkage to serine and threonine hydroxyl moieties on nuclear and cytoplasmic proteins. Proteins that contain O-GlcNAc take part in cellular processes such as transcription, translation, signal transduction, and cytoskeletal assembly, along with other functions. Studies have shown that diseases like diabetes and cancer are strongly affiliated with the major alterations in metabolism, which affect the alterations of O-GlcNAcyclation. Such alterations to O-GlcNAcylation interfere with cellular signaling forces and worsen the disease state.

O-GlcNAc signaling is closely associated with cellular metabolism, and ties very closely to phosphorylation due to post-translational modifications that process swiftly in response to internal and external signs. Furthermore, the similarity between O-GlcNAc and phosphorylation is founded in the sugar’s capability to be dynamically connected or detached based on the changed in the cellular environment activated by stress, hormone, or nutrients. Since O-GlcNAc is linked to the serine and threonine residues, the sugar is strictly competing with phosphorylation. Steric hindrance can be encountered when o-GlcNAcylated and phosphorylated residues are in close proximity to each other. O-GlcNAcylation is catalyzed by an enzyme known as uridine diphospho-N-acetylglucosamine (UDP-GlcNAc): polypeptide β-N-acetylglusaminyltransferase (OGT). In most cells, the enzyme OGT dynamically creates many specific holoenzyme proteins complexes that monitor specific activity approaching the myriad of target protein substrates. On the contrary, phosphorylation includes many individual unique kinases. Just like how there are protein phosphatases that detach phosphorylation, there exist a single cytosolic or nuclear β-N-acetylglucosaminidase (OGA) that aim substrates by shaping transient holoenzyme complexes to remove the sugar component.

An important glucosamine, Hexosamine biosynthetic pathway (HBP), which is a prominent precursor in the synthesis of glycosylated proteins, accommodates a collection of metabolic components that are relative to the formation of Uridine diphosphate N-acetylglucosamine (UDP-GlcNAc). To do this, the HBP integrates various metabolic inputs that will essentially deliver the synthesis of UDP-GlcNAc, which is implicitly the donor substrate for OGT. Seemingly, a rise in cellular glucose and flux to a particular level of concentration via HBP will increase UDP-GlcNAc levels to a certain degree. Slight increases in UDP-GlcNAc concentration will suffice O-GlcNAc to function as a nutrient sensor. This is due to the unique responsive property of O-GlcNAc transferase to UDP-GlcNAc concentrations. Rising levels of flux via the HBP causes a resistance for insulin. Flux through the HBP, which leads to the production of UDP-GLcNAc and O-GlcNAcylation are influenced by diseases like diabetes and cancer through metabolic effects. Studies have proven that O-GlcNAcylation play a strict role in insulin signaling. When there is O-GlcNAc and OGT that disrupt insulin signaling, the antipodal of decreased O-GlcNAcylation and O-GlcNAcase stimulate proper insulin signaling. Under specific diabetic conditions, such low levels of O-GlcNAcylation are advantageous in the sense it relieves hepatic insulin resistance and saves the diabetic cardiomyocytes functionality.

Much clarity is needed to gain a better understanding of O-GlcNAc signaling. For instance, uncertainty in the approach of how OGT and OGA target their substrates remains a struggle in fully grasping the mechanistic abilities of this specific metabolic signaling molecule. Development in such research can unravel metabolic diseases like diabetes and cancer.

Biochemicals

Amino acids, carbohydrates, and lipids are essential for life; therefore, metabolism focuses on the production of these molecules during the creation of cells and tissues, and the digestion and use of them when they are broken down and used as a provider of energy.

Amino Acids/Proteins

When amino acids arrange themselves as a linear chain joined together by peptide bonds, proteins are formed. Many proteins are enzymes that catalyze the chemical reactions involved metabolism.

The name protein came from the Greek word proteios, meaning "first place." In bacterial cells, almost 50% of the dry mass is made up of proteins. Almost all organisms contain proteins. All functions of living organisms are related to proteins and each of their specific functions.

Proteins can be classified based on their functions in the cell:

 1. Enzymatic proteins
      a. Specifically speed up reactions that are endogenic.
      b. This is the largest group of proteins.
      c. Enzymatic proteins are responsible for metabolic related reactions in cells.
      d. Examples:
          1) Digestive enzymes catalyze the hydrolysis of foods. 
          2)  DNA- and RNA-polymerases
          3)  Dehydrogenases
 2. Structual proteins
      a. Support the shape of the cell.
      b. Maintain the structure in tissues.
      c. Examples:
          1) Collagen is a type of fibrous framework that makes up the connective tissues in animals.
          2) Keratin is a type of fibrous proteins that supplements hair, horns, feathers, and skin.
 3. Storage proteins
      a. Store amino acids.
      b. Contain energy that can be released in metabolic reactions
      c. Examples:
          1) Ovalbumin is a protein used as an amino acid source for the developing embryo in egg whites.
          2) The casein protein is a major source of amino acid for baby mammals in milk.
 4. Transport proteins
      a. Transport substances.
      b. Examples:
          1) Hemoglobin is a protein of verberate blood that carries oxygen from the lungs to other parts of the body.
          2) Membrane protein attaches to the membrane of the cell, transporting substances that are unable to cross the membrane themselves.
 5. Hormonal proteins
      a. Regulate an organism's activities.
      b. Can be classified as peptides because they are usually small
      b. Examples: 
          1) Insulin is a hormone secreted by the pancreas that sends signals to the cells to regulate the concentration of sugar in the blood steam for vertebrates.
 6. Receptor proteins
      a. Response of cell to stimuli from chemicals, neighboring cells, etc.
      b. Examples:
          1) Receptors that are attached to cell membranes detect signals from hormonal proteins.
 7. Contractile and motor proteins
      a. Involved in the movement of organelles
      b. Examples:
          1) Actin regulates the contraction of muscles
          2) Cilia is responsible for the movement of organelles
 8. Defensive proteins
      a. Protect against foreign substances in the body.
      b. Examples:
          1) Antibodies.
 9. Motor Proteins
      a. Convert chemical energy to mechanical energy to facilitate movement
      b. Examples:
           1) Actin and myosin are the proteins within muscles that help in movement.
           2) Microtubules help move organelles within the cell, and chromosomes 

Amino acids are organic molecules that contain a carboxyl and amino groups attaching an alpha carbon as the center. The alpha carbon also contains other various groups symbolized by R and a hydrogen. The R groups are usually called the side chain and they differ for each amino acid.

Amnio acids includes the following:

 1.  Glycine (Gly or G)
      - Nonpolar
      - Smallest R group with only a hydrogen atom
      - evolutionary conserves because most other R group cannot fit into the small space
      - alpha carbon is achiral
 2.  Alanine (Ala or A)
      - Nonpolar/aliphatic
      - R group is a methyl
      - alpha carbon is chiral
 3.  Valine (Val or V)
      - Nonpolar/aliphatic
      - alpha carbon is chiral
 4.  Leucine (Leu or L)
      - Nonpolar/aliphatic
      - alpha carbon is chiral
 5.  Isoleucine (Ile or I)
      - Nonpolar/aliphatic
      - alpha carbon is chiral
 6.  Methionine (Met or M)
      - Nonpolar
      - Sulfur containing
      - alpha carbon is chiral
      - First amino acid of proteins
 7.  Phenylalanine (PHe or F)
      - Nonpolar
      - Aromatic
      - alpha carbon is chiral
 8.  Tryptophan (Trp or W)
      - Nonpolar
      - Aromatic
      - alpha carbon is chiral
 9.  Proline (Pro or P)
      - Nonpolar
      - Cyclic
      - alpha carbon is achiral
 10. Serine (Ser or S)
      - Polar
      - Hydroxy containing
      - alpha carbon is chiral
 11. Threonine (Thr or T)
      - Polar    
      - Hydroxy containing
      - alpha carbon is chiral
 12. Cysteine (Cys or C)
      - Polar
      - Thiol containing
      - alpha carbon is chiral
 13. Tyrosine (Tyr or Y)
      - Polar
      - Aromatic
      - Hydroxy containing
      - alpha carbon is chiral
 14. Asparagine (Asn or N)
      - Polar
      - Amide
      - alpha carbon is chiral
 15. Glutamine (Gln or Q)      
      - Polar
      - Amide
      - alpha carbon is chiral
 16. Aspartic acid (Asp or D)
      - Electrically charged (Acidic)
      - alpha carbon is chiral
 17. Glutamic acid (Glu or E)
      - Electrically charged (Acidic)
      - alpha carbon is chiral
 18. Lysine (Lys or K)
      - Electrically charged (Basic)
      - alpha carbon is chiral
 19. Arginine (Arg or R)
      - Electrically charged (Basic)
      - alpha carbon is chiral
 20. Histidine (His or H)
      - Electrically charged (Basic)
      - alpha carbon is chiral

A proteins consist of one or more polypeptides. Polypeptides consist of different level of structures.

 1) Primary structure
     - the unique sequence of an amino acid
 2) Secondary structure
     - hydrogen bonds interact between polypeptide backbones.
     - polypeptides can fold into structures such as alpha helix, beta pleated sheet, etc.
 3) Tertiary structure
     - hydrophobic interaction and disulfide bridges formed between side chains of polypeptides
 4) Quaternary structure
     - overall protein stuctures consist of two or more polypeptide chains that combine into one functional macromolecule.

Carbohydrates

Carbohydrates are the most abundant biological molecules found in organisms and are responsible for the storage and transport of energy, such as starch and glycogen, and structural components, like cellulose in plants or chitin in animals.

Carbohydrates are classes of macromolecules - large polymers built from monomers. Carbohydrates include both sugars and polymers of sugars. The simplest carbohydrates are monosaccharides, also known as simple sugars. Another group of carbohydrates are the disaccharides, which are created by joining together two monosaccharides by a covalent bond. Polysaccharides consist of many monosaccharides used as building blocks.

 1. Monosaccharides
     a. simple sugar with general formulas, (CH2O)n
     b. the most common monosaccharide is glucose
     c. the molecule contains a carbonyl group and many hydroxyl groups attached to the carbon atom
     d. a monosaccharide can either be an aldose (sugar with an aldehyde) or a ketose (sugar with a ketone) depending on the location of the carbonyl group
     e. monosaccharide can be categorized by the number of carbon on the chain starting with carbon 3
 2. Disaccharide
     a. two monosaccharides join together through a glycosidic linkage - a covalent bond that can be formed by a dehydration reaction
     b. examples:
          1) maltose - two molecules of glucose joined together
          2) sucrose - a molecule of glucose joined with a molecule of fructose
 3. Polysaccharides
     a. polysaccharides are polymers of multiple monosaccharides joined by glycosidic linkages
     b. some polysaccharides are storage materials and some are structural

Storage polysaccharides

 1. storage polysaccharides store sugar for later use
 2. usually joined by 1-4 linkages of alpha glucose monomers
 3. usually form a helical shape
 4. starch
     a. a storage polysaccharide found  in plants that is joined by glucose monomers
     b. the simplest form of starch is unbranched amylose
     c. a more complex form of starch is branched amylopectin
 5. glycogen
     a. a storage polysaccharide in animals that is joined by glucose monomers
     b. an amylopectin-like polymer but contains more branches
     c. mainly stored in the liver and muscle cells

Structural polysaccharides

 1. structural polysaccharides usually give protection to the cell as a form of membrane
 2. usually joined by 1-4 linkages of beta glucose monomers
 3. usually form beta sheets
 4. cellulose
     a. major components of cell walls in plants
     b. cellulose is unbranched
     c. hydroxyl groups on the glucose are able to interact with other hydroxyl groups on other molecules to form hydrogen bonds
     d. cellulose molecules are grouped into units called microfibrils. 
     e. humans are unable to digest cellulose

Lipids

Lipids are the most diverse group of biochemicals. Structurally, the lipids' primary function is to be part of biological membranes, such as the cell membrane, or as a source of energy in organisms. Another major class of lipids that is produced in cells are steroids, such as cholesterol.

Lipids are grouped together by their hydrophobic behavior - meaning they mix poorly with water. Lipid consist mostly hydrocarbons.

Different kinds of lipid include the following:

 1. Fats
     a. fats consist of a glycerol and three fatty acids, which are constructed from hydrocarbons and the carbon at one end containing a carboxyl group.
     b. the fatty acids are joined to the glycerol by an ester linkage.
     c. saturated fats are fats that do not contain any double bond on the hydrocarbon chain
     d. unsaturated fats are fats that contain double bonds on different positions of the hydrocarbon chain
          1) most double bonds in the hydrocarbon chain are cis- double bonds
          2) trans fats contain trans-double bonds
 2. Phospholipids
     a. phospholipids consist of a glycerol, two fatty acids and a phosphate group attached to an alcohol group
     b. phospholipids consist of a hydrophilic head and a hydrophobic tail
     c. phospholipids make up the lipid bilayer in the cell membranes
          1) hydrophobic tails dislike water and try to get away from water by getting as close as possible to other hydrophobic tails
          2) hydrophilic heads like contact with water and act as protection for the hydrophobic tails in the bilayer
 3. Steroids
     a. steroids consist of four fused ring skeletons
     b. different steroids consist of the steroid skeleton and other chemical groups attached to it
     c. example
          1) cholesterol
               a) common component of animal cell membranes
 4. Glycolipid
     a. glycolipids consist of a fatty acid unit and a sugar unit
     b. derive from sphingosine
     c. simplest glycolipid is cerebroside
     d. a more complex glycolipid is ganglioside

References

Slonczewski, Joan L. Microbiology. 2nd ed. New York, 2009.

General Information

Anabolic reactions are those that require energy to occur. Cells can couple anabolic reactions together with catabolic ones to form an efficient energy cycle; the catabolic reactions transform chemical fuels into cellular energy, which is then used to initiate the energy-requiring anabolic reactions.

Anabolism can be thought of as a set of metabolic processes, in which the synthesis of complex molecules is initiated by energy released through catabolism. These complex molecules are produced through a systematic process from small and simple precursors. For example, an anabolic reaction can begin with relatively simple precursor molecules (created previously by catabolic reactions) and end with fairly complex products, such as sugar, certain lipids, or even DNA, which has an extremely complex physical structure. The increased complexity of the products of anabolic reactions also means that they are more energy-rich than their simple precursors.

Anabolic reactions are divergent processes. That is, relatively few types of raw materials are used to synthesize a wide variety of end products. This results in an increase in cellular size or complexity—or both. Anabolic processes are responsible for cell differentiation and increases in body size. Bone mineralization and muscle mass are attributed to these processes. Anabolic processes produce peptides, proteins, polysaccharides, lipids, and nucleic acids. These molecules comprise all the materials of living cells, such as membranesand chromosomes, as well as the specialized products of specific types of cells, such as enzymes, antibodies, hormones, and neurotransmitters.

ATP provides the energy needed for anabolism to take place. ATP is a high energy molecule that couples anabolism by the release of free energy. This energy does not come through the breakage of phosphate bonds; instead it is released from the hydration of the phosphate group.The chemical reaction where ATP changes to ADP supplies energy for this metabolic process. Anabolism is the opposite of catabolism, for example, synthesizing glucose is an anabolic process whereas the breaking down of glucose is a catabolic process. The Gibbs free energy for the synthesis of glucose is positive, meaning that the reaction is not spontaneous and will not go to completion in any time frame. However, when coupled with ATP, this reaction becomes more thermodynamically favorable, as is the case with many other endothermic reactions in the body. Anabolism and catabolism must be regulated in a way that does not allow the two processes to occur simultaneously. Each process has its own set of hormones that switch these processes on and off. Anabolic hormones include growth hormone, testosterone, and estrogen. Catabolic hormones include adrenaline, cortisol, and glucagon.

There is a need for cells to separate the metabolism process into anabolic and catabolic pathways. Anabolism requires the input of energy, which can be described as an "uphill" (energy intake) process. Catabolism is a "downhill" process which energy is released as the organism had used up energy. At certain points in the anabolic pathway, the cell must put more energy into a reaction than is released during catabolism. Such anabolic steps require a different series of reactions than are used at this point during catabolism.

Stages of anabolism

• There are three basic stages of anabolism.

Stage 1 production of precursors such as amino acids, monosaccharides and nucleotides.

Stage 2 use energy from ATP to turn the precursors into reactive form.

Stage 3 the assembly of these activated precursors into complex molecules such as proteins, polysaccharides, lipids and nucleic acids.

Examples of Biochemical Anabolic Reactions

1) Polysaccharides: Polysaccharides serve as an example of anabolism because polysacchrides are derived from their subunits of simple monosaccharides. A simple example is the formation of glycogen. Glygocen is a polysaccharide that is composed of subunit glucose monosacchrides connected by glycosidic bonds.

2) Polypeptides: Polypeptides serve as an example of anabolism because polypeptides are derived from their subunits of simple peptides. Polypeptides such as hemoglobin are composed of four different proteins (essentially peptides) that come together and form a completely different protein. Peptides them selves are a result of anabolic reactions themselves due to the condensation reactions that simple amino acids go through to combine and thus form peptide chains.

3) Carbon Fixation: Carbon fixation is another example of anabolism because in photosynthetic organisms such as plants, cyanobacteria, algae and other photoautotrophic organisms fixate carbon dioxide into glycerate 3-phosphate which is then further converted into glucose. Beginning with photosynthesis, synthesis of carbohydrates from sunlight and carbon dioxide. This process uses ATP and NADPH produced by the photosynthetic reactions to convert CO2 into glycerate 3-phosphate. This shows anabolism because from a smaller simpler subunit of carbon dioxide, glycerate 3-phosphate is produced, a much larger and complex biochemical compound.

Examples of Anabolic hormones

Growth hormone:

Growth hormone is a protein-based peptide hormone that stimulates growth, cell reproduction, and regeneration in humans and other animals. Growth hormone is often used to treat children with growth disorders as well as adult growth hormone deficiency.

Insulin:

Insulin is a hormone that is essential for regulating fat and steroids metabolism in the body. It causes the uptake of glucose from the blood by cells in the liver, muscle, and fat tissue. Glucose is then stored as glycogen in the liver and muscle. Insulin also contributes to other body functions such as vascular compliance and cognition.

Testosterone:

Testosterone is a steroid hormone found in mammals, reptiles, birds, and other vertebrates. In mammals, testosterone is mainly secreted in the testes of males and the ovaries of females. However, small amounts of testosterone are also secreted by the adrenal glands. Testosterone is the principal male sex hormone.

Estradiol:

Estradiol is the predominant sex hormone present in females. However, it is also found in males and acts as an active metabolic product of testosterone. Estradiol has major impacts on the reproductive and sexual functions as well as other organs. [7] [Berg]

Side effects may occur when anabolic hormones (or steroids) are used in excessive amount. If used too much by men, it can cause a decline in testosterone secretion, testicular atrophy (wasting away of the testes), and even breast enlargement. For women, excess use of anabolic hormones may cause a decrease in estrogen secretion and the ability to ovulate, the growth of facial hair, as well as regression of the breasts.

References

http://science.jrank.org/pages/319/Anabolism.html#ixzz0UM1KSRXO

http://science.jrank.org/pages/319/Anabolism.html

Berg, Jeremy; Tymoczko, John L.; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. New York. 2007

Overview

Catabolism is the release of energy from a set of metabolic pathways which break down molecules into smaller units, including the breaking down and oxidizing of food molecules. An example would be proteins, nucleic acids, lipids and polysaccharides being broken down into smaller molecules like amino acids, nucleotides, fatty acids, and mono saccharides. By glycolysis, the glucose is broken down into two pyruvates which can be used for later mechanism (Krebs cycle) to produce energy. The oxidation of long-chain fatty acid to acetyl-CoA is a central energy-yielding pathway in many organisms. Its opposite process is anabolism, which combines small molecules into larger molecules. Energy that is released from catabolism will store as ATP within the cell. The cell will then use this source of energy for synthesizing cell components from simple precursors, for the mechanical work of contraction and motion, and for transport of substances across its membrane. Catabolism maintains the chemical energy needed in order to help the cell grow and develop. Some waste products caused by catabolism are carbon dioxide, urea, and lactic acid. Heat is also sometimes released as a by product because these are oxidation processes. Examples of catabolism are the citric acid cycle.

The energy cells contain is liberated through two distinct processes: glycolysis and cellular respiration

GLYCOLYSIS Glycolysis is a series of reactions that break down glucose into two smaller organic molecules.

Glycolytic pathway Glycolysis occurs in the cytoplasm, in the presence or absence of oxygen. The pathway has several steps to convert six-carbon glucose into two molecules of three-carbon pyruvates. The direct generation of ATP from ADP and P_i is known as substrate-level phosphorylation. NAD⁺ is reduced to NADH twice, so we start with two molecules of NAD⁺ and finish with two molecule of NADH.

The net reaction for glycolysis is: Glucose + 2 ADP + 2P_i + 2 NAD⁺ → 2 Pyruvate + 2 ATP + 2 NADH + 2H⁺ + 2H₂O

Most of the chemical energy extracted from the sum is still stored in pyruvate’s bonds. Pyruvate has two potential fates based on the character of the cell’s environment. In aerobic organisms, pyruvate undergoes further oxidation through the mitochondrial electron transport chain. In anaerobic organisms, pyruvate undergoes an oxygen-free process called fermentation. Some cells are obligate aerobes or anaerobes, meaning that they require that designated environment. Others are facultative; they refer one environment over the other but can survive in either.

Fermentation In the glycolysis reaction equation, NAD⁺ is a necessary reagent and must be present for glycolysis to occur. At the end of glycolysis, the coenzyme is present only in its reduced form, NADH. One way to regenerate NAD⁺ is through oxidation in the electron transport chain, but anaerobic organisms would not participate in that. So another method is used: fermentation. Fermentation reduces pyruvate to either ethanol or lactic acid.

Alcohol Fermentation This process occurs in yeast and some bacteria. Pyruvate is first decarboxylated to acetaldehyde, which is then reduced by NADH to ethanol, thereby regenerating the NAD⁺ Pyruvate (3C) → CO2 + Acetaldehyde (2C) Acetaldehyde + NADH + H⁺ → Ethanol (2C) + NAD⁺

Lactic acid Fermentation This process occurs in some fungi and bacteria, as well as in mammalian muscles when oxygen demand exceeds supply. Basically, many glucose molecules are put through glycolysis, yielding twice as many molecules of cellular respiration, so it builds up. Concurrently, NADH builds up, depleting cells’ supply of NAD⁺. To keep muscles working, pyruvate is reduced to lactic acid, and NADH is oxidized back to NAD⁺. Lactic acid decreases the local pH, which we fell as the burn and fatigue effects of strenuous exercise. Once oxygen supply catches up to demand, the lactic acid may be converted back to pyruvate in the process known as the Cori cycle. The amount of oxygen necessary to do this is known as the oxygen debt. Pyruvate (3C) + NADH + H⁺ → Lactic Acid + NAD⁺

CELLULAR RESPIRATION It is the most efficient means of glucose catabolism, generating approximately 36 to 38 ATP per molecule of glucose. Respiration is aptly named; it is an aerobic process using an electron transport chain, with oxygen being the final electron acceptor. There are three key phases: pyruvate decarboxylation, the citric acid cycle, and the electron transport chain. A productive way to keep track of these reactions will be the follow the carbon.

Pyruvate decarboxylation The first step in aerobic respiration is pyruvate decarboxylation. This step itself does not require oxygen, but it only occurs once the cell commits to aerobic respiration—and the commitment is made only in the presence of oxygen. Pyruvate is transported from the cytoplasm into the mitochondrial matrix, where it is decarboxylated. The remaining acetyl (2C) group is bound to a coenzyme A molecule to form acetyl-CoA. Once NAD⁺ is reduced to NADH per pyruvate; in other words, two NAD⁺ molecules are reduced per molecule of glucose. Acetyl-CoA is a key intermediate in the utilization of fat, protein, and other carbohydrate energy reserves.

2 Pryuvate (3C) + 2 CoA + 2 NAD⁺ → 2 NADH + 2 Acetyl-CoA (2C) + 2 CO₂ (1C)

Citric Acid Cycle The citric acid cycle starts with the combination of acetyl-CoA (2C) and oxaloactate (4C) to generate citrate (6C). Through a series of eight reactions, two CO2 molecules are released, and oxaloacetate is regenerated. The citric acid cycle does not directly generate much energy. Each turn of the cycle generates one ATP via substrate-level phosphorylation and a GTP intermediate, for a total of two pyruvates per glucose molecule. The value of the citric acid cycle is its ability to generate high-energy electrons that are carried by NADH and FADH2. For each molecule of acetyl-CoA that enters the cycle, three NADH and one FADH2 are produced by two to account for the fact that the cycle turns twice per molecule of glucose. These enzymes then transport the electrons to the electron transport chain on the inner mitochondrial membrane, where more ATP is produced via oxidative phosphorylation. At the end of the citric acid cycle, oxaloacetate is regenerated in anticipation of the next round.

The overall reaction is: 2 Acetyl-CoA + 6 NAD⁺ + 2 FAD + 2 GDP + 2 P_i + 4 H₂O → 4 CO2 + 6 NADH + 2 FADH₂ + 2 ATP + 4 H⁺ + 2 CoA Structural Biochemistry/Lipids/Fatty Acids Single-molecule techniques have recently become popular in the biophysical department in helping to discover or clarify and better understand some important biochemical properties such as protein-DNA interactions, protein folding, and the functions and capabilities of membrane proteins. Many of the single-molecule techniques were first revealed in the physics and biophysics department, and later were found to be of great assistance to research biological and biochemical molecules. In 1976, the technique called single ion-channel recordings was first discovered and that later became the gateway to recent techniques such as atomic force microscopy (AFM), optical and magnetic tweezers and single-molecule fluorescence spectroscopy. Many of these recent techniques have helped in fields such as protein folding, transcription, replication, translation, molecular motors, membrane proteins, and viral biology. These single-molecule methods have been able to provide information on problems that could not be solved before while also giving a more detailed view into subjects that have already been researched. Not only that, but single-molecule techniques have also been able to steer the biochemists away from the usual averaging of ensembles that results in the use of moles to a more detailed and unitary concept that involves the single molecules and particles.

Static vs. Dynamic Heterogeneity

A useful aspect of single-molecule techniques is that they give the distribution of values for a property instead of an average of the property which is averaged over a large molecule ensemble. This kind of feature allows for more complete and specific data that can speak more about the certain biomolecule that is being assessed than an average which gives a rather unspecific and broad view. The specificity of single-molecule methods also provides for information on the molecular heterogeneity, which is a fundamental aspect of complex biomolecules and their functions. Molecular heterogeneity can be classified as either static or dynamic. Static heterogeneity is when a collection of molecules that have several subpopulations are extremely stable and do not interconvert over the period of time that is being observed. An example of this is inactive molecules. When an ensemble method is being used, one is trying to determine the fraction of active molecules compared to the whole population. However, when using a single-molecule technique, the inactive molecule can be ignored or disregarded because they do not give an experimental signal. Therefore, static heterogeneity allows for the sole research of the species of interest (active molecule) because the inactive ones can be ignored and removed. On the other hand, dynamic heterogeneity is when a sample with subpopulations of molecules interconvert over the time period observed. An example of dynamic heterogeneity is an enzyme that interconverts between two catalytically active states, with each conformation yielding a different affinity for a substrate. When the interconversion of the enzyme is fast compared to the temporal resolution of the single-molecule technique that is being used, the value that comes out will show a weighted time-average mean of the affinity for each of the two states of the enzyme. On the other hand, if the interconversion is slower than the temporal resolution of the single-molecule method, then the observation of the interconversion between the two states can be directly seen. In addition, dynamic heterogeneity can also be observed in enzymes that operate on substrates in multiple steps, with each step having a different rate. With single-molecule analysis, the observation of each single step in the process of the enzyme can be carefully monitored as real time "movies" that show the kinetics of each step and the intermediate structures of the enzyme and substrate.

Kinetics of Single-Molecule Techniques

The type of information that is extracted from single-molecule techniques are also something that can be compared to the ensemble assays that are used frequently. One example is the kinetic information that is provided by the single-molecule methods. Kinetic observation of a reaction by single-molecule methods provide something called dwell times. These dwell times are of a single molecule at each of its certain states along the reaction pathway. Think of two structural states of a molecule called A and B that is being observed by a single-molecule observable. The single-molecule method with enough temporal resolution to look at the two states would be able to determine the dwell time for states A and B. Afterward, a frequency histogram can be plotted to provide information on how long each dwell time lasts. Dwell time distributions can be used to study enzyme-substrate interactions. This example can also be used to reveal the significant aspect of single-molecule technique in its method of identifying the structural states of a biological macromolecule.

Single-Molecule Methods

There is a wide variety of single-molecule techniques that are present right now. The commonly used methods are separated into two different attributes: force and fluorescence. These two groups were separated based upon the different time resolutions, span of observations, and different spatial resolutions. In the force-based detection methods, there are specific techniques called atomic force microscopy, optical tweezers, tethered partical motion, and magnetic tweezers. On the other hand, the fluorescence imaging includes confocal microscopy and total-internal-reflection fluorescence (TIRF).

Force-based Dectection

Atomic force microscope block diagram v2

Atomic Force Microscopy (AFM)

Atomic force microscopy was first used for topographical imaging of molecules on a flat atomic surface. This form of imaging is done by scanning an extremely sharp tip along the sample surface and then measuring the deflection from the tip by using a laser and a quadrant photodetector. Dry and wet solutions can both be observed under this imaging, although the wet solution's temporal resolution is far worse than the dry solution's. The setup of AFM usually has the a small lever (AFM cantilever) that has a sharp nanometer scale tip which is attached to one of the ends of a biomolecule. By using piezoelectric positioning, the surface which the biomolecule is adsorbed can then be scanned in the x,y,z coordinate system (3D spatial) to an angstrom scale resolution. Then, the position of the tip could be measured by deflecting a laser beam off the surface of the tip and onto the photodetector, which will be position sensitive. The lever will act as a linear spring and when the biomolecule is moved relative to the tip, the lever will flex and apply a force on the molecule. Because the lever is very rigid, the force applied can be great and is therefore very useful in measuring the structural properties of folded proteins and chemical bonds.

Optical Tweezers

Generic Optical Tweezers ru

For the method of optical tweezers, the setup involves a high-power infrared laser that is focused tightly because of a high-end microscope lens. The focused infrared laser beam would trap a micron-sized bead at its focal spot, which allows for the control of the position of the bead. Then, the trap acts like a spring, which would exert a higher force on the bead when it moves farther away from the laser axis. And by measuring the position of the bead in relation to the focal spot, the trap's applied force can be determined (this can be done with the quadrant photodetector, like that in AFM). There are two main modes of optical tweezers: constant force and constant position. When it is in the constant force mode, a constant trap force can be maintained by a feedback loop that will displace the optical trap or the sample coverslip surface, which will keep the position of the bead constant within the zone that it is trapped in. In the constant position mode, the trap position's center is held and the bead will experience a progressively growing force as it is being pulled out of the trap.

Magnetic Tweezers

The setup for magnetic tweezers has the small bead that is used in optical tweezers replaced by a small magnetic bead. The small magnetic bead will be controlled by a pair of magnets that is close to the sample. This setup allows for the tethered biomolecule to have a force applied on it and a rotation that will be easily imposed. Magnetic tweezers function in the constant force mode because the force is an important control parameter for the interaction of interest. In addition, compared to the optical tweezers, the magnetic tweezers cannot obtain the exquisite spatial sensitivity. However, it is easier to use and can be compatible with long term analysis of single biomolecules. Magnetic tweezers is usually used for studying the structural properties of DNA and protein-DNA transactions.

Tethered Particle Motion

This method is used for studying the interation of proteins and polymers. The main idea is one bead is bound to the surface while the other bead is attached to the end of the polymer. These are then put into an aqueous solution and it can be seen that the bead has restricted movement in which it moves in a Brownian motion. The position of this bead is then recorded by optical microscopes; the recordings can tell us information about the DNA during transcription. Different bead types have also improved ways of analyzing polymers. Firstly, metal beads are sometimes used due to their high intensity of gold light. Secondly, polystyrene beads are also used in conjunction with the optical tweezers because these beads are less intense than the metal ones.

Fluorescence-based Methods

Fluorescent biomolecules contain several properties that can be used by fluorescent microscopy and spectroscopy to look at their location, structure and dynamics. There are two main methods for single-molecule fluorescence: confocal microscopy for point detection and wide-field imaging for area detection. Confocal microscopy for point detection tries to collect fluorescence that is emitted by a diffraction-limited volume of ~1 femtoliter. On the other hand, for wide-field imaging, there is total-internal-reflection fluorescence (TIRF) microscopy. TIRF is capable of observing hundreds of molecules that are immobile on the surface for extended periods of time through using evanescent-wave excitation within a thin layer just above the surface and using a ultrasensitive camera above surface for imaging.

FIONA or fluorescence imaging with one nanometer accuracy, has enabled motion tracking with precision of 1 nanometer. This can be done with any single fluorophore or light-scattered particle (example: small gold particle) that is attached to the molecule of interest. With that set up, one can identify the presence of a molecule and track its movement on molecular tracks in vitro or its diffusion in vivo. This remarkable tracking ability has been able to reveal how molecular motors like kinesin and myosin move on their tracks. The combination of this precision in tracking and the founding of complicated methods of switching on and off the fluorescence of some probes allows for techniques such as PALM (photoactivated localization microscopy) and STORM (stochastic optical reconstruction microscopy). PALM and STORM have been characterized as "super-resolution microscopies" because of their ability to achieve spatial resolutions that are better than 50 nanometers.

With two probes attached to a molecule, there are several new capabilities that come forth. FRET (fluorescence resonance energy transfer) is distance dependent and can be exploited to measure nanometer distances and the change in distances within single molecules. For FRET, the first probe would act as the FRET donor while the second probe would be the FRET acceptor. The first probe would be fluorescent, and the second probe would quench the donor in a distance-dependent manner (it can also be fluorescent). Because of this setup, the movements that change the donor-acceptor's distance of separation would also change their fluorescence. These changes are used to observe the kinetics of conformational changes or molecular association/disassociation. However, single-molecule FRET is also affected by relative orientation and rotational freedom of the fluorophores, which can be mistaken for conformational changes. This difficulty can be examined by using MFD (multi-parameter fluorescence), which is a method that can tell the fluorescence properties of single molecules such as fluorescence intensity, anisotropy, and lifetime at different wavelengths. Another method that can help the complication is ALEX (alternating laser excitation), which uses two lasers to measure FRET and relative probe stoichiometries.

Combination of Force and Fluorescence Approaches

The combination of force and fluorescence techniques is being pursued right now by single-molecule groups because of it will allow for simultaneous manipulation and visualization of single molecules as they interact and react. This combination works because the force and fluorescence techniques are highly complementary due to the fact that force-based methods can achieve timescales of 50-100ms while fluorescence-based can be many times faster and are not restricted by an applied force. In addition, whereas force-based methods provide information more on the global structural and mechanical rearrangements in biomolecules, fluorescence-based methods reveals the local conformational changes. This complementary ability of the two types of single-molecule method can lead to great discoveries in the future.

References

Kapanidis, A.N. and Strick, T., 2009, Biology, one molecule at a time, Trends in Biochemical Sciences, v. 34, p. 234-243.

Chemistry is the study of composition, structure and properties of matter. This includes studying changes that are observed when undergoing various chemical reactions or understanding the basic structure of complex macromolecules. It provides an understanding for atoms, molecules, crystals which make up everything in this world, while incorporating the concepts of energy and entropy.

Chemical Foundations of Biochemistry

3d Lewis representation of hydrogen, oxygen, nitrogen and carbon to use as organic building blocks

Chemistry is always interconnected with the other sciences, such as astronomy, physics, material science, and biology. In terms of biochemistry, it strives to find chemical explanations to biological form and function. Chemistry shows that all forms of life share common origins through the similarity of various chemical pathways between different organisms. For example, all living cells have the common molecules of amino acids, carbohydrates, lipids, and nucleic acids which all perform the same functions in every cell.

Nature is predominantly made of four elements: Hydrogen (H), Oxygen (O), Nitrogen (N), and Carbon (C). Together they make up over 99% of all living cells. These elements are the lightest atoms that are capable of making up to four stable and strong bonds. There are other elements that are also essential to living organisms such as Sodium (Na), Potassium (K), Calcium (Ca), and Sulfur (S). These are called trace elements and usually are needed to help specific proteins to function.

Organic Chemistry as the Backbone of Biochemistry

Organic chemistry is the study of the structures, properties, synthesis, and reactions of chemical compounds consisting of mainly carbon and hydrogen atoms. It is the backbone science that explains many of the interactions that occur in a cell at the most basic level. The majority of biochemistry, from a chemical point of view, involves the interactions between organic molecules that exist in organisms and how they are utilized in a versatile fashion at a cellular level. The arrangement of bonds within a molecule determines its overall shape, conformation, and the molecules ability to perform specific functions within a cell. Organic compounds that are of particular interest in biochemistry are carbohydrates, nucleic acids, and proteins.

Carbon is unique from the other essential elements for it has the capability to form four stable, covalent bonds and stable double and triple bonds. Carbon can bond with other carbons to create long chains with or without branches or cyclic structures allowing the formation of an endless array of molecules with different shapes, sizes, and compositions. Replacing hydrogen with other elements or functional groups gives the organic molecule different properties such as polarity.

Importance of Chemical Structure

Biochemistry chemistry is stereospecific meaning that the 3-D orientation of a molecules matters and only specific arrangements will interact properly. Stereoisomers are molecules that have the same chemical bonds but have a different spatial arrangement. They may look very similar but they cannot be the same molecule with out breaking any of the covalent bonds.

Human hands are mirror images of each other just like enantiomers.

There are many places where there can be varying spatial configuration. A molecule can have a chiral center which means that an atom (usually carbon) that is attached to four different substituents. These four can arrange themselves any way they want to around the carbon but if the arrangements are not exactly the same then they are stereoisomers. One example is if the molecule is a mirror image of itself. No matter how you rotate the molecule it will never superimpose on to its original form. This specific type of stereospecificity is called enantiomers. If they are not mirror images of each other they are called diastereomers.

Another way is to have a double bond. Double bonds, unlike single bonds, are stiff and do not rotate freely. The four substituents that surround the bond can form either trans, where the two highest ranked substituents are opposite sides of the bond, or cis, where the highest ranked are on the same side of the bond, relationships. This is called either geometric isomers or cis-trans isomers.

While enantiomers have almost the same chemical properties only the specific orientation will have the desired effect in a biological system. In nature, chiral compounds are usually found in only one orientation. For example, in carbohydrates the second to last carbon always has a "D" orientation and amino acid's chiral carbon is always in an "L" orientation. Enzymes can discriminate between enantiomers therefore is essential to have the right configuration.

Water and its Effects on Structure

The field of biochemistry also encompasses the special properties of water. Because all living organisms have cells that interact with aqueous solutions, the properties of water are fundamental to the environment in which biochemical reactions occur. The structure of a molecule determines whether it is hydrophilic or hydrophobic. Its interaction in water determines how these molecules function in a cell, one example being the lipid bi-layer in the cell membrane and how its chemical properties determines its function in the cell (in this case acting as a barrier in a cell). Water also has unique features such as high boiling and melting points. Its liquid form is also more dense than its solid form which allows marine life. An important feature of water in regards to biochemistry is its ability to form hydrogen bonds.

Introduction

Atoms form bonds by gaining, losing, or sharing electrons. Typically, they seek to achieve the electron configuration of a noble gas. Bonding occurs because it lowers the energy of a system and makes the atom more stable. Atoms also a unit of matter, the smallest unit of an element, consisting of a dense, central, positively charged nucleus surrounded by a system of electrons and an equal in number to the number of nuclear protons. The entire structure having an approximate diameter of 10-8 centimeter and characteristically remaining undivided in chemical reactions except for limited removal, transfer, or exchange of certain electrons.

An atom cannot be broken down further without changing the chemical nature of the substance. For example, if you have 1 ton, 1 gram or 1 atom of oxygen, all of these units have the same properties. We can break down the atom of oxygen into smaller particles, however, when we do the atom looses its chemical properties. For example, if you have 100 watches, or one watch, they all behave like watches and tell time. You can dismantle one of the watches: take the back off, take the batteries out, peer inside and pull things out. However, now the watch is no longer behaves like a watch.

Atoms are made up of 3 types of particles: electrons, protons and neutrons. Each of these particle has different properties. Electrons are tiny, very light particles that have a negative electrical charge (-). Protons are much larger and heavier than electrons and have the opposite charge. So, protons have a positive charge (+). Neutrons are large and heavy like protons, however, neutrons have no electrical charge. Anyway, each atom is made up of a combination of these particles. The proton and electron stay together just like the two magnets (the opposite electrical charges attract each other).

The electron that is constantly spinning around the center of the atom (called the nucleus). The centrigugal force of the spinning electron keeps the two particles from coming into contact with each other much as the earth's rotation keeps it from plunging into the sun. In an electrically neutral atom, the positively charged protons are always balanced by an equal number of negatively charged electrons. As we have seen, hydrogen is the simplest atom with only one proton and one electron. Helium is the 2nd simplest atom. It has two protons in its nucleus and two electrons spinning around the nucleus. With helium though, we have to introduce another particle. Because the two protons in the nucleus have the same charge on them, they would tend to repel each other, and the nucleus would fall apart. To keep the nucleus from pushing apart, helium has two neutrons in its nucleus. Neutrons have no electrical charge on them and they act as a sort of nuclear glue, holding the protons, and thus the nucleus together.

As you add more electrons, protons and neutrons, then the size of the atom increases. We can measure an atom's size in two ways: 1) By using the atomic number (Z) or using the atomic mass (A, also known as the mass number). The atomic number describes the number of protons in an atom. For hydrogen, the atomic number (Z) is equal to 1. For helium Z = 2. Since the number of protons equals the number of electrons in the neutral atom, Z also tells you the number of electrons in the atom. The atomic mass tells you the number of protons plus neutrons in an atom. Therefore, the atomic mass, A, of hydrogen is 1. For helium A = 4.

Electrically neutral atoms are the atoms that have no positive or negative charge on them. Atoms, however, can have electrical charges. Some atoms can either gain or lose electrons (the number of protons never changes in an atom). If an atom gains electrons, then the atom becomes negatively charged. If the atom loses electrons, then the atom becomes positively charged (because the number of positively charged protons will exceed the number of electrons). An atom that carries an electrical charge is called an ion.

While the number of protons for a given atom never changes, the number of neutrons can change. Two atoms with different numbers of neutrons are called isotopes. For example, an isotope of hydrogen exists in which the atom contains 1 neutron (commonly called deuterium). Since the atomic mass is the number of protons plus neutrons, two isotopes of an element will have different atomic masses (however the atomic number, Z, will remain the same).

The quantum mechanical basis for the formation of chemical bonds is an overlap in the probability densities of two or more wavefunctions; this means that electrons have a non-zero chance of being found in a region of space that is forbidden by classical physics.

Ionic Bonds

An ionic bond is the transfer of electrons between a metal and a nonmetal. An example of an ionic bond is that between sodium and chlorine atoms. The sodium atom transfers its lone electron in the 3s state to the chlorine atom. After the electron transfer, the sodium atom bears a +1 charge while the chlorine atom now bears a -1 charge. With this transfer of electrons, the sodium now has the electron configuration of the noble gas neon, while chlorine now has that of the noble gas argon.

Ionic bonding will occur only if the overall energy change for the reaction is favorable – when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond. The low electronegativity of metals and high electronegativity of non-metals means that the energy change of the reaction is most favorable when metals lose electrons and non-metals gain electrons.

Pure ionic bonding is not known to exist. All ionic compounds have a degree of covalent bonding, which means, ionic bond could be consider as a special type of covalent bond. The larger the difference in electronegativity between two atoms, the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.

Covalent Bonds

Covalent bonds are another type of chemical bond used to achieve a noble gas configuration, or an octet of electrons. Covalent bonds are formed between nonmetals, usually from the Boron, Carbon, Nitrogen, Oxygen, and Halogen families. Metals are rarely involved in covalent bonds. Each covalent bond consists of two electrons, one usually from each atom involved in the bond. The atoms form enough covalent bonds that when the electrons in the bonds are added with the valence electrons, they will have an octet. The key difference between ionic and covalent bonds lies in how the electrons are distributed between the two atoms. In ionic bonds, the electrons are transferred from one atom to the other, giving the atoms effective +1 and -1 charges. However, in covalent bonds, the valence electrons from both of the two atoms are shared between two atoms. Thus, neither atom is given a full positive or negative charge. Instead, the electrons shared between the two atoms - whether it be 2, 4, or 6 electrons - varies from molecule to molecule.

There are two types of covalent bonds: pure covalent bonds and polar covalent bonds. Pure covalent bonds exist when there is no difference between the two atoms sharing the electrons. The electronegativity of the two atoms is identical. Because the electronegativity values do not differ, they pull the electrons that are being shared between them with the same force. Thus, the electrons are shared equally and none of the atoms bears a partial positive or negative charge. An example of a pure covalent bond is a Cl-Cl or a Br-Br bond. Pure covalent bonds rarely exist for bonds that are not between identical atoms. Another example would be the covalent bonds between the carbons in long alkane chains.

Polar covalent bonds are those that exist between atoms of different electronegativities. The electrons in the bond are still being shared, but not equally between the two atoms. Though the exact ratio of the electron density that each atom bears cannot be determined easily, it is very easy to determine which atom pulls more electron density towards itself. The more electronegative atom will pull the shared electrons more, causing it to now bear a slightly negative charge. Because charge has to be conserved, the less electronegative atom must now bear a slight positive charge, equal in magnitude to the negative charge. As an example, consider a bond between carbon and chlorine. Chlorine is much more electronegative than carbon, thus it pulls more of the electrons towards itself. This gives the chlorine a slightly negative charge and the carbon a slightly positive charge. If the difference between the two atoms is so great causing one of the two atoms to posess a lot of the electron density, the bond becomes increasingly ionic and less covalent. For this reason, though H-Cl is considered a covalent bond, it is classified as a very strong acid, meaning it dissociates completely. Because the electronegativity difference is so vast, the chlorine molecule pulls all the electron density towards itself, thereby dissociating into H+ and Cl- ions in the presence in water.

However, it is important to note that a molecule that contains polar bonds can be nonpolar. For example, take the molecule carbon tetrachloride. This molecule has four polar C-Cl bonds. However, due to the orientation of the polar bonds, they cancel out and the molecule as a whole is nonpolar.

Hydrogen Bonds

A hydrogen bond is a bond created by the dipole-dipole interaction of a hydrogen atom and an electronegative atom such as an oxygen or nitrogen atom due to dipole dipole interactions. A common example of this is water where the electronegativity of the oxygen allows it to have a slight negative charge while the two hydrogen atoms have a slight positive charge. The negative charge on the oxygen forms a weak bond with the slight positive charge of another water molecule's hydrogen. This type of bonding is also present in organic fluorine compounds between C and F groups. This force is weaker than covalent bond and ionic bonds, but stronger than Van der Waals interactions.

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C), or most of the solutions that use water as the solvent. This is because of the strong hydrogen bond, as opposed to other group 16 hydrides. Intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids.

Role of Noncovalent Interactions in Macromolecules

In macromolecules such as proteins, DNA, and RNA, noncovalent interactions are essential. Noncovalent interactions include hydrogen, ionic, hydrophobic, and Van Der Waals bonding. These interactions are described in more specificity in the list that follows this group. When compared to covalent bonds, noncovalent bonds are weak and continuously form and break bonds. However, when several noncovalent bonds are formed, there is a net increase in bond strength. Their combined participation in a macromolecule makes a difference (i.e. substrate binding to enzyme and the lipid bilayer's role in transport). With several hydrogen bonds, ionic, and hydrophobic interactions existent at the same time, it is unlikely that these several weak interactions will break the substrate and enzyme without external energy. This property is the reason why enzymes have specific catalytic power. Protein folding and the unique properties and structures of proteins also depend on these noncovalent interactions.

Metallic Bonding

Metallic bonding is the bonding between metal and metal. The bonding involves electron pooling and metallic bonding. Since metal atoms are larger and can easily lose their valence electrons, the electrons on their outer shell can be pooled and be distributed evenly with other metal atoms. Differently from the covalent bonding, electrons in the metallic bonding are delocalized which means that the electrons can move freely through the metal.

Metallic Bonding

1) Metallic bonding in sodium: Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. Even a metal like sodium (melting point 97.8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the periodic table. Sodium has the electronic structure 1s22s22p63s1. When sodium atoms come together, the electron in the 3s atomic orbital of one sodium atom shares space with the corresponding electron on a neighbouring atom to form a molecular orbital. This is the same sort of way that a covalent bond is formed. The difference, however, is that each sodium atom is being touched by eight other sodium atoms, the sharing occurs between the central atom and the 3s orbitals on all of the eight other atoms. And each of these eight is in turn being touched by eight sodium atoms, which in turn are touched by eight atoms and so on, until you have taken in all the atoms in that lump of sodium. All of the 3s orbitals on all of the atoms overlap to give a vast number of molecular orbitals, which extend over the whole piece of metal. There have to be huge numbers of molecular orbitals, because each orbital can only hold two electrons. The electrons can move freely within these molecular orbitals. Each electron becomes detached from its parent atom. The electrons are said to be delocalized. The metal is held together by the strong forces of attraction between the positive nuclei and the delocalized electrons.

2) Metallic bonding in magnesium: Magnesium has the outer electronic structure 3s2. Both of these electrons become delocalized, so the "sea" has twice the electron density as it does in sodium. The remaining ions also have twice the charge and so there will be more attraction between ions and sea. Each magnesium atom has 12 protons in the nucleus compared with sodium's 11. In both cases, the nucleus is screened from the delocalized electrons by the same number of inner electrons,the 10 electrons in the 1s2 2s2 2p6 orbitals. That means that there will be a net pull from the magnesium nucleus of 2+, but only 1+ from the sodium nucleus. So not only will there be a greater number of delocalized electrons in magnesium, but there will also be a greater attraction for them from the magnesium nuclei. Magnesium atoms also have a slightly smaller radius than sodium atoms, and so the delocalized electrons are closer to the nuclei. Each magnesium atom also has twelve near neighbors rather than sodium's eight. Both of these factors increase the strength of the bond.

3)Metallic bonding in transition elements: Transition metals tend to have particularly high melting points and boiling points. The reason is because they can involve in the 3d electrons in the delocalization as well as the 4s. The more electrons can involve, the stronger the attraction.

4) The metallic bond in molten metals: In a molten metal, the metallic bond is still present, although the ordered structure has been broken down. The metallic bond is not fully broken until the metal boils. That means that boiling point is actually a better guide for the strength of the metallic bond than melting point. In melting, the bond is loosened and is not broken.

Relative Strength of Chemical Bonds and Intermolecular Forces

All the bonds known to chemistry can be described the various dissociation energies needed to break the bonds. It is these dissociation energies that rank the strength of the various bonds found in chemistry. Ionic bonds, being the strongest bonds, have a dissociation energy of >400 kcal/mol. Covalent bonds where are the second strongest bonds have a dissociation energy of about 400 kcal/mol. Hydrogen bonds, dipole-dipole, and london (van der waals) dispersion bonds are in a sub category of bonds labeled intermolecular forces. These forces are significantly weaker than ionic and covalent bonds because of their nature of being interactive forces between compounds rather than physical bonds. Hydrogen bonds are the strongest of said forces and overall third strongest in all bond interactions with a dissociation energy of 12-16 kcal/mol. Dipole Dipole interactions are the second strongest intermolecular force but the fourth strongest bond interaction with a dissociation energy of 0.5-2 kcal/mol. Lastly comes the london vander waal forces that the weakest interaction in chemical bonding with a dissociation energy of <1 kcal/mol.

References

Berg, Jeremy M. Biochemistry. 6th ed. W.H. Freeman, 2007.

Introduction

This is a common form of a covalent bond where the hydrogens both share one electron each

Covalent bonds are chemical bonds that are formed by sharing valence electrons between adjacent atoms. This type of bonding is mostly seen in interactions of non-metals. Covalent bonds allow elements the ability to form multiple bonds with other molecules and atoms - a fundamental necessity for the creation of macromolecules. In the covalent bond, as the distance between the nuclei decreases, each nucleus starts to attract the other atom's electron, which lowers the potential energy of the system. Anyway, when the attraction increases, the repulsions between the nuclei and between the electrons increase as well. In covalent bonding, each atom achieves a full outer (valence) level of electrons. Each atom in a covalent bond counts the shared electrons as belonging entirely to itself. Most covalent substances have low electrical conductivity because electrons are localized and ions are absent. Overall, the atoms in a covalent bond vibrate, and the energy of these vibrations can be studied with the IR spectroscopy.

Octet Rule

A general rule to follow when looking at covalent bonding is the octet rule, also known as the noble gas configuration. An atom participating in covalent bonding must (with few exceptions) follow the octet rule, which states that an atom must have eight electrons around it. These electrons can be shared or unshared. The two atoms do not need to share their electrons equally; an electron pair can be donated from one atom instead of each atom donating one electron. A periodic table can be used to determine the number of valence electrons an atom. The general rule is that all atoms will be stable if they can have eight electrons around them. Therefore different atoms can share their unpaired electrons with other atoms with unpaired electrons to gain an octet.

There are quite a few exceptions to this rule. Two very important ones are Hydrogen (H) and Helium (He). These atoms do not have octets and only need a total of two electrons to be stable. This is because hydrogen and helium only contain a 1s electron shell, which can only hold two electrons. Other exceptions occur when the total number of electrons in a molecule or between two molecules is an odd number. These molecules tend to be very reactive. Also, atoms past the second row on the periodic table can have more than eight electrons surrounding them. For example, in phosphorus pentafluoride (PF₅) the phosphorus is bonded to 10 electrons, and in sulfur hexafluoride (SF₆) the sulfur atom is bonded to 12 electrons. Molecules can also be electron deficient, meaning there are not enough available electrons to complete full octets around all the atoms in the molecule. An example of an electron deficient molecule is boron trichloride (BCl₃). In this molecule, the boron atom is only bonded to three electron pairs, while the chlorides are surrounded by full octets.

Types of Covalent Bonds

Multiple covalent bonds can be formed between atoms, which are stronger than single bonds and have higher bond energy and shorter bond lengths. The bond order is used to determine the number of pairs of electrons in a covalent bond. When a molecule has double and single covalent bonds, it can have different chemical forms of equal energy as resonance structures, which has more stability and the bond is the average of the double and single covalent bond. The characteristics of a covalent bond can also be effected by the two atoms it joins.

Single Bonds

An electron dot diagram of a covalent bond between chlorine and hydrogen

Single bonds are one of the weaker types of covalent bonds. Single covalent bonds are also called sigma bonds. These are made when only two electrons are shared. This leads to an overlap of the orbitals and a merging of the electron density clouds. Single bonds tend to be very flexible allowing atoms to rotate around the bond. An example of a single bond is a carbon-carbon (C-C) covalent bond has a bond length of 1.54 A and bond energy of 356 kJ/mol.

Note that the properties of a single bond depends not only on the two atoms that is bonds but also on the atoms surrounding those atoms. Sigma bonds have no nodal planes.

Some covalent single bonds will also have double bonds properties, which are shorter, rigid and non-rotated. One example is peptide bond in proteins which connect each amino acid together to form polypeptide. The peptide-bond is 1.32 A which is shorter than 1.54 A (C-C). The energy that needs to break the peptide bond is much higher than the single bond and this non-rotated single bond contributes the planar property in the polypeptide chain, which also makes the peptide bond more stable than the normal single bond. The double bond properties are contributed by the resonance structure of the pepetide bond.

Double bonds

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Double bonds occur when a covalent bond consists of four shared electrons. A double covalent bond contains a sigma bond and a pi bond. Pi-bonds apply to the overlapping p-orbitals. The orbitals can only overlap in a side-by-side arrangement leading to one nodal plane on the internuclear axis. A single covalent bond only contains a sigma bond. Double bonds tend to be shorter than their single bond equivalents and stronger. Double bonds also create electron density around the bond. Unlike single bonds, double bonds are not flexible and the two adjoining atoms cannot rotate about the bond.

Triple Bonds

C₂H₂.

A triple covalent bond contains one sigma bond and two pi bonds where six electrons are being shared. These bonds are stronger than double bonds and shorter. They are more rigid than double bonds and have a larger electron density. The most common triple bonds are on carbons like C2H2. The skeletal form to draw a triple bond is three straight lines connecting the two atoms.

Polar Covalent Bonds

Covalent bonds can be polar or non-polar depending on the electro-negativity value of the atoms bonded together. If there is a very large difference between the two atoms' electro-negativity values, a polar covalent bond is formed. The atoms do not need to possess the same electro-negativity values, or be of the same element, but they need to be relatively close in their values. If the electro-negativity values are closer, the co-valency between the atoms will be stronger. An exception to this rule is when a molecule possesses symmetry. When the overall dipole moment is zero, such as linear molecule of CO2, the molecule is considered non-polar. The more electro-negative atom will attract the electrons, making itself have a partial negative charge and giving the other atom a partial positive charge. These partial negative and positive charges are what account for the dipole-dipole, dipole-induce dipole, and induced dipole-induced dipole interaction. This attraction-to-repulsion stability is what gives the covalent bonds stability. In addition to the electro-negativity differences between atoms, covalent bonding depends on the angles of adjacent atoms relative to each other.

Typical accepted values for determination of type of bonds:

Difference in electronegativity - X < 0.5 - Non-polar covalent bond

Difference in electronegativity - 0.5 ≤ X ≤ 1.9 - Polar covalent bond

Difference in electronegativity - 1.9 < X - Ionic bond

Specific Types Of Covalent Bonds

Disulfide Bonds In Chemical interactions, certain compounds can react to create a disulfide bond, which is a type of covalent bond that is usually derived by the coupling of two thiols (-S-H). These interactions can also be called SS-bonds or disulfide bridges, with the connectivity of these interactions mainly being R-S-S-R .

Role in Protein Folding Disulfide bonds can play a vital role in the tertiary structure of proteins in the effect they have on protein folding and stability. These disulfide bonds between proteins usually are formed between the thiol groups of cysteine residues. The other amino acid group in which sulfur appears is methionine, which cannot form disulfide bonds.

Formation of a Disulfide Bond.

Disulfide bonds help to stabilize the tertiary structure of a protein molecule in several ways, for example, The disulfide bonds destabilize the unfolded form of a protein by lowering its overall entropy, or state or chaos. Also, when the disulfide bonds link two segments of the protein chain, this increases the effective local concentration of protein residues and lowers the effects of water in a that specific region. Since water molecules are known to attack amide-amide bonds, lowering the effects of water in these disulfide bond- regions helps to stabilize a protein.

Covalent Bond: Bond Length and Bond Energy

The bond energy (BE) is the energy required for the attraction or breakage between the atoms. Since it is the energy needed to break the attraction between the atoms, the bond energy is endothermic and positive. However, the energy required for the formation of the bond is exothermic and a negative value. The bond length is the distance between the nuclei of two covalent bonded atoms. It can be calculated based on the total radii of the bonded atoms. As a result, the bond length increases when the covalent radius increases. And the shorter the bond length, the higher bond energy will be needed to break the attraction between the atoms because shorter distance between the atoms means the bond will be stronger and harder to break. On the other hand, the longer the bond length is, the lower bond energy is needed to break a weaker bond. One can use bond energy to determine the ΔH_rxn. In a reaction, when two atoms react with each other to form the product of different atoms, there are two types of bond energy. One is the energy required for the reactant to be broken and the other one is the energy required for products to be formed. As a result, the difference between the two bond energy is the enthalpy or the work of the reaction. ΔH⁰_rxn= ΔH_{reactant bonds broken} + ΔH_{product bonds formed}

References

Silberberg, Martin S. Chemistry "The Molecular Nature of Matter and Change." Fifth Edition.

Background Information

A noncovalent bond is a type of chemical bond that typically bond between macromolecules. They do not involve sharing a pair of electrons. Noncovalent bonds are used to bond large molecules such as proteins and nucleic acids. Noncovalent bonds are weaker than covalent bonds but they are crucial for biochemical processes such as the formation of double helix. There are four commonly mentioned fundamental noncovalent bond types. They include electrostatic interactions, hydrogen bonds, van der Waals interactions, and hydrophobic interactions. Each type differs in geometry, strength, and specificity.

Types of Noncovalent Bonds

Electrostatic interactions

The energy of an electrostatic interaction is given by Coulomb’s Law: E = kq1q2 / Dr2 where E is the energy, q1 and q2 are the charges of two atoms, r is the distance between the two atoms, D is the dielectric constant, and k is a proportionality constant. k = 1389 for energy that’s measured in kilojoules/mol or k = 332 for energy that’s measured in kcal/mol. A charged group on a molecule can attract an oppositely charged group from another molecule. By contrast, an attractive interaction has a negative energy. The dielectric constant is important for the medium.

Hydrogen bonds

A hydrogen bond is the interaction of a hydrogen atom with an electronegative atom. The electronegative atom can be nitrogen, oxygen, fluorine that comes from another chemical group. Hydrogen bonds are responsible for specific base-pair formation in the DNA double helix. The hydrogen atom in a hydrogen bond is shared by two electronegative atoms such as nitrogen or oxygen. A hydrogen-bond donor is the group that includes an electronegative atom where the hydrogen atom is more tightly bound to and a hydrogen-bond acceptor is when an electronegative atom is less tightly bound to the hydrogen atom. The electronegative atom where the hydrogen atom covalently bonds can pull electron density away from the hydrogen atom creating a positive electronegativity charge. The hydrogen atom can also interact with an atom that has negative electronegativity charge.

van der Waals interactions

Van der Waals interaction is the distribution of electronic charge around an atom that fluctuates with time. It is the sum of the attractive or repulsive forces between molecules. The charge distribution is not perfectly symmetric. The attraction increases as two atoms come closer to each other, until they are separated by the van der Waals contact distance. When the distance of the energy is shorter than the van der Waals contact distance, a very strong repulsive force becomes dominant.

Hydrophobic interactions

Hydrophobic interaction is the physical property of a molecule that is repelled from a mass of water. They are also called hydrophobic exclusions. It is the tendency of hydrocarbons to form intermolecular aggregates in an aqueous medium, and analogous intramolecular interactions.

Ion Induced Dipole

An ion induced dipole is a noncovalent bond interaction that results when the approach of an ion induced a dipole in an atom or in a nonpolar molecule by distributing the arrangement of electrons in the nonpolar species.

Dipole Induced Dipole

A dipole-induced dipole interactions results a nonpolar molecule has it's electron density shifted by a charged molecule. If a negatively charged species described by its electron density (for example, in water electron density is centralized on the oxygen atom) is brought near a species with an evenly distributed charged molecule, the electrons go towards the side where the positive species is, and away from the opposite side, creating an artificial dipole.

http://upload.wikimedia.org/wikipedia/commons/thumb/a/a6/Dipole_interactions.png/320px-Dipole_interactions.png

References

Berg, Jeremy M., John L. Tymoczko, and Lubert Stryer. Biochemistry. 6th ed. New York: W. H. Freeman and, 2006. Print. A hydrogen bond is formed by a dipole-dipole force between an electronegative atom (the hydrogen acceptor) and a hydrogen atom that attaches covalently with another electronegative atom (the hydrogen donor) of the same molecule or of a different molecule. Only nitrogen, oxygen, and fluorine atoms can interact with hydrogen to form a hydrogen bond donor; this is different than a hydrogen covalent bond. The hydrogen bond acceptor however can be any atom which is in a polar bond, is electronegative, and has a lone pair. In a hydrogen bond, the lone pair electrons on oxygen, nitrogen, or fluorine interact with the partial positive hydrogen that is covalently bonded to one of those atoms. The hydrogen atom in a hydrogen bond is shared by two electronegative atoms such as oxygen or nitrogen.) Hydrogen bonds are responsible for specific base-pair formation in the DNA double helix and a major factor to the stability of the DNA double helix structure. A hydrogen-bond donor includes the hydrogen atom and the atom to which it is most tightly linked with. The hydrogen-bond also play a very important roles in proteins' structure because it stabalizes the secondary, tertiary and quaternary structure of proteins which formed by alpha helix, beta sheets, turns and loops. The hydrogen-bond connected the amino acides between different polypeptide chains in proteins structure. The hydrogen-bond acceptor is the atom that is less tightly linked to the hydrogen atom.

Hydrogen bonds are fundamentally electrostatic interactions and are much weaker than covalent bonds. They are, however, the strongest kind of dipole-dipole interaction. The electronegative atom to which the hydrogen atom is bonded with pulls electron density away from the hydrogen atom, developing a partial positive charge. Therefore, the hydrogen atom can then interact with a partial negatively charged atom through an electrostatic interaction.

Hydrogen bond with ammonia

Components

Hydrogen bonding is a form of electrostatic interaction between a hydrogen atom bonded to two electronegative atoms; one of which is the hydrogen-bond donor that has a stronger bond between itself and the hydrogen. These electronegative atoms are nitrogen, oxygen, and fluorine; this electronegative atom pulls electron density away from the hydrogen atom, giving it a partially positive charge. This partial positive charge is attracted to the partial negative charge of the hydrogen bond acceptor (an electron density rich atom). The chemical bond formed between the hydrogen-bond donor, hydrogen atom, and hydrogen-bond acceptor has a straight, linear structure.

H-bonding

Hydrogen bonding (H-bond) is a non-covalent type of bonding between molecules or within them, intermolecularly or intramolecularly. This type of bonding is much weaker and much longer than the covalent bond and ionic bonds, but it is stronger than a van der waals interaction. It also carries some features of covalent bonding: direct and straight. In other words, H-bond donor and H-bond acceptor lie along the straight line. In order to form an H-bond, an H-bond donor and H-bond acceptor are required. The H-bond donor is the molecule that has a hydrogen atom bonded to a highly electronegative, small atom with available valence (N, F, and O follow the above description the best because they are very electronegative, making H, which is covalently attached to them, very positive). The H-O, H-N, and H-F bonds are extremely polar; as a result, the electron density is easily withdrawn from the hydrogen atom towards the electronegative atom. The partially positive hydrogen in one molecule attracts to partially negative lone pair of the electronegative atom on the other molecule and H-bond forms as a result of such an interaction. All the hydrogen bonds vary in strength

Other important facts about hydrogen bonding are as follows. The small sizes of nitrogen, oxygen, and fluorine are essential to H bonding for two reasons. One is that it makes those atoms electronegative that their covalently bonded H is highly positive. Other reason is that it allows the lone pair on the other oxygen, nitrogen, or fluorine to come close to the H. Also, hydrogen bonding has a profound impact in many systems. Hydrogen bonding is also involves in the action of many enzymes [The Molecular Nature of Matter and Change].

Properties of Water Due to Hydrogen Bonding

Ammonia, water, and hydrogen fluoride all have higher boiling points than other similar molecules, which is due to hydrogen bonds. Bonds between hydrogen and these strongly electronegative atoms are very polar, with a partial positive charge on hydrogen. This partially positive hydrogen is strongly attracted to the partially negative oxygen on the adjacent molecule. In general, boiling points rise with the increase molecular weight, both because the additional mass requires higher temperature for rapid movement of the molecules and because heavier molecules have a greater London forces. Water's freezing point is also much higher than other similar molecules. An unusual feature is that it decreases in density when it freezes. The tetrahedral structure around each oxygen atom, with two regular bonds to hydrogen and two to other molecules. This requires a great amount of space between the ice molecules. Clathrates are molecules trapped in holes of solid, like ice, that is theorized to be able to be used as anesthesia.

Hydrogen bond and physical properties

Hydrogen bonding has a significant influence on a molecule's boiling points. The boiling point usually increases with the increase of the molar mass. However, molecules that are involved in intermolecular H-bonding bonding have much higher boiling points in comparison with the molecules of the same molar mass that are not involved in H-bonding. This is because the unusually strong H-bonding forces allow for stronger interaction between water molecules and therefore creating a stronger bond and higher boiling point. In addition, H-bonding is responsible for many unusual proprieties of water, such as its high boiling point, melting point, heat of vaporization, high dielectric constant, surface tension, capillary action etc.

Hydrogen bonding can occur between hydrogen and four other elements. Oxygen(most common), Fluorine, Nitrogen and Carbon. Carbon is the special case in that it only really interacts in hydrogen bonding when it is bound to very electronegative elements such as Fluorine and Chlorine.

Hydrogen bonding is an important component of the three major macromolecules in biochemistry such as proteins, nucleic acids, and carbohydrates. The H-bonding is responsible for the structure and properties of proteins(enzymes). Hydrogen bonding is applicable in these biomolecules because of functional groups present. Some such are the carboxylic acid, alcohol or even amine groups. These provide either an hydrogen, oxygen or nitrogen for possible hydrogen bonds.

Hydrogen bond in proteins

Hydrogen bonding within the "green fluorescent protein" 1RRX.

As previously mentioned, hydrogen bond can be intermolecular (ex. the bonding of water molecules) as well as intramolecular (ex. the bonding of protein and DNA). The secondary structure of protein forms as a result of H-bonding between amino acids. For example, an α-helix is a rod-like secondary structure that forms as a result of H-bonding between the carboxyl group of (i) amino acid to the amino group of (i+4) amino acid. The turn (loop family) is a secondary structure which forms as a result of H-bond between carboxyl group of (i) amino acid and amino group of (i+3) amino acid. The β-sheet is a secondary structure which forms as a result of H-bonding between two or more β-strands. An anti-parallel β-strands forms hydrogen bonds that are straight due to the carbonyl group and the amino group being directly aligned, while a parallel β-strand forms hydrogen bonds that are slightly weaker in comparison to the anti-parallel because the carbonyl group and the amino group don't align perfectly, which forms a longer and weaker hydrogen bond.

The solubility of proteins in water is dependent on the ability to form hydrogen bonds with the protein surface. Proteins that have a greater hydrophilic surface content are generally more capable of forming hydrogen bonds with the surrounding water. The alteration of salt concentration of the solution, as is performed in salting out/in, creates a shielding effect that reduces the ability to form an H-bond with the hydrogens in water. The protein precipitation method of salting out utilizes this concept in protein fractionation.

Hydrogen bonding in water

Hydrogen bonding in water

The simplest example of a hydrogen bond can be found in water molecules. A water molecule consists of one oxygen atom attached to two hydrogen atoms. A hydrogen bond can be formed between two molecules of water. In the case of liquid water where there are many water molecules present, each water molecule could potentially hydrogen bond with up to 4 other molecules (2 through its 2 hydrogen atoms with each hydrogen bonding to another oxygen and another 2 through its 2 lone pairs on the oxygen that can hydrogen bond to 2 other hydrogen atoms).

Although water has a low molecular mass, it has an unusually high boiling point. This property can be attributed to the large number of hydrogen bonds that exist within the water. Since these bonds are difficult to break, water’s melting point, viscosity, and boiling point are relatively high in comparison to other liquids that are similar but lack the hydrogen bonding. Water contains substantially more hydrogen bonds (up to 4) relative to certain other liquids that also have hydrogen bonding. An example would be ammonia in which the nitrogen only has one lone pair but 3 hydrogen atoms and thus only capable of forming up to 2 hydrogen bonds.

Hydrogen bonding can also explain why the density of ice is less than the density of liquid water. In water's liquid form, the hydrogen bonding that keeps the molecules close together is constantly being broken and remade repeatedly at room temperature. But as the water turns into ice, the hydrogen bonding causes the water molecules to form a rigid, lattice structure, which causes large gaps between the molecules, resulting in it's smaller density yet larger volume.

Hydrogen bonding also accounts for water's high surface tension. The large availability of hydrogen bonding between water molecules (4 hydrogen bonds to one water molecule) proves how well they can stick to each other, forming a strong and stretchy surface. Common examples from which this characteristic can be observed include a cup filled slightly over the top without spilling over, or small organisms that are able to stay on top of the water without breaking its surface.

Water has a different number of hydrogen bonds depending upon the temperature. It is estimated that at 0oC each water molecule has an average of 3.69 hydrogen bonds, while at 25oC it has an average of 3.59 hydrogen bonds, and at 100oC it has an average of 3.24 bonds. The decreasing hydrogen bonds with an increase in temperature can be attributed due to the increase of molecular motion.

Hydrogen bonding in DNA

DNA contains four bases: Guanine, Cytosine, Adenine, and Thymine. The complementary base pairs of guanine with cytosine and adenine with thymine connect to one another using hydrogen bonds. These hydrogen bonds between complementary nucleotides are what keeps the two strands of a DNA helix together. Each base can also form hydrogen bonds with the external environment such as with water. Although these internal and external hydrogen bonds are fairly weak, the consolidated power of all the millions of hydrogen bonds in DNA make it a stable molecule. Also, the hydrogen bonds on the phosphate groups on each nucleotide interact inducing two strands of DNA to conform to a helical structure.

The base pairing in the DNA (one purine and one pyrimidine base) can be explained in more details. In addition to holding the DNA strands together, the hydrogen bonding between the complementary bases also sequester the bases in the interior of the double helix. Therefore, the hydrogen bonding between the bases reinforces the hydrophobic effects that stabilize the DNA. The hydrophobic bases are again kept in the inside of the helix, whereas the polar exterior is touching the solvent water. The hydrogen bonding is a weak molecular force, but it is an additive effect that stabilizes the DNA molecule. The bases are precisely held by hydrogen bonding with the energy of 1 to 5 kcal/ mol (4 to 21 kJ/mol).

The hydrogen bonding in the DNA bases of one purine (guanine and adenine) and one pyrimidine (cytosine and thymine) creates a similar shape. The pairing of guanine and cytosine shape and structure is very similar to that of the pairing of adenine and thymine. Cytosine and Guanine are held together by three hydrogen bonds. The pairing of adenine and thymine share two hydrogen bonds, thus the bond is slightly weaker and slightly longer.

G-C hydrogen bonding

A-T hydrogen bonding

References

Silberberg, Martin S. Chemistry "The Molecular Nature of Matter and Change." Fifth Edition. 2009.

Berg, Jeremy; Tymoczko, John; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. 2007. (8) http://en.wikibooks.org/wiki/Structural_Biochemistry/Chemical_Bonding/Dipole_interaction http://en.wikibooks.org/wiki/Structural_Biochemistry/Nucleic_Acid/DNA http://en.wikibooks.org/wiki/Structural_Biochemistry/Chemical_Bonding/Covalent_bonds http://en.wikibooks.org/wiki/Structural_Biochemistry/Proteins http://en.wikibooks.org/wiki/Structural_Biochemistry/Nucleic_Acid http://en.wikibooks.org/wiki/Structural_Biochemistry/Carbohydrates http://en.wikibooks.org/wiki/Structural_Biochemistry/Proteins/Structures http://en.wikibooks.org/wiki/Structural_Biochemistry/Proteins/Amino_Acids http://en.wikibooks.org/wiki/Structural_Biochemistry/Nucleic_Acid/Nitrogenous_Bases/Purines/Guanine http://en.wikibooks.org/wiki/Structural_Biochemistry/Nucleic_Acid/Nitrogenous_Bases/Purines/Cytosine http://en.wikibooks.org/wiki/Structural_Biochemistry/Nucleic_Acid/Nitrogenous_Bases/Purines/Adenine http://en.wikibooks.org/wiki/Structural_Biochemistry/Nucleic_Acid/Nitrogenous_Bases/Purines/Thymine http://en.wikibooks.org/wiki/Structural_Biochemistry/Water

Definition

The tendency of nonpolar molecules in a polar solvent (usually water) to interact with one another is called the hydrophobic effect. The interactions between the nonpolar molecules are called hydrophobic interactions. The relative hydrophobicity of amino acid residues is defined by a system known as hydrophobicity scales.

General Information

Hydrophobic oil immiscible in water

The interactions between nonpolar molecules and water molecules are not as favorable as interactions amongst just the water molecules, due to the inability of nonpolar molecules to form hydrogen bonding or electrostatic interactions. When nonpolar molecules are introduced to the water molecules, the water molecules will initially surround the nonpolar molecules, forming a "cages" around the molecules. However, the tendency of nonpolar molecules to associate with one another will draw the nonpolar molecules together, forming a nonpolar aggregate.

Based on the second law of thermodynamics, the total entropy of the system plus its surrounding must always be increasing. Therefore, it is favorable for the nonpolar molecules to associate without the interference of water. The water molecules that initially "caged" the nonpolar molecules are released from the nonpolar molecules' surfaces, creating an increase in entropy in the surrounding. The favorable release of water molecules from nonpolar surfaces is responsible for phenomenon of the hydrophobic effect.

Problem

Hydrophobic interactions can also be seen in the clustering of amphipathic/amphiphillic molecules such as phospholipids into bilayers and micelles. The hydrophobic areas of amphipathic molecules cluster together to avoid the ordered "cage" of water molecules that would surround them and orient the hydrophillic ends as a shield-like outer structure that interacts amicably with the polar water molecules. Micelles occur when a spherical fatty acids structure is formed with a hydrophobic core and hydrophillic outer shell. Bilayers can be commonly seen in cell membranes with hydrophillic outer (outside the cell) and inner (inside the cell) linings has hydrophobic (inside the membrane) center. The Lipid bilayer is a more favored formation in nature due to the micelle formation may contain bulky fatty acids causing hindrance in its formation.

Lipid Bilayer(1) and Micelle (2)

Electric Properties of Plasma Membrane

Most cell membranes are electrically polarized, such that the inside is negative [typically -60 millivolts (mV)]. Membrane potential plays a key role in transport, energy conversion, and excitability. For example, membrane transport. Some molecules can pass through cell membranes because they dissolve in the lipid bilayer. Additionally, most animal cells contain a high concentration of K1 and a low concentration of Na1 relative to the external medium. These ionic gradients are generated by a specific transport system, an enzyme that is called the Na1–K1 pump or the Na1–K1 ATPase. The hydrolysis of ATP by the pump provides the energy needed for the active transport of Na1 out of the cell and K1 into the cell, generating the gradients. The pump is called the Na1–K1 ATPase because the hydrolysis of ATP takes place only when Na+ and K+ are present. This ATPase, like all such enzymes, requires Mg2+

Solution

When two nonpolar molecules come together, structured water molecules are released allowing them to interact freely with bulky water. The release of water from such cages is favorable. The result is that non-polar molecules show an increased tendency to associate with one another in water compared with others - less polar and less self-associating solvents. This tendency is called the hydrophobic effect and the associated interactions are called hydrophobic interaction.

The release from the cage-like clathrates is more favorable because it increases the entropy of the system.

References

http://en.wikibooks.org/wiki/Structural_Biochemistry/Second_law http://en.wikibooks.org/wiki/Structural_Biochemistry/Water http://en.wikibooks.org/wiki/Structural_Biochemistry/Lipids/Biological_Membranes Van Der Waals interactions (also known as London Dispersion forces) are weak attractions that occur between molecules in close proximity to each other. The basis of these interactions is that the distribution of electronic charge around an atom fluctuates with time. As two atoms come closer to each other, this attraction increases until they are separated by the van der Waals contact distance. When two molecules are too close to each other, the potential energy due to repulsion is very high, which means that it is unstable. There is repulsion even though the molecules are neutral, because there is an electron cloud surrounding each molecule. When these molecules get too close to each other, repulsion between the molecules occur. As the molecules move farther apart from each other, the potential energy due to repulsion decreases. This force was named after Johannes Diderik van der Waals, a Dutch physicist who studied them extensively.

An example of interaction energy-distance relation graph obtained from argon dimer. Because of London dispersion forces, the interaction energy greatly increases when the distance between two atoms is smaller than 3.8A

Van der Waals radius

Vander Waals attraction

Two molecules can interact by Van der Waals forces when they are at a certain distance apart. The molecules are stabilized by Van der Waals interaction at the Van der Waals contact distance because the potential energy of the system at this distance is at its lowest. In the potential energy diagram shown to the right, the minimum potential energy point corresponds to the Van der Waals contact distance.

As the distance between the molecules increases, the weakly bonded molecules lose their stability and are no longer affected by the Van der Waals forces due to their large distance apart. When the distance between the two molecules decreases, however, the stabilization is decreased as well due to the electrostatic repulsion between the molecules. This level of repulsion is felt more drastically and more intensely in Van der Waals interactions than in ionic interactions, where the level of repulsion is felt more gradually.

Energy in Van der Waals interaction

Energies associated with van der Waals interactions are quite small. Usually, they are about 2 to 4 kJ/mol per atom pair. When the surfaces of two large molecules come together, a large number of atoms are in Van der Waals contact, and the net effect, summed over many atoms pairs can be substantial. Macromolecules such as proteins and DNA contain numerous sites of potential van der Waals interactions that the cumulative effect of these small binding forces can be enormous; hence the most stable structure for macromolecules is that where weak interactions are maximized.

Van der waals interaction in proteins structure

In addition to hydrogen-bonds and disulfide bonds, protein structure can also be stabilized by Van der waals interactions. In the coiled-coil protein, there are Heptad repeat which form by the side chain interaction between each alpha helix; hepad-repeat is repeated in every 7th residues. If these repeating residues are hydrophobic, such as leucine, van der waals interaction will be formed to stabilize this protein structure.

References

Berg, Jeremy; Tymoczko, John; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. 2007. (8)

Nature of Dipole Bonding

The nature of dipole bonding begins when atoms differ in their electronegativity, which quantifies an individual atom's ability to attract electrons to itself. A classic example of an electronegative atom is fluorine.

When a covalent bond forms between two atoms, the electrons will be distributed between the two atoms unequally; the more electronegative atoms will have the larger electron density. This unequal sharing of electrons creates a charge separation, and the molecule under inspection will develop partial charges where the electronegative atom will develop a partial negative charge and its adjacent atom will develop a partial positive charge. The molecule is then said to be polarized due to this charge separation. When molecules exhibit this charge separation, there is a pseudo-electrostatic force between the partial charges of molecules.

The key to dipole bonding is charge separation within a molecule.

Permanent dipole-dipole interactions

Dipole moment

When the covalent bonds in a molecule are polarized so that one portion of the molecule experiences a positive charge and the other portion of the molecule experiences a negative charge. This separation of opposite charges creates an electric dipole. Depending on the orientation of two dipoles, molecules can be attracted to each other as the partial negative charge is attracted to the partial positive charge.

• Ex. Carbon Monoxide (CO)

The permanent and induced dipole interaction of nonpolar molecules bonded to nonpolar molecules in polar molecules, which forms more stable structure.

Besides, when the polar molecules lie near one another, as in liquids and solids, their partial charges act as tiny electric fields that orient them and give rise to dipole-dipole forces. So, the dipole-dipole forces is the positive pole of one molecule that attracts the negative pole of another. The greater the dipole moment, the greater the dipole-dipole forces between the molecules are.

Induced dipole interactions

Induced dipoles occur due to the fact that electrons are in constant motion within a molecule. For nonpolar molecules, the averaged charge distribution within a molecule would suggest that there is no permanent dipole. However, instantaneous charge distributions indicate the presence of a transient dipole. Imagine that if you take a photograph of a molecule, you can see an electron and a nucleus in a fixed arrangement. This fixed arrangement implies that there is a dipole moment, if only for a brief moment. We call this type of instantaneous dipole an induced dipole. Brief partial negative charges will be attracted to brief partial positive charges.

Induced dipoles are typical of nonpolar molecules such as fatty acid chains, aliphatic chains, and aromatic hydrocarbons.

Hydrophobic interactions

Hydrophobic effect

When nonpolar molecules clump together when surrounded by polar molecules.

• The nonpolar molecules are not held together by bonds which is why it is called an instantaneous dipole-induced dipole interaction. This phenomenon is due to the reduction of the number of water molecules surrounding the hydrophobic portion of molecules.

clathrate

The structure that is formed when hydrophobic interactions take place (when the hydrophilic molecules line the outside of the hydrophobic molecules). This forces the hydrophobic molecules to become highly ordered

• Example: When water molecules surround lipids

NaCl forming an Ionic Bond.png

An Ionic Bond is a specific type of chemical bond formed between a "metal" and a "nonmetal." "Metals" involved in ionic bonds are usually the Alkali and Alkaline-Earth metals - also known as the first two columns on the period table - as well as several transition state metals. The non-metals usually involved in the ionic bonds are the halogens.

The goal of forming chemical bonds is to have an octet. "Octet" means that an element has eight electrons. Another way of putting this phenomenon is to state that each element wants to have the electron configuration of a noble gas. For the non-metals, the goal is to achieve the electron configuration of the noble gas in the same row. For the metals, the goal is to have the electron configuration of the noble gas in the row directly preceding it. For example, chlorine wants the electron configuration of Argon while sodium wants that of Neon. Ionic bonds help achieve the octet because the metal effectively transfers its valence electrons to the nonmetal. In this way, both metal and nonmetal achieve noble-gas electron configurations.

However, pure ionic bonds do not exist. There is a level of covalent-bond character in each ionic bond. As a general rule of thumb, the larger the electronegativity difference between the metal and nonmetal, the more ionic the bond, and therefore, the less covalent the bond. Electronegativity is the ability to draw electron density while in the presence of another atom. The more electronegative the atom, the greater its ability to pull electron density towards itself.

Electrostatic interactions are also known as Charge-charge interactions and Ionic interactions. An electrostatic attraction exists when there are closely packed ions of opposite charges. An electrostatic repulsion is present between different ions that have the same charge.

Coulomb's Law

The force between two point charges can be calculated by Coulomb’s law,

F_1on2 = F_2on1 = kq₁q₂/r²

In other words, the bond energy is directly proportional to the charges of the two atoms and inversely proportional to the square of the distance between the two atoms. F has the unit N (Newtons), r is the distance between the center of the two-point charges in meter, q₁ and q₂ are the charges (in C, Coulombs) of each atom respectively. k is the constant, approximately equal $9.00x10^{9}$ N•m²•C⁻². If the force is negative (F < 0), it represents the existence of attraction since the only possible way to have F smaller than zero is if the sign of q₁ and q₂ are different, or opposite, meaning the sign of charges for these two atoms are opposite. In other words, if the force is positive (F > 0), it means the two charged atoms repel each other, due to the same sign of charges (both positive or negative charges.)

Ionic bond energy

The ionic bond energy between two-points charges can be calculated using the formula of electrical potential energy,

U_ele= Kq₁q₂/r²

In other words, the bond energy is directly proportional to the charges of the two atoms and inversely proportional to the distance (bond length in microscopic level) between the two charged atoms. U_ele has the unit J (Joules), r is the distance between the ion centers in nanometers, and q₁ and q₂ are the numerical ion charges. K is the constant, approximately equal $2.31x10$⁻¹⁹ J•nm. It is also equal to kQ² x 10⁻⁹ (Conversion to nanomemter, nm) or Q²/4$\pi \epsilon _{0}$ x 10⁻⁹ where k is approximately equal to $9.00x10^{9}$ Nm²C⁻² and Q is equal to the charge of a proton, $1.60x10$⁻¹⁹C; $\epsilon _{0}$ approximately equal to $8.85x10$⁻¹² C²N⁻¹•m⁻² or F•m⁻¹, representing the permittivity of free space.

Lattice Energy

Lattice energy (ΔH⁰_lattice) is the released energy from the interaction between ions. The positive sign of this energy indicates that the energy is needed for ions to form a solid. On the other hand, the negative sign of the energy indicates that the energy is needed for ionic solids to be separated into its gaseous ions. In addition, ionic interactions can be explained based on the lattice energy because the qualitative number of the energy indicates the ions' hardness, solubility, and melting point. Lattice energy can be predicted based on the effect of ionic size and ionic change. As one goes down the periodic table, the ionic radius increases. And as the radius increases, there would be a decrease in the electrostatic energy between the positively charged atom and negatively charged atom. As a result, there would be a decrease in the lattice energy. In addition, ionic charge can also determine the lattice energy. A greater amount of energy will be needed to bring a larger charged ions together than to form the smaller charged ions.

Introduction

A disulfide bond, also called an S-S bond, or disulfide bridge, is a covalent bond derived from two thiol groups. In biochemistry, the terminology R-S-S-R connectivity is commonly used to describe the overall linkages. The most common way of creating this bond is by the oxidation of sulfhydryl groups. (2 RSH → RS-SR + 2 H+ + 2 e-) This process of oxidation can produce stable protein dimers, polymers, or complexes, in which the sulfide bonds can help in protein folding. The process mostly occurs with the thiol groups in cysteine.

Formal depiction of disulfide bond formation as an oxidation.

Disulfide bonds can occur in two ways: intramolecularly and intermolecularly. Intermolecular disulfide bonds occur between polypeptide chains while intramolecular disulfide bonds occur within a polypeptide chain and are usually responsible for stabilizing tertiary structures of proteins. On the other hand, intermolecular disulfide bonds are attributed to stabilizing quaternary protein structures.

Disulfide Bonds in Proteins

Two cysteine residues can be linked by a disulfide bond to form cystine.

Disulfide bonds in protein membranes are found in both bacteria and eukaryotes. Extracellular proteins often have several disulfide bonds, whereas intracellular proteins usually lack them. In proteins, these bonds form between the thiol groups of two cysteine amino acids. Cross-linkage between multiple linear polypeptide chains is not uncommon in proteins. Most of the cross-linkages are from disulfide bonds formed by the oxidation of two cysteine amino acids. The result is a disulfide bond called cystine connecting the polypeptide chains. The cysteine amino acid group is the only amino acid capable of forming disulfide bonds, and thus can only do so with other cysteine groups. These bonds are responsible for the stabilizing the globular structure and are the strongest type of bond that a protein can possess and are one of the major forces responsible for holding proteins in their respective conformations, and therefore have an important role in protein folding and stability.

The typical bond dissociation energy of a disulfide bond ranks at 60 kcal/mole and has a bond length of 2.05 Å. Fairly low energy is required to produce rotations about the S-S bonds, thus these rotations are common. At dihedral angles near 90°, the bonds tend to be more stable. However, the bonds become significantly better oxidants at angles approaching 0° and 180°. Disulfide bonds have been identified in the protein folding in E. Coli. They are used in many processes, including DNA replication.

Disulfide Bonds in cyclic peptides

Most cyclic peptide bonds are formed between disulfide bonds. As a result, the denaturation of cyclic peptides can often be attributed to the stability of disulfide bonds. In the study with the peptide 1 (cyclo(1,4)-Cys-Gly-Phe-Cys-Gly-OH), where it was conducted in buffer solutions between pH 1-11 at 70 degrees C. It was found that the most stability came from pH ~ 3 and a Vshape between pH ~1-~5. As the pH goes from neutral to basic, degradation was found between Gly2-Phe3, which is due to the breaking of disulfide bonds.

Disulfide bonds hold proteins, such as this one, in its conformation

Making disulfides

Multiple ways to make disulfides

In the journal article “Multiple ways to make disulfides” by Neil J. Bulleid and Lars Ellgaard, they discuss how disulfides can be formed in the endoplasmic reticulum (ER), and the different enzymes that catalyze the pathways of formation. Disulfides increase the stability of the protein and also “regulate redox-dependent functions,” and over the years, our ideas of how disulfides form in proteins have drastically changed. [2]

Disulfides are created in the presence of enzymes in the protein disulfide isomerase (PDI) family. They act as a oxidizing agent, oxidizing the thiol group on a protein. If the protein's amino acid residues, specifically cysteines, are close to one another they will form a disulfide bond even if it is not properly folded. If a disulfide bond forms when the protein is not properly folded, they call this a non-native disulfide. This could be a misfolded protein, or it could be one of the intermediates before the protein folds into its native state. PDIs help non-native disulfides become native disulfides by acting as a catalyst to the isomerization process (they have to help brake the non-native disulfide bond so that the protein can finish folding properly before they can form the native disulfide bond) (Figure 1). [2]

Figure 1: http://imgur.com/sEvem

Ero1

Bulleid and Ellgaard studied an enzyme in yeast to understand how disulfides were formed de novo (Latin for “in the beginning”). The enzyme they studied was ER oxidoreduct in (Ero1p). Experiments showed that “Ero1p and the mammalian homologues ERO1(alpha) and ERO1(beta) are able to catalyze oxidation by coupling de novo disulfide formation to the reduction of oxygen to hydrogen peroxide (H2O2).” Ero1p was shown to oxidize PDI, which allowed PDI to exchange disulfides on the protein. Using knockout experiments they were able to show that while in yeast, knockout of Erop1 interrupted disulfide formation, in higher eukaryotes (for example, mice and humans) knockout of ERO1(beta) only caused misfolding in proinsulin and the double knockout of ERO1(alpha) and ERO1(beta) did not do much worse than just the knockout of ERO1(beta). In fact, after some time, “double knockout cells re-established normal ER redox conditions after a strong reductive challenge, albeit at a slower rate than in wild-type cells.” This tells us that ERO1 is not as necessary in higher eukaryotes as it is in yeast and implies that there are other pathways to forming disulfide bonds. [2]

PRDX4

Because hydrogen peroxide is produced when disulfides are formed via ERO1 catalysis, and H2O2 can cause damage biomolecules, Bulleid and Ellgaard believed there had to be other proteins in order to remove the H2O¬2¬. This is where peroxiredoxin (PRDX4) comes in. PRDX4 is a group of enzymes located in the ER that both removes H2O2 and also forms disulfides. In this process the peroxidatic cysteine in PRDX4 takes an oxygen from H2O2 to make water and a –SOH group, this then reacts with the adjacent –SH group to form a disulfide bond (Figure 2). This can now be exchanged with the –SH groups on some PDI proteins so it can then exchange with substrate proteins (Figure 1). [2]

Figure 2: http://imgur.com/oyzmc

References

1. He, HT. "Synthesis and chemical stability of a disulfide bond in a model cyclic pentapeptide: cyclo(1,4)-Cys-Gly-Phe-Cys-Gly-OH." The University of Kansas, Lawrence, Kansas 66047, USA.

2. Neil J. Bulleid, Lars Ellgaard, Multiple ways to make disulfides, Trends in Biochemical Sciences, Volume 36, Issue 9, September 2011, Pages 485-492, ISSN 0968-0004, 10.1016/j.tibs.2011.05.004.(http://www.sciencedirect.com/science/article/pii/S096800041100082X)

Overview

Researchers initially thought that enzyme ERO1, endoplasmic reticulum oxidoreductin 1, couples oxygen reduction to de novo formation of disulfides, but it was recently discovered that mammals that are deficient in this enzyme still survive and form disulfides. This suggests that there exist alternative pathways to forming disulfides. Discoveries have found that peroxiredoxin 4 is involved in peroxide removal and disulfide formation. Many other different pathways for disulfide formation in the mammalian ER include quiescin sulfhydryl oxidase, ER-localized protein disulfide isomerase peroxidases and vitamin K epoxide reductase. These various pathways for forming disulfides are regulated by glutathione.

Disulfide-bond

Disulfide formation in the endoplasmic reticulum

Many important secretory and cell-surface proteins like antibodies, plasma membrane receptors and channels, extracellular matrix proteins, and blood clotting factors contain disulfide bonds because disulfides heighten protein stability and control redox-dependent functions. Disulfides in the endoplasmic reticulum undergo a procedure catalyzed by membranes of the PDI, protein disulfide isomerase, family and then undergo co-translational translocation to the ER. As the polypeptide starts to fold, cysteine residues that come into close contact form disulfides, even if they do not form in the end product. These disulfides that do not end up folding into the final product are sources of problems for misfolded proteins but may also serve as intermediates in normal folding processes. In order to obtain the final correctly folded disulfide, incorrectly folded disulfides must be disintegrated in a reaction catalyzed by the PDI family. The PDI family thus plays a vital role in the course for the correct formation and reduction of disulfides in order for proper folding of proteins that enter the endoplasmic reticulum

The correspondence of a disulfide between the enzyme and substrate is needed to cause a catalytic reaction to form a disulfide. Since PDI family members each accommodate at least one thioredoxin domain, a CXXC motif alternates between dithiol and disulfide states at their active site. When a disulfide is transferred to the substrate protein, a reduction of the active site occurs. It is necessary that the active site be deoxidized in order for the enzyme to carry out further oxidation. It was initially thought that GSSG, glutathione disulfide, was relevant in how the active site is reoxidized. In Vitro experiments were conducted and showed that a GSH/GSSG ratio similar to that found in the endoplasmic reticulum would efficiently oxidize active site cysteines in PDI, thus transferring disulfides onto substrate proteins. These experiments didn’t show how disulfides were introduced de novo. With the revelation of ERO1p, an enzyme in yeast called endoplasmic oxidoreduction, it was found that this enzyme was necessary for the formation of disulfide. ERO1p plays an important role in oxidizing PDI instead of secreting proteins or low-molecular-weight molecules like GSH and catalyzing oxidation by coupling de novo disulfide formation to the reduction of oxygen to hydrogen peroxide. While securely regulated in order to prevent overproduction of reactive oxygen species, ERO1 showed how disulfides could be formed de novo and also identified the ultimate electron acceptor for the pathway.

Ero 1 is essential in yeast but not in higher eukaryotes

Knockout of the gene encoding ERO1p in various organisms showed different conclusions of its importance. When the gene encoding ERO1p was knocked out in yeast, it demonstrated its importance as an essential protein in yeast. Knockouts in higher eukaryotes showed different results. When knocked out in D. mlanogasterI, this led to a relatively mild phenotype with a certain problem in its folding of the cell-surface receptor Notch. Two ERO1 paralogs, pair of genes ERO1α and ERO1β, exist in mice and humans. When knocked out of ERO1 β, a defect in the folding of proinsulin occurred. ERO1α is thought to drive disulfide development in other tissues. It was found that a double knockout of ERO1α and ERO1β did not result in a more severe reaction than when ERO1 β is knocked out by itself. This showed that there existed an ERO1-independent pathway for the formation of disulfides in mammalian cells and double knockout cells helped re-establish normal endoplasmic redox conditions.

Potential pathways to generate disulfides de novo in the ER

PRDX4

Since H2O2 is also generated when a reaction is catalyzed by ERO1, this indicates that additional proteins might be present in the endoplasmic reticulum to remove this reactive oxygen species. Enzymes called peroxiredoxins metabolize H2O2 ensuing disulfide formation. PRDx4 is active in both H2O2 removal and disulfide formation. Studies show a new responsibility for PRDx4 in de novo disulfide formation. Since PRDX4 and several PDI family members within the ER and equivalent, this shows that the enzyme is an abundant endoplasmic resident protein. Under circumstances where PRDX4 was efficiently reduced by some PDI family members during incubation of the two proteins at equal concentrations, some PDI protein was oxidized. This showed that even though GSH was a reductant on its own, when a PDI is included, the efficiency of PRDx4 reduction is enhanced. This proposes that disulfide exchanged between PRDx4 and GSH to form GSSP depends on the presence of PDI. If PRDX4 were to be reduced by PDI family members, a rapid disulfide formation in secretory proteins could occur. These results are confirmed by in vitro evidence. Since PRDX4 enhances a temperature-sensitive mutant of ERO1, viability and disulfide formation in yeast occurs at non-permissive temperatures. However, overexpression of PDI family members can also lead to decrease or increase in the ability of PRDX4 to be reduced in the endoplasmic reticulum. In conclusion, by combining the ERO1 and PRDX4 pathways, every oxygen molecule that is reduced causes two disulfides to be introduced, therefore making the whole process more effective than using ERO1 alone.

GPX7 and GPX8

GPX7 and GPX8, homologous enzymes, belong to the family of thioredoxin GPX-like peroxidases that can also reduce H2O2. They are PDI peroxidases that couple the reduction of H2O2 to oxidation of certain PDI family members. In the presence of these GPXs, certain PDI family members are easily oxidized. With the presence of both GPX and PDI, oxidative refolding of a reduced model protein mediated by H2O2 proceeds faster. Physical associations between ERO1α and both GPXs in cells are shown by bimolecular fluorescence complementation. With the addition of GPX7, the in vitro rate of oxygen consumed by ERO1α increased, indicating a more efficient process in its presence. These biochemical results show an important role for GPX7 and GPX8 in disulfide formation.

QSOX

Erx2p, a sulfhydrl oxidase, when overexpressed could suppress the ero1-1 mutation. This different pathway shows that alternative proteins could potentially fulfill the essential function of ERO1 in yeast and other organisms. QSOX, a flavoprotein known to introduce disulfides into proteins in vitro, is similar to Erv2p. QSOX plays a part in catalyzing de novo disulfide formation by coupling disulfide oxidation to the reduction of oxygen in order to form H2O2. Unlike ERO1, which specially oxidizes only PDI family members, QSOX has a much broader substrate specificity that can introduce disulfides into protein substrates. However, PDI has the ability to greatly enhance native disulfide formation since QSOX cannot isomerize non-native disulfides. QSOX has the ability to complement Δero1 yeast strain when overexpressed. This indicates that QSOX is involved in disulfide formation when in vivo. When ERO1 and QSOX are both knocked down, a more severe phenotype occurs than when ERO1 is knocked down alone indicating that QSOX might provide some function when ERO1 is absent. QSOX currently remains a candidate for de novo disulfide formation independent of ERO1 due to its promiscuous substrate specificity and location in the secretory pathway.

VKOR

The VKOR enzyme exhibits another potential ERO1-independent pathway for disulfide formation . VKOR is a four transmembrane helix protein in the endoplasmic reticulum. VKOR functions by catalyzing the two steps in the reduction of vitamin K epoxide so as to generate vitamin K hydroquinone. When vitamin K epoxide is reduced, a CXXC motif in VKOR is then oxidized to form a disulfide bond. Members in VKOR family exchange this disulfide with thiredoxin-like oxidoreductases in order to oxidize substrate proteins. Since human VKOR does not contain a thioredoxin domain, PDI family members instead serve as VKOR substrates. WHen overexpression occurs with active-site CXXA mutants, VKOR can be trapped in a mixed-disulfide complex with PDI family members, mainly the transmembrane-bound TMX and TMX4. Results of these experiments show that VKOR sustaining disulfide formation is still unclear but proves a potential pathway.

Relative contribution and interplay between oxidative pathways

With all these potential pathways for disulfide formation, now their relative contribution has to be figured out. ERO1 pathway is an important pathway for disulfide formation under normal physiological conditions. TRDX4 hyperoxidation in cells expresses a deregulated form of ERO1 and shows that the enzyme is active in the ER and has the ability to produce hydrogen peroxide. Even though mammals do no necessarily need this pathway, if it didn’t exist, it would compromise disulfide formation. ERO1 is important in determining the redox status of PDI in cells and also forms mixed disulfides with PDI. This shows that ERO has the ability to oxidize PDI. If animals did not have ERO1 to survive and still formed disulfide-bonded proteins, this shows that other pathways exist. Their relative contributions still have to be explored.

PRDX4-dependent pathway doesn’t provide a significant contribution to disulfide formation. If it did, it would be predicted that the Prdx4 knockout mouse should have a severe phenotype and even though the Prdx4 knockout mouse is viable, it is sterile due to oxidative stress in the testis. PRDX4 pathway is then determined to be crucial for correct function of specific tissues though not essential for survival. If there were a higher presence of PRDX4 in higher eukaryotes, it would show mild phenotypes of ERO1 knockouts as an alternative source of hydrogen peroxide that is available in order to make the PRDX4 pathway work. More research is needed to determine which sources provide the H2O2 necessary to drive disulfide formation. The importance of the PRDX4 pathway needs to be determined, even if PRDX4 provides an alternative pathway to ERO1 for disulfide formation.

An endogenous level, enzymes GPx7 and GPx8 cannot substitute for PrdxIV in ERO1 knockout cells. These enzymes could have a similar function to PRDX4. Research shows that proteins of the peroxiredoxin family have a higher rate of reactivity towards H2O2than GPXs. Evidence shows that interaction between Ero1α with the GPXs in vivo might compensate for their relatively slower rates of reactivity. Further understanding of their characterization is needed of the enzymes in order to explain their roles in de novo disulfide formation. Current research shows that the cellular function of VKOR in disulfide formation remains essentially uncharacterized. In VKOR, the mixed disulfide trapped with PDI family members does suggest a role in the process though. The pathway of oxidation of PDIs by VKOR being coupled strictly to γ-carboxyglutamate formation is not important due to its low flux. This would not be the case if the hydroquinone could be reoxidized through an alternative electron acceptor. This could bring about a more important role in the pathway.

As shown in vitro experiments, direct oxidation by H2O2 of cysteines in PDI and substrate proteins to form disulfides could potentially occur. Although, these pathways are shown to not be important due to faster kinetics of substrate refolding when the reaction mixture contains H2O2 and PDI family members together with either of the PDI peroxidases or PRDX4. Dehydroascorbate is an addition source of oxidizing equivalents in the ER. DHa can be moved into the ER from the cytosol, generated in the ER, and like H2O2, DHA directly oxidizes PDI and unfolds reduced proteins in vitro . Due to its slow rate, the former pathway is shown to not be a main route for the reduction of DHA. A faster pathway occurs with a faster PDI-independent oxidation of protein substrates.

References

Bulleid NJ, Ellgaard L. Trends Biochem Sci. 2011 Sep;36(9):485-92. Epub 2011 Jul 19. Review. PMID: 21778060 [PubMed - indexed for MEDLINE]a http://www.ncbi.nlm.nih.gov/pubmed/21778060

Overview

Valence bond theory – this is used to describe hybrid orbitals and electron pairs. It is an extension of the electron dot and bybrid orbital representations. Crystal field theory – this is used to describe the split in metal d-orbitals, which approximates the energy levels from the ultraviolet and visible spectra, but it does not describe bonding Ligand field theory – this is an expansion of the crystal field theory that can be used to describe bonding between the metal ion and the ligands by focusing on the orbital interactions. Angular overlap method – this is used to estimate the orbital energies in a molecular orbital calculation.

Valence Bond Theory

Valence bond theory was originally proposed by Pauling as a way of hybridization between atomic orbitals. This was one of the major theory used to describing the bonding of coordination compounds, but it's rarely used today. The valence bond theory is filled by the Aufbau principle, which states that electrons are most stable when they filled each orbital with one electron before further filling it. This leads to the Madelung rule, which states that orbitals filled by n+l are filled first before orbitals filled by n+l of higher energy orbitals.

Period 4

Period 5

Period 6

Period 7

Element

Z

Electron Configuration

Element

Z

Electron Configuration

Element

Z

Electron Configuration

Element

Z

Electron Configuration

Lanthanum

57

[Xe] 6s2 5d1

Actinium

89

[Rn] 7s2 6d1

Cerium

58

[Xe] 6s2 4f1 5d1

Thorium

90

[Rn] 7s2 6d2

Praseodymium

59

[Xe] 6s2 4f3

Protactinium

91

[Rn] 7s2 5f2 6d1

Neodymium

60

[Xe] 6s2 4f4

Uranium

92

[Rn] 7s2 5f3 6d1

Promethium

61

[Xe] 6s2 4f5

Neptunium

93

[Rn] 7s2 5f4 6d1

Samarium

62

[Xe] 6s2 4f6

Plutonium

94

[Rn] 7s2 5f6

Europium

63

[Xe] 6s2 4f7

Americium

95

[Rn] 7s2 5f7

Gadolinium

64

[Xe] 6s2 4f7 5d1

Curium

96

[Rn] 7s2 5f7 6d1

Terbium

65

[Xe] 6s2 4f9

Berkelium

97

[Rn] 7s2 5f9

Scandium

21

[Ar] 4s2 3d1

Yttrium

39

[Kr] 5s2 4d1

Lutetium

71

[Xe] 6s2 4f14 5d1

Lawrencium

103

[Rn] 7s2 5f14 7p1

Titanium

22

[Ar] 4s2 3d2

Zirconium

40

[Kr] 5s2 4d2

Hafnium

72

[Xe] 6s2 4f14 5d2

Rutherfordium

104

[Rn] 7s2 5f14 6d2

Vanadium

23

[Ar] 4s2 3d3

Niobium

41

[Kr] 5s1 4d4

Tantalum

73

[Xe] 6s2 4f14 5d3

Chromium

24

[Ar] 4s1 3d5

Molybdenum

42

[Kr] 5s1 4d5

Tungsten

74

[Xe] 6s2 4f14 5d4

Manganese

25

[Ar] 4s2 3d5

Technetium

43

[Kr] 5s2 4d5

Rhenium

75

[Xe] 6s2 4f14 5d5

Iron

26

[Ar] 4s2 3d6

Ruthenium

44

[Kr] 5s1 4d7

Osmium

76

[Xe] 6s2 4f14 5d6

Cobalt

27

[Ar] 4s2 3d7

Rhodium

45

[Kr] 5s1 4d8

Iridium

77

[Xe] 6s2 4f14 5d7

Nickel

28

[Ar] 4s2 3d8 or [Ar] 4s1 3d9 (disputed)

Palladium

46

[Kr] 4d10

Platinum

78

[Xe] 6s1 4f14 5d9

Copper

29

[Ar] 4s1 3d10

Silver

47

[Kr] 5s1 4d10

Gold

79

[Xe] 6s1 4f14 5d10

Zinc

30

[Ar] 4s2 3d10

Cadmium

48

[Kr] 5s2 4d10

Mercury

80

[Xe] 6s2 4f14 5d10

Crystal Field Theory

Crystal field theory was originally developed to describe the structure of metal ions in crystals. The energies of the d orbitals are split by electrostatic field. This was developed in the 1930s, which ignored the covalent bonding since ionic crystals didn't describe it. When the ligands come close to the metal, a big destabilization occurs. The dx^2-y^2 and dz^2 have eg symmetry whereas the dxy, dxz, dyz have t2g symmetry, which implies that they're triply degenerate. eg accounts for 40% of the energy difference while t2g accounts for 60% of the energy. While eg's energy rises, t2g's energy falls to accounts for the difference in rising so that their energies cancel out. This energy difference is called crystal field stabilization energy. This also leads to the idea of high spin and low spin. High spin occurs when Δo > the pairing energy with a weak ligand field and low spin occurs when Δo < pairing energy with a strong ligand field.

Factors that influence Δo: 1) charge : the greater the charge on the central ion, the greater the pull from the ion, which results in an increase in oxidation state.

2) identity of the metal : the greater #d orbitals will have a greater Δo. For example, the 5d orbital can interact more efficiently than 3d, which results in a greater Δo.

3) identity of the ligand : stronger Lewis bases have greater Δo

Ligand Field Theory

Ligand field theory is used to describe ligand-metal orbital interactions. The metal's d-orbitals match the irreducible representations Eg and T2g. The metal's s and p orbitals have the symmetry A1g and T1u. This leads to a total of 3 bonding orbitals.

Coordination Numbers

Factors involved in determining the overall shape of a coordination compound

1) the number of bonds – since most bonds are exothermic, the creation of bonds will lead to a greater stabilization of the structure

2) VSEPR considerations

3) Occupancy of d-orbitals – the number of d electrons affect the geometry of the coordination compound

4) Steric interference – occurs when large ligands surround the central metal

5) Crystal packing effects – the regular shape is distorted when packed into a crystal. This would hinder the understanding of the original structure since it'd be unclear as if the distortion came from the crystallization process or naturally.

Coordination number 4 – this is usually obtained from square planer and tetrahedral structures with four ligands. Some examples are: CrO4 2-, Ni(CO)4, [Cu{py)4]+. Coordination number 5 – this is usually obtained from trigonal bipyramid, square pyramid, and pentagonal planes. Some examples are: Fe(CO)5 and PF5. Coordination number 6 – this is the most common coordination number since it is usually obtained from the structure of an octahedral. Some examples are: [Co(en)3]3+ and [Co(NO2)6]3-. Coordination number 7 – this is usually obtained from pentagonal bipyramid, capped trigonal prism, and capped octahedron. Some examples of molecules in this structure are: [NiF7]2- and [NbF7]2-. Coordination number 8 – this is usually only obtained in simple ionic lattice structures like CsCl.

References

Miessler, Gary. Inorganic Chemistry. 4th Edition.

Essentiality of Aluminum and Silicon

Darwin, natural selection and the biological essentiality of aluminum and silicon

By definition, natural selection is a competition in which winners and losers are defined by selection pressures which act upon competitors that are constrained within specific boundaries or arenas. Through experimentation and discovery, it has been assessed that natural selection can be viewed as a force of nature which is as important in biochemical evolution as it is in speciation. Exley takes a closer look at the biological essentiality of both aluminum and silicon. Aluminum in particular is deemed critical and also the most widely abundant metal on earth. In fact, it is the third most abundant element in the Earth’s crust. However, an element’s wide abundance does not correlate with its biological importance. Instead, aluminum is severely hindered by its biological unavailability. Selection of essential metals for biological use is mainly based on several factors including reaction kinetics and reaction thermodynamics. For example, kinetic constraints involve how a biochemical reaction comes to equilibrium which dictate reaction kinetics and ultimately which biochemical pathways are most efficient or favorable and thus predominate. Equilibrium constants dictate properties of reaction products by ways of solubility equilibrium and complex stability.Living systems however concentrate on the importance of kinetics rather than thermodynamics. Specifically kinetics rely on concentration of reactants, products, competitors, and interferences which all aid living systems in its attempt to avoid chemical equilibrium. Thus, kinetics affect the natural selection of which metals become essential to living systems.

Natural Selection of Aluminum

This takes us to a more enhanced view at the metal of aluminum. Despite its abundance, it has no essential role in any biological system in an organism. Silicon, on the other hand, is second most abundant, but is indeed viewed as an essential element. It was later found evident that aluminum’s lack of an indispensable biological role could be attributed to its non-participation in addition to the possibility of simply being selected out of systems altogether. The absence of aluminum is actually quite unfortunate, due to its versatility as a biologically reactive element. With further investigation, it was deemed that the lack of aluminum was not because it was selected out of biochemical systems but because the lack of available biologically-reactive aluminum present for selection. The unavailability of aluminum can be explained by the lack of aluminum in biotic cycles. Less than .001% of aluminum which is cycled through abiotic processes such as rain-fuelled dissolution of mountains, is actually cycled through biotic processes. Even more puzzling is that there are no known biological mechanisms that specifically keep aluminum out from biota, nor are there "biological footprints" left behind in evolutionary encounters with biologically reactive aluminum.

Aluminum has the ability to bind to oxygen-based functional groups, participate in critical redox reactions, and serves the role as an excellent immunogen as an antigen to enable widespread use of particular vaccines. However, aluminum was not selected --due to slow ligand exchange rates. While it is able to bind to oxygen-based functional groups it does not do so quickly enough to efficaciously serve as a metal co-factor for enzymes. Also, the prevalence of the biologically reactive form of Silicon, Silicic acid, reacted with biologically reactive Aluminum Al3+ and thus reduced the amount of Al3+. Exley asserts due to this, other less abundant metals were able to outcompete aluminum.

Exley went on to study salmon, which unveiled an interesting fact about the emergence of silicic acid, which sought out to protect against the toxicity of aluminum. In more familiar terms, it can be said that silicic acid took geochemical control over the availability of aluminum. Silicic acid is the only available biological form of silicon because silicon’s bonds are extremely tough to break and is selected against when it comes to participation in reactions. Silicon is essential, however, but does not possess any biochemical importance. Silicic acid is a weak acid that participates in an interaction with aluminum hydroxide to produce HAS. In doing so, it was successful in significantly reducing the biological availability of aluminum and has further promoted less selection for metals as well. The metals, however, that were selected have been deemed essential. Metals such as Magnesium, Iron, Calcium, Zinc, and Copper have created a cloud over Aluminum’s head, metaphorically, of course. The aluminum environment has lately been progressing through human activities and a sort of biochemical evolution that must account for a biologically reactive aluminum.

Non-Selection of Silicon

While aluminum was selected against due to its lack of availability of its biologically reactive form Al3+, Silicon's biologically reactive form Si(OH)4 has always been available for selection. Silicon has been seen in biological systems such as in the form of silicic acid which can pass through permeable biological membranes mimicking water and protecting against the toxicity of Aluminum. However today, there is little evidence in biotic systems to prove the essentiality of silicon to living systems.

While there is no silicon biochemistry, silicon still asserts evolutionary pressures being an essential element to life. However the only form of biologically reactive Silicon exist in the form of Silicic acid. Silicic acid is a polyprotic weak acid which loses it's first proton at relatively high pH 10. The majority of biochemical reactions occur around neutral pH which thus limits the bioorganic and bioinorganic chemistry of silicon which only exists as a small neutral molecule.

Sequence, Structure and Biophysical Properties of Proteins

Environmental pressures can shape the sequence, structure and other properties of proteins. Proteins which occupy extreme environments such as proteins which are present in high temperature environments need to be more stable. This stability is mostly determined by the thermodynamic stability of the protein which is dictated by the energy gap between the native state and the unfolded and/or the misfolded states.

Changes in sequence and structure that affect both native and unfolded states are both observed by thermophilic proteins. Characteristics such as higher compactness, tightly packed secondary structure are observed within these proteins. These features are naturally selected for higher tolerances to thermal environmental pressures.

Adaptation of Viruses

Viruses are another example to adaptation to extreme environments such as presence outside host cells. However once within a host, they must avoid countermeasures of the host and survive. Virus genomes are usually packed tightly rather than loosely packed. Viruses are also prone to high rates of mutation exemplified by overlapping reading frames in RNA viruses which enable a single mutation to affect more than one protein. Natural selection must select how the protein is packed and thus its sequence and thus the mutational tolerance of the virus. A virus while having a compact protein structure, it is not as compact as thermophilic proteins. While thermophilic proteins are resistant to mutation due to its tightly packed proteins, a virus with looser packed proteins can thus mutate more frequently observed in the high rates of viral mutations.

Reference

Tokuriki N, Oldfield CJ, Uversky VN, Berezovsky IN, Tawfik DS. Department of Biological Chemistry, Weizmann Institute of Science, Rehovot 76100,Israel.

Darwin, natural selection and the biological essentiality of aluminium and silicon. Exley C. The Birchall Centre, Lennard-Jones Laboratories, Keele University, Staffordshire, ST5 5BG, UK.

Overview

Valence bond theory – this is used to describe hybrid orbitals and electron pairs. It is an extension of the electron dot and bybrid orbital representations. Crystal field theory – this is used to describe the split in metal d-orbitals, which approximates the energy levels from the ultraviolet and visible spectra, but it does not describe bonding Ligand field theory – this is an expansion of the crystal field theory that can be used to describe bonding between the metal ion and the ligands by focusing on the orbital interactions. Angular overlap method – this is used to estimate the orbital energies in a molecular orbital calculation.

Valence Bond Theory

Valence bond theory was originally proposed by Pauling as a way of hybridization between atomic orbitals. This was one of the major theory used to describing the bonding of coordination compounds, but it's rarely used today. The valence bond theory is filled by the Aufbau principle, which states that electrons are most stable when they filled each orbital with one electron before further filling it. This leads to the Madelung rule, which states that orbitals filled by n+l are filled first before orbitals filled by n+l of higher energy orbitals.

Period 4

Period 5

Period 6

Period 7

Element

Z

Electron Configuration

Element

Z

Electron Configuration

Element

Z

Electron Configuration

Element

Z

Electron Configuration

Lanthanum

57

[Xe] 6s2 5d1

Actinium

89

[Rn] 7s2 6d1

Cerium

58

[Xe] 6s2 4f1 5d1

Thorium

90

[Rn] 7s2 6d2

Praseodymium

59

[Xe] 6s2 4f3

Protactinium

91

[Rn] 7s2 5f2 6d1

Neodymium

60

[Xe] 6s2 4f4

Uranium

92

[Rn] 7s2 5f3 6d1

Promethium

61

[Xe] 6s2 4f5

Neptunium

93

[Rn] 7s2 5f4 6d1

Samarium

62

[Xe] 6s2 4f6

Plutonium

94

[Rn] 7s2 5f6

Europium

63

[Xe] 6s2 4f7

Americium

95

[Rn] 7s2 5f7

Gadolinium

64

[Xe] 6s2 4f7 5d1

Curium

96

[Rn] 7s2 5f7 6d1

Terbium

65

[Xe] 6s2 4f9

Berkelium

97

[Rn] 7s2 5f9

Scandium

21

[Ar] 4s2 3d1

Yttrium

39

[Kr] 5s2 4d1

Lutetium

71

[Xe] 6s2 4f14 5d1

Lawrencium

103

[Rn] 7s2 5f14 7p1

Titanium

22

[Ar] 4s2 3d2

Zirconium

40

[Kr] 5s2 4d2

Hafnium

72

[Xe] 6s2 4f14 5d2

Rutherfordium

104

[Rn] 7s2 5f14 6d2

Vanadium

23

[Ar] 4s2 3d3

Niobium

41

[Kr] 5s1 4d4

Tantalum

73

[Xe] 6s2 4f14 5d3

Chromium

24

[Ar] 4s1 3d5

Molybdenum

42

[Kr] 5s1 4d5

Tungsten

74

[Xe] 6s2 4f14 5d4

Manganese

25

[Ar] 4s2 3d5

Technetium

43

[Kr] 5s2 4d5

Rhenium

75

[Xe] 6s2 4f14 5d5

Iron

26

[Ar] 4s2 3d6

Ruthenium

44

[Kr] 5s1 4d7

Osmium

76

[Xe] 6s2 4f14 5d6

Cobalt

27

[Ar] 4s2 3d7

Rhodium

45

[Kr] 5s1 4d8

Iridium

77

[Xe] 6s2 4f14 5d7

Nickel

28

[Ar] 4s2 3d8 or [Ar] 4s1 3d9 (disputed)

Palladium

46

[Kr] 4d10

Platinum

78

[Xe] 6s1 4f14 5d9

Copper

29

[Ar] 4s1 3d10

Silver

47

[Kr] 5s1 4d10

Gold

79

[Xe] 6s1 4f14 5d10

Zinc

30

[Ar] 4s2 3d10

Cadmium

48

[Kr] 5s2 4d10

Mercury

80

[Xe] 6s2 4f14 5d10

Crystal Field Theory

Crystal field theory was originally developed to describe the structure of metal ions in crystals. The energies of the d orbitals are split by electrostatic field. This was developed in the 1930s, which ignored the covalent bonding since ionic crystals didn't describe it. When the ligands come close to the metal, a big destabilization occurs. The dx^2-y^2 and dz^2 have eg symmetry whereas the dxy, dxz, dyz have t2g symmetry, which implies that they're triply degenerate. eg accounts for 40% of the energy difference while t2g accounts for 60% of the energy. While eg's energy rises, t2g's energy falls to accounts for the difference in rising so that their energies cancel out. This energy difference is called crystal field stabilization energy. This also leads to the idea of high spin and low spin. High spin occurs when Δo > the pairing energy with a weak ligand field and low spin occurs when Δo < pairing energy with a strong ligand field.

Factors that influence Δo: 1) charge : the greater the charge on the central ion, the greater the pull from the ion, which results in an increase in oxidation state.

2) identity of the metal : the greater #d orbitals will have a greater Δo. For example, the 5d orbital can interact more efficiently than 3d, which results in a greater Δo.

3) identity of the ligand : stronger Lewis bases have greater Δo

Ligand Field Theory

Ligand field theory is used to describe ligand-metal orbital interactions. The metal's d-orbitals match the irreducible representations Eg and T2g. The metal's s and p orbitals have the symmetry A1g and T1u. This leads to a total of 3 bonding orbitals.

Coordination Numbers

Factors involved in determining the overall shape of a coordination compound

1) the number of bonds – since most bonds are exothermic, the creation of bonds will lead to a greater stabilization of the structure

2) VSEPR considerations

3) Occupancy of d-orbitals – the number of d electrons affect the geometry of the coordination compound

4) Steric interference – occurs when large ligands surround the central metal

5) Crystal packing effects – the regular shape is distorted when packed into a crystal. This would hinder the understanding of the original structure since it'd be unclear as if the distortion came from the crystallization process or naturally.

Coordination number 4 – this is usually obtained from square planer and tetrahedral structures with four ligands. Some examples are: CrO4 2-, Ni(CO)4, [Cu{py)4]+. Coordination number 5 – this is usually obtained from trigonal bipyramid, square pyramid, and pentagonal planes. Some examples are: Fe(CO)5 and PF5. Coordination number 6 – this is the most common coordination number since it is usually obtained from the structure of an octahedral. Some examples are: [Co(en)3]3+ and [Co(NO2)6]3-. Coordination number 7 – this is usually obtained from pentagonal bipyramid, capped trigonal prism, and capped octahedron. Some examples of molecules in this structure are: [NiF7]2- and [NbF7]2-. Coordination number 8 – this is usually only obtained in simple ionic lattice structures like CsCl.

References

Miessler, Gary. Inorganic Chemistry. 4th Edition.

The science of studying carbon-containing molecules is known as organic chemistry. One of the properties of the carbon atom that makes life possible is its ability to form four covalent bonds with other atoms, including other carbon atoms. This binding ability with comes from having four electrons in the carbon’s outer shell, causing it to need four additional electrons for its outer shell to be full.

Role of Carbon in organic chemistry

In living organisms, carbon atoms most commonly form covalent bonds with other carbons and with hydrogen, oxygen, nitrogen and sulfur atoms. Bonds between two carbon atoms, between carbon and oxygen, or between carbon and nitrogen can be single or double in organic compounds. Bonds of a higher order between these atoms can be found in inorganic compounds however. The combination of carbon with itself and with different elements and different types of bonds allows a vast number of organic compounds to be formed from only a few chemical elements. This is made all the more impressive because carbon bonds may occur in configurations that are linear, ring like, or highly branched. Such molecular shapes can produce molecules with a variety of functions. One last feature of carbon that is important in biochemistry is that carbon bonds are stable at the different temperatures associated with life. This property arises in part because the carbon atom is very small compared to most other atoms, and therefore the distance between carbon atoms forming carbon – carbon bonds is quite short. Shorter bonds tend to be stronger and more stable than longer bonds between two large atoms. Thus, carbon atoms are compatible with what we observe about life today, namely that living organisms can inhibit environments ranging from the earth’s icy poles to deep-sea vents. Aside from the simplest hydrocarbons, most organic molecules and macromolecules contain functional groups – group of atoms with special chemical features that are functionally important. Each type of functional group exhibits the same properties in all molecules in which it occurs. For example, the amino group (NH₂) acts like a base. At the pH found in living organisms, amino groups readily bind H⁺ to become NH₃⁺, thereby removing H⁺from an aqueous solution and raising the pH.

Synthesis of Carbon-Carbon Bond

The synthesis of new carbon-carbon bonds in organic reactions is an important synthetic organic technique that leads to the production of artificial chemicals such as new drugs and plastics. In the carbonyl chemistry many synthetic techniques are based on natural processes for the formation of carbon-carbon bonds in biological systems. Some examples of organic reactions forming new carbon-carbon bonds include Aldol reactions, Claisen condensation, Diels–Alder reaction, and Michael reaction.

Claisen Condensation

An Aldol reaction is a powerful technique forming new carbon-carbon bonds in organic chemistry, since it unites two simple molecules into a complex one. The reaction combines two carbonyl compounds to form a new β-hydroxy carbonyl compound. Producsts of such reactions are called aldols, known as the product of aldehyde + alcohol. A typical Aldol reaction involve the nucleophilic addition of a ketone enolate to an aldehyde. Aldol condensation takes place when the aldol product lose a water molecule to form an α,β-unsaturated carbonyl compound. Nucleophiles that can be employed in the aldol reaction include the enols, enolates, and enol ethers of ketones, aldehydes, and other compounds carrying the carbonyl function, whereas the electrophilic reagent is usually an aldehyde or ketone. When different nucleophile and electrophile are used, the reaction is called a crossed aldol reaction. On the other hand, a reaction in which the same nucleophile and electrophile are employed is called aldol dimerization.

Typical Aldol Reaction Mechanism

A Claisen condensation occurs between two esters or one ester and another carbonyl compound in the presence of a strong base, which results in a β-keto ester or a β-diketone. In Claisen condensation, attack of an ester enolate on a carbonyl group generates a new carbon-carbon bond. The reaction mechanism involves ester enolate formation by reacting ethyl acetate with a stoichiometric amount of reagents with ester function, nucleophilic addition of another ester molecule that furnishes a ketoester, elimination of the alkoxide group, and finally the deprotonation of ketoester followed by protonation upon aqueous work-up. The overall process is endothermic and all steps before the deprotonation of ketoester are reversible. The deprotonation of ketoester drives equilibrium, since it removes the base needed to catalyze the previous steps. To prevent transesterification, both the alkoxide and ester are usually derived from the same alcohol.

Claisen Condensation Mechanism

In a Diels-Alder reaction, a conjugated diene adds to a substituted alkene to yield substituted cyclohexane derivatives. Such reaction is a special case of cycloaddition reactions between pi systems; four conjugated atoms containing four pi electrons reacts with a double bond containing two pi electrons. The four-carbon component is called diene and the alkene added is called dienophile. The reaction is also called a [4+2]cycloaddition. This type of reaction can still be carried out in the absence of carbon in the newly formed ring. Diels-Alder reactions that are reversible are called the retro-Diels-Alder; for example, the decomposition reaction of the cyclic system.

Diels-Alder Cycloaddition Mechanism

The Michael reaction is the nucleophilic addition of a carbanion or another nucleophile to an alpha, beta unsaturated carbonyl compound. The stabilized anions derived from β-dicarbony compounds with α,β-unsaturated carbonyl compound leads to 1,4-additions. It is one of the most useful methods for the mild formation of new carbon-carbon bonds. A Michael addition is base-catalyzed and works with α,β-unsaturated ketones, aldehydes, and other carboxylic acid derivatives; they are known as Michael acceptors. A Michael donor is an electron-withdrawing group on the nucleophile such acyl and cyano. [Vollhardt] [8]

Michael Reaction Mechanism

Isomers in organic chemistry

Organic molecules also have isomers. Isomers are molecules that contain the same number of atoms and also the same kind of atoms. However, they have different bonding arrangements. Types of isomers include constitutional (or structural) isomers, and stereoisomers.

Constitutional isomers (structural isomers) are the compounds that have the same molecular formula but differ in how the atom are arranged and connected. Chain isomers, positional isomers, and functional group isomers are constitutional isomers.

Chain isomers

Example: Pentane and 2-Methylbutane

vs.

Positional isomers

Example: 3-Hexanone and 2-Hexanone

vs.

Functional isomers

Example: Ethanol and Dimethylether

Stereoisomers include conformational isomers and configurational isomers. Conformational isomers are compounds that posses the same molecular formula and atomic connectivity but differ in a rotation about a bond. In other words, conformational isomers can be interconverted by rotation about single bonds. They are not separable at room temperature. There are different kinds of conformational isomers. They are eclipsed, staggered, anti, and gauche conformations. Configurational isomers are those isomers which can only be interconverted by breaking bonds. There are two different types of configuration isomers. They are enantiomers and diastereomers. Enantiomers are non-superimposable mirror images. Diastereomers are non-superimposable non-mirror images.

Glucose and fructose, which both have the same chemical formula of C₆H₁₂O₆ but a different arrangement in atoms, would be a good example of constitutional/structural isomers. Enantiomers would be a good example of a type of stereoisomer. Stereoisomers are isomers that have exact same bonding between atoms, but differ in their specific spatial arrangements. For example, enantiomers are isomers that are mirror images of each other. They can be superimposed on each other (achiral) or not (chiral), and may exhibit either R (clockwise) or S (counter-clockwise) configurations. Each mirror image of a chiral molecule expresses different properties than its counterpart. A way to determine whether a molecule is chiral or not is by looking for and identifying chiral centers. A chiral center for carbon, will have four different groups bonded to it. It must be sp3 hybridized and be tetrahedral in shape. A molecule's chiral center can also be referred to as a stereocenter.

Fullerenes are organic molecules that consist only of carbon atoms.

The role of organic synthesis research in Biochemistry

Organic synthesis is the science of constructing molecules. There are two major areas of research in organic synthesis: exploratory and target oriented. Research in both fields requires innovation, imagination, and artistic creativity.

Exploratory research involves the development of new organic reactions. Many researchers in this field focus on the optimization of previously known reactions. There are many factors when developing new organic reactions such as reactant, solvent, temperature, pH, etc. The main goals for any researcher in this field are to maximize yield of desired product, minimize side reactions/products, and be reliable for a broad spectrum of starting material. The advancement of the methodology aspect of organic synthesis expands the tools and techniques used by target oriented research.

Target oriented research involves the development of organic molecules through a series of organic reactions. Researchers in this field use preexisting reactions and commercially available materials to synthesis desired products. “Target” molecules are either natural products or designed molecules. A linear synthesis (a linear series of reactions conducted one after another) will suffice with simple molecules. Other approaches such as convergent synthesis (independent synthesis of key intermediates) are used for complex molecular structures. Methods such as solid phase synthesis are exceptionally useful in the synthesis of proteins.

“Target” molecules are biologically, medicinally, and/or theoretically interesting products. The field of bioorganic synthesis began with the synthesis of urea by Friedrich Wöhler in 1828. Natural product synthesis has been recognized with Nobel Prize in Chemistry on several occasions. Target oriented natural products are immensely useful for medical research. Cancer inhibiting molecules have been found in several natural products including marine natural products.

New reactions are constantly being developed and optimized. Many of these reactions are particularly useful in the synthesis of drugs and biomarkers. Reactions such as the palladium cross-coupling reaction, winner of the 2010 Nobel Prize in Chemistry, have been immensely useful in the synthesis of drugs, biomarkers, and other useful molecules.

The simplest Suzuki coupling reaction involves a palladium cross-coupling reaction of phenylboronic acid and bromobenzene to yield phenylbenzene.

Straying away from traditional synthetic methods has brought monumental advancements in alternate molecule building techniques in the last few decades. An example of an innovative technique is solid phase synthesis, a method in which molecular building blocks are attached to a bead and the “target” molecule is obtained with a linear series of reactant solutions. There are several advantages of solid phase synthesis compared to traditional solvent-based synthesis. Some advantages are that functional groups can be easily protected and also it is easier to extract unwanted byproducts or reactants from the desired “target” molecule. Solid phase synthesis is exceptionally useful with the synthesis of peptides, deoxyribonucleic acid (DNA), and other sequence-based molecules.

“There is excitement, adventure and challenge, and there can be great art in organic synthesis” – R.B. Woodward (Nobel Prize, 1965)

Synthetic Approach to Activated Amino Acids

The use of coupling agents in peptide synthesis is way of activating the carbonyl carbon of one amino acid, thus rendering it more reactive to the adjacent amino acid's amine group. Dicyclohexylcarbodiimide and various uronium salts are predominantly used as coupling agents in the field [1]. Although the use of coupling agents isn't a daunting task, constructing an intrinsically activated amino acid would circumvent the use of these reagents, saving time, money, and reaction yield.

The general synthetic approach at obtaining an activated amino acid will be described herein. Reaction of a specific isocyanide with an arbitrary aldehyde yields and alpha-hydroxy indole species [2].

Indole

Converting the hydroxy function on this species to an amino group would yield an activated amino acid. Notice that the indole function is aromatically stabilized, acts as a great leaving group, and thus activates the carbonyl carbon. It is imperative that converted amine function is protected, so that cross reaction between monomers doesn't occur. Carrying out this task isn't easy, because many chemical syntheses require extreme conditions which aren't suitable for peptide synthesis. High temperature or pH fluctuations can easily break peptide bonds, rendering the targeted peptide destroyed. Therefore, mild conditions in carrying out this displacement are sought after. [9] [Gianneschi]

Click Chemistry

The ability of generating molecular modification with high selectivity is invaluable for studies of chemical and biological systems. Click chemistry is the chemical philosophy of synthesizing molecules from a core group of highly effective reactions developed by Sharpless, Finn, and Kolb in 2001. The inspiration for the development of Click chemistry came from the idea that nature tends to produce substances from smaller subunits. The logic behind click chemistry is to bind small molecular units together to produce products with reactions that proceed rapidly in high yields under ambient conditions. These reactions are required to have a high thermodynamic driving force that is orthogonal to other functional groups that may be present in biomolecules. Click chemistry reactions are effective for labeling biomolecules. They also proceed in biological conditions with high yield. An important aspect of the reactions is that they are bioorthogonal, meaning that they don’t react with functional groups in the biological systems. Some examples of Click chemistry reactions are (a) Azide-Alkyne Cycloaddition, (b) Copper-Free Azide Alkyne Cycloaddition, and (c) Staudinger Ligation shown in the schemes below.

The major goals of Click Synthesis were to simplify the methods of how molecules are synthesized and, consequently, improve the process of identifying and synthesizing molecules with biological importance. These methods have proven to be beneficial in modern drug development; particularly useful for in situ fragment-based drug design. In situ drug development via Click chemistry has been extended to the selective generation of potent inhibitors of carbonic anhydrase and HIV-1 protease. A schematic representation of the process of inhibitor development with aid of Click chemistry is shown below.

These techniques have been extremely useful in combinatorial drug development. The product of the azide-alkyne cycloaddition, triazole, is favorable in drug design because it possesses a variety of useful properties. The highly tuned properties of these reactions possessing high yields, selectivity, and ability to undergo transformation in mild (biological) environments allow the products to be directly analyzed for activity without purification (which is a significant shortcut!). Example of direct screening of click chemistry products with alpha-1,3-fucosyltransferase (fuc-T) show below.

Click chemistry designs include reactions that are opposed to the reaction with biological molecules. This is a very useful property for selectively labeling molecules to detect in biological systems. The Staudinger ligation and azide-alkyne cycloadditions have proven to be very helpful for tasks that were very challenging prior.

There are is an extensive range of applications, a brief description and example schematic depictions are presented below:

(1) Introduction of Unnatural Amino Acids Bearing Reactive Tags into Proteins.

(2) Labeling of Viral Surfaces.

(3) Incorporation of Labeled Probes onto Proteins via Post-Translational Modification.

(4) Labeling of Nucleotides for Imaging DNA and RNA.

(5) Derivatization of Lipid Probes.

(6) Activity- Based Protein Profiling.

The benefit of these methods is clear from the broad spectrum of applications in chemical biology. The scope of these reactions is quite broad and takes innovation of the researchers using these methods to maximize its potential. These methods expand to pure organic chemistry lab as well. Without constraints of biological environments, these reactions have obvious beneficial aspects with precise control of target structures. The progression of bioorthogonal chemistry (has and) continues to produce new and effect tools for the future of research. [Best]

Wöhler's Synthesis of Urea

Wöhler's synthesis is one of the examples of synthesis in which molecules are made. Carbon compounds, organic products, in this reaction are made from inorganic salt.

Pb(OCN)₂ + 2 H₂O + 2 NH₃ -----> 2 H₂N(C=O)NH₂ + Pb(OH)₂

Lead cyanate Water Ammonia Urea Lead hydroxide

The effect of acid on food digestion and stomach acid

The normal human stomach has about 0.02 M of HCl per day. If there's an increase of HCl in the stomach, the pH of stomach juice would falls from 2.5 to 1.0 . HCl would mess up the normal folded shapes of protein molecules in food. As a result, the acid would destroy many digestive enzymes in the stomach. Therefore, to protect itself from the increase of such acid, the cells in the stomach must work together to prevent it from happening. First, the stomach tissues are made of protein molecules with the interior covered with layers of gastric mucosa. When stimuli such as smelling and tasting activate the cells in the gastric mucosa, the signal molecule, histamine, would make the parietal cells to hide the acid juices in the stomach. As a result, there would be an increase in the acid juices production in the stomach. In order to prevent it from happening, active ingredients such as cimetidine, famotidine, and ranitidine help to reduce the acid by obstructing the histamine from helping the parietal cells.

References

http://en.wikipedia.org/wiki/Claisen_condensation

http://en.wikipedia.org/wiki/Aldol_reaction

http://en.wikipedia.org/wiki/Claisen_condensation

http://en.wikipedia.org/wiki/Diels%E2%80%93Alder_reaction

http://en.wikipedia.org/wiki/Michael_reaction

Vollhardt and Schore. Organic Chemistry. 6th edition. New York: W.H. Freeman and Company

http://en.wikipedia.org/wiki/Carbon%E2%80%93carbon_bond

http://en.wikipedia.org/wiki/Peptide_synthesis

Gianneschi, C. N., Rubinshtein, M., James, R.C., Kobayashi, Y., Yang, J., Young, J., Yanyan, J.M. Org. Lett., 2010, 12 (15), pp 3560–3563

Best, M.D. Biochemistry, 2009, 48 (6571), pp 6571–6584

General information

The term “functional group” is one that is used almost exclusively in organic chemistry. In organic chemistry, functional groups are a set of specific atoms within a molecule that determine the molecule’s overall reactivity and properties. Small differences in functional groups, like the difference between an aldehyde and a ketone, can result in drastic differences in the properties and reactivity of the molecules. Compounds that have the same functional groups will have the same type of reactivity, though small deviances will result if other substituents or functional groups are present on the molecule.

An example of a Grignard reaction

There are 7 important functional groups in the chemistry of life: Hydroxyl, Carbonyl, Carboxyl, Amino, Thiol, Phosphate, and aldehyde groups. {Not for 10 class you'll be confused}

1) Hydroxyl group: consists of a hydrogen atom covalently bonded to an oxygen atom. The hydroxyl group is denoted by -OH in chemical structures and has a valence charge of -1 when in the hydroxyl ion form. It is present in Alcohols and Carboxylic Acid molecules.

2) Carbonyl group: is written as a covalent C=O double bond. It is a very polar molecule and the carbonyl carbon can serve as a reaction site for many reactions. It is present in Aldehydes, Ketones, Esters, Anhydrides, and Carboxylic Acids.

3) Carboxyl group: is the monovalent group -COOH, consisting of a carbonyl group bound to a hydroxyl group. It is the main functional group in organic acids (carboxylic acids), in which a proton can dissociate and lead to a strong stabilizing resonance that forms between the two oxygens.

4) Amino group: The amino group (–NH2) consists of one atom of nitrogen attached by covalent bonds two atoms of hydrogen, and a spot where it is attached to an alkyl group, or another hetero atom. This leaves a lone valence electron pair on the nitrogen which is available for bonding to another atom.

5) Thiol group: a functional group containing a sulfur atom bonded to a hydrogen atom. General formula: -SH. The amino acid cysteine contains a thiol group. It is the thiol derivative of a hydroxyl group, and can undergo similar reactions.

6) Phosphate group: A functional group or radical comprised of phosphorus attached to four oxygen (a diester), and with a net negative charge, thus represented as PO4-. The phosphate group is important in living things in varying ways. Firstly, it is an important structural backbone component of nucleotides, which is the basic structural unit of DNA and RNA. Secondly, it is used as an electron transfer component of energy-rich molecules, such as ATP. Thirdly, it is also bound to coenzymes like NADP / NADPH involved in anabolic reactions (such as photosynthesis in plants and lipid synthesis in animals). It is also a part of the hydrophilic head of phospholipids in biological membranes.

7) Aldehyde group: Aldehydes are the second compound containing the carbonyl group (C=O). In aldehydes, one of the two groups attached to the carbonyl carbon is an alkyl group, while the other is a hydrogen atom. This means that aldehyde carbonyl occur at the end of a chain of carbon atoms (i.e. in sugars, R-CHO is at the end of a chain, while a carbonyl inside the chain attached to two carbons is a ketone). They can form acetals and hemiacetals when exposed to water, or in the presence of hydroxy groups. When an aldehyde group is the highest priority functional group present in the molecule, it is names as an alkanal (note, the -e is dropped). The numbering scheme used will be the one that gives the carbonyl carbon atom the lowest possible number (i.e. number 1). Other functional groups are located by this numbering scheme. Since the carbonyl group is always at carbon number 1, there is no need to indicate the location. It is possible to have the -CHO group directly attached to a ring. If it is the highest priority group in the molecule, the ring is a substituent to the aldehyde and it is named as a cycloalkyl carbaldehyde or cycloakanecarbaldehyde (an older way is as a carboxaldehyde). There are a number of compounds which were named before IUPAC developed the standardized nomenclature rules. Many of these compounds are still referred to by these common names.

Other common functional groups include:

• Haloalkanes: can also be named as alkyl halides despite the fact that the halogens are higher priority than alkanes. The alkyl halide nomenclature is most common when the alkyl group is simple.

Example of Haloalkane

• Ethers: If both groups are simple alkyl groups, then the ether is usually named as alkyl alkyl ether. If the two alkyl groups are the same, then it's a dialkyl ether.
• Thiols: A sulfur-containing organic compound having the general formula RSH, where R is another element or radical.
• Alkenes: The root name is based on the longest chain containing both ends of the alkene unit, the C=C. The chain is numbered so as to give the alkene unit the lowest possible numbers. The locant for the first carbon of the alkene is used in the name.
• Alkynes: The alkynes are the third homologous series of organic compounds of hydrogen and carbon, where there is at least one triple-bond between the atoms in the molecules.
• Aromatic Compounds: a hydrocarbon containing one or more benzene rings that are characteristic of the benzene series of compounds.

Aromatic Compound With Alcohol Substituent

• Anhydrides: are used in the preparation of esters. Ethyl acetate and butyl acetate (from butyl alcohol and acetic anhydride) are excellent solvents for cellulose nitrate lacquers. Acetates of high-molecular-weight alcohols are used as plasticizers for plastics and resins. Cellulose and acetic anhydride give cellulose acetate, used in acetate rayon and photographic film. The reaction of anhydrides with sodium peroxide forms peroxides (acetyl peroxide is violently explosive), used as catalysts for polymerization reactions and for addition of alkyl halides to alkenes. In Friedel-Crafts reactions, anhydrides react with aromatic compounds, forming ketones such as acetophenone.

Anhydrides react with water to form the parent acid, with alcohols to give esters, and with ammonia to yield amides; and with primary or secondary amines, they furnish N-substituted and N,N-disubstituted amides, respectively.

Read more: http://www.answers.com/topic/acid-anhydride#ixzz2DkMTM1Db

Read more: http://www.answers.com/topic/acid-anhydride#ixzz2DkMNN84B

• Esters: is any of a class of organic compounds corresponding to the inorganic salts and formed from an organic acid and an alcohol.
• Nitriles: are typically undergo nucleophilic addition to give products that often undergo a further reaction. The chemistry of the nitrile functional group, C=N, is very similar to that of the carbonyl, C=O of aldehydes and ketones.

Basic Information

The hydroxy-group is one of many functional groups studied in organic chemistry. The presence of a hydroxy group indicates that the molecule is either an alcohol or a carboxylic acid. The chemical representation of a hydroxy group is ^–OH, indicating an oxygen atom covalently bonded to a hydrogen atom.

Within the general functional group of alcohol, there are three subgroups: primary, secondary, and tertiary alcohols. An alcohol is classified as a primary, secondary, or tertiary depending on the carbon it is attached to. An alcohol is primary if it is attached to a carbon with only one carbon-carbon bond. By the same token, an alcohol is secondary if the carbon to which it is attached is bound to two other carbons, and tertiary if the carbon’s remaining three bonds are bonded to other carbons. The following examples of primary, secondary, and tertiary alcohols are ethanol, isopropanol (IUPAC nomenclature: propan-2-ol), and tert-butyl alcohol (IUPAC nomenclature: 2-methylpropan-2-ol), respectively.

Ethanol, a primary alcohol

Isopropanol, a secondary alcohol

t-Butyl Alcohol, a tertiary alcohol

Physical Properties of Alcohols

The bond existing between oxygen and hydrogen is significantly shorter than the bonds formed between carbon and hydrogen. Hydroxy groups are inherently polar. The oxygen in the covalent bond is highly electronegative; thereby, pulling the majority of the electron density shared in the covalent bond towards itself. From this information it becomes clear that the oxygen bears a significant partial negative charge, while the hydrogen bears a partial positive charge of the same magnitude.

The presence of a highly electronegative group allows alcohols to hydrogen bond with other molecules. For example, in ethanol, the negatively charged oxygen on one ethanol molecule is attracted to a positively charged hydrogen molecule on another ethanol molecule. The ability of alcohols to hydrogen bond amongst themselves explains their high melting points. Alcohols also have higher boiling points than the hydrocarbons and ethers which contain a comparable number of carbon atoms.

The solubility of alcohols in water is interesting because alcohols have hydrophobic and hydrophilic centers, making them amphipathic or amphiphilic. The hydrocarbon portion of the alcohol is nonpolar, and thus resists dissolving in water; the polar alcohol group is hydrophilic and thereby promotes dissolution in water. In a small molecule such as propanol, the alcohol dissolves because the polar alcohol group outweighs the very short hydrogen chain. However, as the hydrocarbon chain increases beyond butanol to pentanol and hexanol, alcohols become virtually immiscible in water.

Alcohols are amphoteric species that may exist as both acids and bases. At low pH, or when surrounded by strong acids, they may be protonated to alkyloxonium ions. In contrast, at high pH and when surrounded by strong bases, they may exist as alkoxide ions.

Phenols

Structure of Phenol

Phenols are hydroxy-arenes. They have some similarities to alcohols in chemical reactivity at oxygen, but they are much more acidic and much less basic. They appear to be enols (hydroxy attached to an alkene) and should be unstable relative to the keto form, but the gain in aromaticity in being in the enol favors the enol form relative to the keto. Examination of the resonance forms indicates that the hydroxyl donates electron density to the ring, hence the hydroxyl group is activating and ortho-para directing.

Nomenclature of Phenols

Substituted phenols are named as derivatives of the parent compound phenol with the hydroxyl-bearing carbon being designated as C-1. Some examples are:

2-Methyl-4,6-dinitrophenol	Benzene-1,3-diol

Acidity of Phenols

The pka of phenols is around 8-10. Phenols are substantially more acidic than alcohols. This is because phenoxide (the corresponding conjugate base) is resonance stabilized.

Synthesis of Phenols: Nucleophilic Aromatic Substitution

Direct concerted displacement of leaving groups on aromatic rings is not possible. In appropriately substituted rings, leaving groups may be replaced by hydroxyl involving nucleophilic processes via a 2-step mechanism: nucleophilic aromatic substitution.

Alcohol-like Reactivity of Phenols

Phenols undergo many of the same reactions as alcohols, especially where the nucleophicity of the corresponding phenoxide (alkoxide) is concerned.

Phenol alkylation via Williamson-Type Reaction

Synthesis of Alcohols

In industry, methanol is made on a multibillion-pound scale from a pressurized mixture of CO and H₂ called synthesis gas.The reaction involves a catalyst consisting of copper, zinc oxide and chromium(III) oxide.

Changing the catalyst to rhodium or ruthenium leads to 1,2-ethanediol (ethylene glycol), an important industrial chemical that is the principal component of automobile antifreeze. Ethanol is prepared in large quantities by fermentation of sugars or by the phosphoric acid-catalyzed hydration of ethene (ethylene).

On a smaller than industrial scale, we can prepare alcohols from a wide variety of starting materials. For example, conversions of haloalkanes into alcohols by Sn2 and Sn1 processes featuring hydroxide and water. These methods are not as widely used as one might think, however, because the required halides are often accessible only from the corresponding alcohols. They also suffer from the usual drawbacks of nucleophilic substitution: Bimolecular elimination can be a major side reaction of hindered systems, and tertiary halides from carbocations that may undergo E1 reactions. Some of these drawbacks are overcome by the use of polar, aprotic solvent.

Conceptually, the easiest way to reduce a carbonyl group would be to add hydrogen, H-H, across the carbon-oxygen double bond directly. Although this can be done, it requires high pressures and special catalysts. A more convenient way is a polar process, in which hydride ion, and a proton, are delivered to the double bond, either simultaneously or sequentially.

Organometallic Reagents

Organometallic reagents of magnesium and lithium are used in the synthesis of alcohols. Since the metal in organometallic compounds are extremely electropositive, the carbon can react as if it were carrying a negative charge (giving it nucleophilic behavior). These organometallic compounds can attack the carbonyl group of aldehydes and ketones. An important feature of this reaction is that a new carbon-carbon bond is formed in the process.

Synthesis of Tertiary Alcohol From Ketone

The first step of this reaction involves the nucleophilic carbon attacking the carbonyl carbon of a ketone or aldehyde to form the carbon-carbon bond. Simultaneously, two electrons from the carbon-oxygen double bond (carbonyl) are pushed by the electrons from the attacking carbon to form a metal alkoxide with the metal from the organometallic compound. Upon aqueous work-up, an alcohol is formed by hydrolyzing the metal alkoxide.

Different types of alcohols can be formed by changing the compound reacting with the organometallic compound. For example, a reaction between an organometallic compound and formaldehyde will result in a primary alcohol. Reaction with any other type of aldehyde will produce secondary alcohols; reaction with ketones will yield tertiary alcohols.

Hydroxylation

An important function of metabolic degradation is the introduction of hydroxy groups into unfunctionalized parts of molecules in our body (hydroxylation). A set of proteins, known as the cytochrome proteins, are vital biomolecules that are present in almost all living cells, which assist with the process of hydroxylation. An example of a cytochrome protein is cytochrome P-450. Cytochrome P-450 uses O₂ to add hydroxy groups to alkyl substituents. The importance of hydroxylation can be signified by the excretion of drugs to prevent toxic accumulation. Hydroxylation helps increase the water solubility of compounds through addition of hydroxy groups to specific regions of the compounds. Thus in the liver, hydroxylation helps to accelerate the excretion of medicine that we ingest by making it more soluble. If it weren't for this process, ingested drugs would stay in our bodies much longer and have a much greater chance of toxic accumulation.

Hydroxylation can also be selective in the addition of hydroxy groups to a molecule. Protein's undergoing hydroxylation can control the order in which hydroxy groups are added, as well as, the positions to which they are added. This is a key function of the process of hydroxylation, as it allows molecules to undergo oxidation at specific regions of the molecule. Also it allows molecules to bind or react with other compounds in specific orientations.

Nomenclature

The system used to name organic molecules – IUPAC – is relatively simple to use for alcohols, though it can get more complicated if the molecule containing the alcohol group contains substituents containing other functional groups, such as alkenes, alkynes or halogens. The simplest case is when the molecule contains only hydrocarbons and the hydroxy group is the sole functional group, making it purely an alcohol.

In general, the presence of functional groups is indicated by a characteristic suffix that is distinct for each functional group and a number designating the location of the functional group on the molecule. For alcohols, the characteristic suffix used is –ol.

Case 1: Hydrocarbons and alcohols In this case, the longest chain of the hydrocarbons is counted – this becomes the primary hydrocarbon chain of the molecule. Then the longest carbon chain is numbered with the goal of putting as small a number as possible on the hydroxyl group. The longest chain is named as it is for alkanes, but the suffix for alcohols is –ol. Thus, for a propane chain which contains an alcohol group on the second carbon is named propan-2-ol or 2-propanol. Because the alcohol is the only functional group, it does not matter which of the two names is used. It is assumed that the “2” corresponds with the location of the hydroxy group. However, this rule does not always hold true when other functional groups are present.

Case 2: This is more complex because there is a hierarchy of functional groups in organic chemistry. Each functional group when isolated wants to be labeled as the smallest number. However, some functional groups are of “higher importance” and thus are numbered first. Also, if the alcohol is no longer the most “important” functional group, the molecule does not the suffix for the alcohol. This may present a slight problem initially, because the –ol was the only hint that the molecule possessed an –OH group. However, the presence of the –OH group is designated by using the word “hydroxy” after the number of the carbon on which the –OH group is bound to. An example is: 1-chloro-2-hydroxybenzene.

If there are multiple –OH groups present, the nomenclature is modified only slightly. The suffix used in the name is the number of alcohol groups followed by –ol. For example, a compound containing two –OH groups has a suffix of “diol”; for a compound with three –OH groups, the suffix is a “triol.” The numbers of the carbons containg the alcohol groups are stated in the same manner as if there were only one; giving the hydroxy group the lowest number possible. For example, a diol would be named as 3,3 pentandiol. This name indicates that the third carbon in the five-carbon chain is bonded to two –OH groups.

Unique Infrared Spectroscopy of Alcohols and Phenols

Infrared Spectroscopy is a technique used by chemists to elucidate structural information about an organic compound. Organic molecules absorb electromagnetic radiation in the infrared region and promote specific vibrational states. These vibrations are called fundamental absorptions, and they arise from the excitation from the ground state to the lowest-energy excited state. These modes of vibration are then picked up by a machine which interprets the radiation in the form of a spectrum. Since every type of covalent bond has a different natural frequency of vibration, two different molecules will never share the same infrared spectrum. The absorption of certain bonds have very distinct vibrational frequencies. A small range of absorption can be allocated to a specific type of bond. For instance, an absorption around 1715 is normally due to the carbonyl group.

Alcohols and Phenols have unique Infrared spectrum and usually show strong and broad streching bands between 3400 and 3300 cm⁻¹. In solution, it is sometimes possible to observe a free stretching band at around 3600⁻¹ with a sharp but weaker peak.

The broad peak that is well known as the alcohol or phenol peak is created by the stretching vibrations between the O-H. Intermolecular hydrogen bonding usually weakens the O-H bond, and shifts the band to a lower frequency. Phenols are known to have a broader O-H band than alcohols. However, it is hard to determine the difference by looking at the O-H band due to their similar shapes. The bending vibrations between the C-O-H usually yield weak and broad peaks around 1440–1220 cm⁻¹. However, these peaks are difficult to see due to interference by similar bands from strong CH₃ bending peaks. Strong C-O bond stretches are observed from 1260 – 1000 cm⁻¹. Since the C-O bond is adjacent to the C-C bond, the shape and size of the peak can be used to determine whether an alcohol is primary, secondary, or tertiary. It can also be used to distinguish between phenols and alcohols. In addition, alcohols that are present in amides, ethers, and aldehydes can also be distinguished in IR spectrum. Infared spectrum is a useful analytical tool that is used to identify which function groups are present or which ones are not.

NMR of alcohol groups

Nuclear magnetic resonance (NMR) is a technique used to identify organic compounds. Since the atomic nucleus has a spinning charge due to the positively charged protons, it can generate a magnetic field. Because of this, the protons will have two orientations, a β and α spin state that have different energies in the presence of an external magnetic field. When the compound is exposed to electromagnetic radiation, resonance (when the nuclei is flipped back and forth between orientations) will occur and when the amount of radiation matched the difference in energy of the two spin states, an NMR signal is recorded and the NMR detector records peaks.

One of the most common forms of NMR is the proton NMR (H⁺ NMR). The NMR spectrum is read by the chemical shift of the different groups. In a proton NMR a compound with an -OH group will have a chemical shift (peak) at around 0.5-5.0 ppm. For example, ethanol will have a peak at around 4.0 ppm indicating the hydroxide groups while its two other peaks at around 1.0 ppm and around 3.0 ppm will indicate the position of the C-H protons in accordance to their position in the compound.

References

1. Schore, Neil E. (2011). Organic Chemistry Structure and Function 6th edition. W. H. Freeman.

2. Hammond, Christina N., Mohrig, Jerry R., Schatz, Paul F.. (2010). Techniques in Organic Chemistry 3rd Edition. W.H. Freeman.

3. Atkins, Peter, Jones, Loretta. (2008). Chemical Principles: The Quest for Insight 4th Edition. W.H. Freeman.

General Information

A ketone. Carbonyl is in between two hydrocarbons.

An aldehyde. R can be a anything but a hydroxyl group or alkoxy group attached to a carbonyl.

A carbonyl group is a functional group that is comprised of a carbon atom doubly bonded to an oxygen atom. The carbon attached to the oxygen can have single bonds to different atoms. The atoms that the carbon is bound to distinguish it as a ketone, aldehyde, carboxylic acid, ester, or amide.

==The Nomenclature of Carbonyl Compounds==

1, Naming Carboxylic Acids:

The functional carboxylic acid is called a carboxyl group. In systematic nomenclature, a carboxylic acid is named by replacing the last alphabet "e" with "oic acid." For instance, a two carbon alkane is called ethane and two carboxylic acid is ethanoic acid. Notice that a carboxylic acid with 6 or less carbons are usually known by their common names. ( methanoic acid = formic acid, ethanoic acid = acetic acid, propanoic acid= propionic acid, butanoic acid= butyric acid, pentanoic acid= valeric acid, hexanoic acid= caproic acid)

2, Naming acyl halides

Acyl halides have a halide in place of the -OH group in a carboxylic acid. Acyl halides are named by replacing "-ic acid" in the acid name with "-yl halide" (eg; -yl chloride, -yl bromide, -yl flouride etc). If acid has a name with "-carboxylic acid," then replace "carboxylic acid with "-carbonyl halide" (e.g. -carbonyl chloride, -carbonyl bromide, -carbonyl fluoride)

3, Acid anhydrides

An acid anhydride is formed by two molecules of carboxylic acid reacting with each other to lose one water molecule. An anhydride is a symmetrical anhydride if the two reacting acid molecules are the same. If two reacting acid molecules are different then they are going to from a mixed anhydride. Symmetrical anhydrides are named by replacing "acid" in acid name with "anhydride." Mixed anhydrides are named by stating the names of both acids in alphabetical order then followed by "anhydride."

4, Naming Esters

If there is an -OR group in place of the -OH group of a carboxylic acid, the name of the group attached to the carboxyl oxygen should be

stated first, then state the name of acid by replacing "-ic acid" with "-ate."

Salts of carboxylic acids are named in the same way. The cation is named first then name the acid with replacing "-ic acid" by "-ate"
Notice that cyclic esters are called lactones. Their common names are derived from the common name of the carboxylic acid, which designates the length of the carbon chain, and a Greek letter to indicate the carbon to which the carboxyl oxygen is attached.

5, Naming Amides

An amide has an -NH2, -NHR, or -NR2 group in place of -OH group of a carboxylic acid. They are named by replacing "-oic acid", "-ic acid", or "-ylic acid" of the acid name with "-amide." If a substituent is bonded to the nitrogen, the name of the substituent should be stated first (if there is more than one substituent bonded to the nitrogen, they should be stated alphabetically), then state the name of the amide. The name of each substituent is preceded by capital N to indicate that the substituent is bonded to a nitrogen. Cyclic amides are called lactams. Their nomenclature is similar to that of lactones.

Carbonyl Compounds

Compound	Aldehyde	Ketone	Carboxylic Acid	Ester	Amide
Structure

Each of the characteristic specific to each compound is detailed on their specific pages.

General Reactivity

Oxygen is far more electronegative than carbon, so the electron density is higher near the oxygen and lower near the carbon. This creates a dipole moment, where the oxygen bears a negative charge and the carbon bears a positive charge of the same magnitude. This distribution of charges makes carbon an electrophile and oxygen, the resulting nucleophile.

The reactivity of each carbonyl compound is dependent on the group attached directly onto the alpha carbon. This is because of the resonance stabilized structures which form due to the donation of an electron by this group. The stronger the contribution of this resonance structure, the stronger the stability of the carbonyl. An example of this can be seen in carboxylic acids which upon deprotonation create a degenerate structure and thus increases the acidity of the proton. Peptide bonds are unusually strong and carry more of an $sp^2$, planar hybridization due to the resonance contribution of the amide electron onto the alpha carbon. This decreases the reactivity of peptide bonds considerably and thus requires much energy or a protease to catalyze the bond. Carbonyls by themselves are very stable bonds and the energy of their formation are usually very high. This makes the formation of carbonyls in organic synthesis to be highly thermodynamically favorable and usually the creation of a carbonyl bond as the end product will drive a reaction to formation.

The carbonyl group has a short, strong, and very polar double bond. Its reactivity of its double bond is very different from the double bond of the alkenes because of their oxygen's electronegativity along with the lone pair of electrons. The carbonyl carbon is also electron withdrawing since it is so close to the highly electronegative oxygen. The polarization of aldehydes and ketones also alters the physical constants. The polarization of the carbonyl group is the reason why their boiling points are higher than those of the hydrocarbons of similar molecular weight. the carbonyl group's polarity causes the smaller molecules to be completely miscible in water. Carbonyl compounds with more than six carbons are considered large, which is insoluble in solution. The larger the compound, the larger its hydrocarbon chain, the more hydrophobic the molecule is, so its solubility would then decrease.

Inductive Effect Take important notice that the electrophilicity of the carbon is highly dependent on the nearby atoms and the atoms it is directly bounded to. For example, the carbonyl carbon of a carboxylic acid will not be as electrophilic as a carbonyl carbon of a ketone because of resonance stabilization. On the other hand, acetyl chloride (the carbon is bound to an R-group, doubly bonded to the oxygen, and then bound to the fluorine), will have a more electrophilic carbonyl carbon because the electronegative chlorine will increase the induced positive dipole of the carbon. Electrophilicity plays a key role in chemical reactions, and less electrophilic carbonyl carbons are not as readily reactive.

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Biochemical Synthesis of Carbonyl Groups

Nicotinamide Adenine Dinucleotide

Within biological systems, carbonyl compounds can be formed by the oxidation of alcohols. An oxidation reaction in organic chemistry is one that is characterized by a process which either adds electronegative atoms or removes hydrogen from a molecule. An example of this oxidation includes the oxidation of ethanol by the oxidizing agent nicotinamide adenine dinucleotide (NAD⁺). NAD⁺ is composed of a pyridine ring, two ribose molecules, and the heterocycle adenine. When the two enantiometers of 1-Deuterioethanol are reacted with NAD⁺, in the presence of enzyme alcohol dehydrogenase, the biochemical oxidation is found to be stereospecific (NAD⁺ and only removes the hydrogen attached to the C₁ atom in 1-Deuterioethanol. Similar to the oxidation of 1-Deuterioethanol described above, other alcohols can be oxidized biochemically to form carbonyl groups. Another example includes the oxidation of methanol to formaldehyde.

'HNMR

In 'HNMR spectroscopy, aldehyde compounds have a very unique chemical shift which appears to be between 9-10 ppm. Which means that the aldehyde formyl hydrogen is also very strongly deshielded. Aldehyde C₂ hydrogens are also slightly deshielded because of the electron withdrawing from the oxygen from the carbonyl group. In ketones, this is also similar. The alpha-hydrogens also experience this deshielding; which has a chemical shift between 2-2.8 ppm.

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¹³CNMR

The carbon-13 NMR spectra for aldehydes and ketones are the same since it is of the chemical shift of the carbon participating in carbonyl group. Since that carbon is bound to an oxygen, it will appear at a lower field approximately 200 ppm. The carbons adjacent to the carbonyl carbon are also deshielded (just like the hydrogens in 'HNMR). The carbons further away from the carbonyl group are less deshielded.

IR Spectroscopy

The IR absorbtion bands for all carbonyl compounds absorb in the 1760-1665 cm^-1 region, which is due to the streching vibration of the C=O bond. Generally, carbonyl groups have a high intensity and narrow regions, making them useful for diagnostic purposes.

Range	Type of Compound	Such as:
1750-1735 cm-1	saturated aliphatic esters	CH3-CH2-COOR
1740-1720 cm-1	saturated aliphatic aldehydes	CH3-CH2-COH
1730-1715 cm-1	α, β-unsaturated esters	R'(R")C-CH-COOR
1715 cm-1	saturated aliphatic ketones	CH3-CH2-CO-CH2-CH3
1710-1665 cm-1	α, β-unsaturated aldehydes and ketones	R'(R")C-CH-COH

Introduction

A carboxyl group consists of a carbon double-bonded to an oxygen and also bonded to a -OH group. Compounds with carboxyl groups are called carboxylic acids or organic acids. The carboxyl group can act as an acid when by donating a proton (H⁺) to a solution and becoming ionized. Under biological conditions at pH~7, carboxyl groups are usually deprotonated, meaning they lose a H⁺, and become negatively charged. An example of a carboxyl group in the body would be carbonic acid, formed from the hydration of a carbon dioxide. Under biological conditions, carbonic acid usually dissociates into bicarbonate ion.

An ester functional group

Properties

Carboxyl groups have an electronegative oxygen atom double bonded to a carbon atom. This carbon-oxygen bond is very polar and the fact that its a double bond increases the polarity of the bond. As a result of the polarity, compounds containing carboxyl groups usually have higher melting points, boiling points and have hydrophilic centers. Moreover, the higher melting point and higher boiling point can be attributed to the fact that they can form hydrogen bonds both in the liquid and solid state. Fatty acids are examples of compounds that have hydrophilic centers due to their carboxyl groups. Also, carboxyl groups, especially when present in molecules with a low molecular weight tend to be highly volatile and therefore tend to have strong odors. The pKa of carboxyl groups usually range from 4-5.

Nomenclature

In naming organic molecules with multiple functional groups, the carboxyl group takes precedence in naming over any other functional group. Therefore when naming a molecule such as an alkane that contains a carboxyl group, the -e on the alkane is replaced by -oic acid. Also, when numbering the chain of the organic molecule that contains a carboxyl group, the carboxy carbon is labeled as the number 1 carbon. Molecules with two carboxyl groups would use instead the -dioic suffix.

Bonding

The polarity of the carbon-oxygen bond makes the carbon very susceptible to nucleophilic attack. Upon attack, the electrons of the double bond will migrate to the oxygen atom in order to maintain the octet for the carbon atom. The oxygen will now be negatively charged and a tetrahedral intermediate has been formed. The double bond will reform when the migrated electrons on the oxygen atom move back into the double bond to oxygen while the carbonyl carbon attacked expels the -OH group as a leaving group. While the expulsion of an -OH group is energetic unfavorable, the formation of the energetically favorable carbon-oxygen double bond helps overcomes this obstacle. Other ways to overcome this obstacle is to convert the -OH group into a better leaving group. The polarity giving the oxygen a partially negative charge also makes the carboxylic acid susceptible to electrophilic attack. An example of this is the hydrolysis of a carboxylic acid under acidic conditions where a proton acts as an electrophile and attacks at the oxygen which is doubly bonded to the carbon.

Amino Acids

The carboxyl group is a major component of amino acids. The carboxyl group, along with the amino group, allows amino acids to be zwitterions where both the amino group and the carboxyl group are charged. Since the carboxyl group can be deprotonated, it can impart a negative charge onto the amino acid. The carboxyl group is also key in the formation of peptide bonds. The carboxyl group of an amino acid can be attacked by the amino group of another amino acid. The nitrogen group of the amino group acts as the nucleophile and attacks the carbon of the carboxyl group. Carboxyl groups are also present on the side chains of two amino acids, Aspartate and Glutamate. These amino acids allow for hydrogen bonding and the formation of salt bridges, which help stabilize the structure of proteins.

From Nature

• Formic acid (HCOOH), the simplest carboxylic acid with only one carboxyl group, is primarily responsible for the pain caused by insects' bites(mostly Hymenoptera, like bees and ants).
• Acetic acid (CH₃COOH) can be biologically synthesized by either aerobic or anaerobic fermentation, a process used to make vinegar. The aerobic process requires warm ethanol (CH₃CH₂OH) and oxygen with Acetobacter. The anaerobic process requires only sugar(C₆H₁₂O₆) as input chemical, and acetogen can then give carboxylic acids as output. It should be noted that aerobic process is still dominantly applied, because acetogens used for anaerobic processes show less tolerance to acidic environments. In other words, acetogens will be killed if too much acid is produced.
• Propanoic acid (CH₃CH₂COOH,or C₂H₅COOH). It can be formed by breakdown of fatty acids with odd numbers of carbon atoms. During such metabolic process, propanoic acid undergoes condensation reaction with the thiol end of CoA (coenzyme A) to form propionyl-CoA. Propanoic acid can also be synthesized by anaerobic respiration of bacteria called propionybacterium.
• Butanoic Acid (C₃H₇COOH) can be found in some naturally occurring esters, such as hexyl butanoate from oil of Heracleum and octyl butanoate in parsnip. The fermentation (biological) method of butanoic acid production was discovered by Louis Pasteur in 1861. The overall process takes one mole of glucose and the products are one mole of butanoic acid, two moles of carbon dioxide, two moles of hygrogen gas and three moles of ATP. During such process, glucose is cleaved into two pyruvate molecules first. The pyruvate is then oxidized to acetyl CoA releasing carbon dioxide and hydrogen as by-products. The ATP is released after the acetyl CoA undergoes various enzyme reactions.
• Benzoic acid (C₆H₅COOH) is the main component of benzoin resin. However, it will be fairly pricey to extract benzoic acid directly from benzoin resins. Therefore most benzoic acid in the market is manufactured industrially.
• Citric acid (HOOCCH₂-COH(COOH)-CH₂COOH).The simple structure is similar to glycerol, which is also a biologically abundant molecule in fat. Citric acid is a well known organic acid in variety of citrus fruit such as lemon and lime. It is the essential material for citric acid cycle, a very important metabolic process. It is also very useful for modern pharmaceutical, cosmetics and other industries that include chemical processes. The industrial production of citric acid experienced a transformation from juice extraction to biosynthesis. Industrial citric acid production began in 1890 using extractions from Italian citrus exports, following the first crystallization of citric acid by Swedish chemist Carl Wilhelm Scheele in 1784. The biological methods of production was discovered in 1893 using Penicillium mold and sugar, but such process was not popular until World War I cut the Italian citrus exports. A more efficient biological production was then discovered by American food chemist James Currie by using cheap sugary mixture and mold called A. niger. Such industrial process has been used by major pharmaceutical companies like Pfizer. In addition, the extracted or biologically synthesized citric acid is precipitated by calcium hydroxide for isolation and acid is converted back from precipitate at the end.
• Oxalic acid (HOOC-COOH) is found in kidney stones as calcium oxalate, and it leads to kidney failure. It is therefore risky to eat carambola (commonly known as starfruit) and monstera due to their high oxalate content. Oxalate is also included in citric acid cycle.

Introduction

The mechanism of peptide bond formation.

Amino groups are composed of a N atom bonded to two H atoms. Amino groups can act as a base because they can pick up an H⁺ from a solution. Amino groups can be ionized with a 1⁺ charge under basic condition. Organic compounds with an amino group are called "amines"; organic compound containing an amino group and carboxyl group are called amino acids which are the building blocks of proteins.

Amines can readily form hydrogen bonds. The amines in the bases of DNA form hydrogen bonds with nearby nitrogen or oxygen atoms and keep the two strands together.

The amino group of lysine[[|]], for example, has be proven useful in the regulation of genes. In particular when the amino group of Lysine in the histones is acetylated, it can no longer function as it regularly does. This is a regulating step that is involved with gene expression and replication.

Another role amino groups play in biochemistry is in enzymes. In the case of a protease, an enzyme which cleaves amino acids, a tetrahedral transition state is formed when the hydroxy group of serine attacks the carbonyl carbon of the amino acid. Because the tetrahedral transition state has a negative charge, the positive NH₃⁺ charges help to stabilize the transition state, forming what is called an oxyanion ring.

Amines can also act as nucleophiles because of its lone pair of electrons. This is the basis by which peptide bonds are formed, with the carbonyl carbon acting as the electrophile in a dehydration reaction.

Structure and Physical Properties

Inversion of amine

The amine nitrogen is sp3-hybridized and tetrahedral shape. The nitrogen non-bonded pair acts like a substituent—the geometry is tetrahedral around nitrogen and the bond angles are all around 109°. Unlike carbon, sp3 hybridized nitrogen is not rigid and undergoes rapid inversion at room temperature. It is similar to an umbrella flipping inside-out and similar to the inversion of configuration which occurs in an Sn2 reaction.

Nomenclature

All the usual IUPAC rules are followed. The suffix –amine is added to the name at the end and the position of the amino group must be specified. For 2° and 3° amines, the largest alkyl substituent is chosen as the parent and the other alkyl groups on nitrogen are named as substituents with the prefix N- to denote that they are attached to nitrogen. Because the amine functional group has the lowest priority in naming, it is often named as a substituent on more highly functionalized molecules.

Acidity of Amines

Amines are much less acidic than alcohols with a pKa ~ 35 and Keq = 10-35. A very strong base like alkyllithium must be used to completely deprotonate an amine.

Basicity of Amines

Amines are the most basic of the common organic functional groups, but are still fairly weak bases. Protonation occurs on the non-bonded electron pair exclusively. The basicity of amines is directly dependent on the “electron density” at the nitrogen atom. Both inductive and resonance effects can alter the basicity of a nitrogen atom.

Hybridization on the N also affects basicity. An increase in s character on an atom increases the electronegativity of that atom which favors acidity and therefore disfavors basicity. Hence sp³-hybridized nitrogen is more basic than either sp² or sp hybridized nitrogen.

Synthesis of Amines

Gabriel synthesis to 1° amines

1. Cyanide displacement of alkyl halides followed by reduction

2. LiAlH4 reduction of amides

3. Displacements with azide followed by reduction

4. Gabriel synthesis to 1° amines

5. Reductive amination

IR Spectroscopy for Amines

Primary Amines will give two short, sharp equal peaks at around 3200 cm-1 – 3500 cm-1

Secondary Amines will give one short, sharp peak at around 3320 cm-1

Tertiary amines will not give a peak in any region.

Reference

1. http://orgchem.colorado.edu/Spectroscopy/irtutor/aminesir.html

General Simple Amide

Amide with -R Groups on Nitrogen

An amide functional group consists of a carbonyl group bonded to a nitrogen. In simple amides, two hydrogen atoms are bonded to the nitrogen (-CONH2) while in more complex amides, the nitrogen is bonded to one or two aliphatic or aromatic groups (-CONR).

Nomenclature

Naming amides is very similar to naming carboxylic acids. For IUPAC nomenclature, first name the carboxylic acid, and then drop the -oic acid and add amide. For example, propanoic acid would become propanamide and ethanoic acid would become ethanamide. The first part of the name depends on the carbon chain the amide is attached to.

Properties

Basicity

Differing from similar amines, amides show no measurable basicity due to the delocalised lone pairs on the nitrogen. In normal amines or -NH2 functional groups, the lone pair on the nitrogen can accept hydrogen atoms acting as a base. However, the pi bond formed between the double bonded carbon and oxygen in an amide contains p orbitals that are positioned almost parallel to the lone pair on the nitrogen. This causes the electron pair to be delocalised and shared throughout the carbonyl part of the molecule. Delocalization reduces an amide's basicity because the electron pair is not associated with a single atom reducing the intensity and focus of its proton drawing ability. Delocalization also helps to stabilize the overall structure of the amide and as a result more energy would be required to break the shared electron structure.

Melting Point

Methanamide is a liquid at room temperature while other amides remain solid. Relative to their size, amides have relatively high melting points due to the hydrogen bonding between the partially positive hydrogen atoms the -NH2 group and another electronegative oxygen. Each simple amide has two partially positive hydrogen atoms and two pairs of electrons on the oxygen allowing for multiple possible sites of hydrogen bonding. A lot of energy is required to break these hydrogen bonds increasing the melting point of amides.

Solubility

Small amides are soluble in water because they may have hydrogens bond with water molecules. Larger amides have trouble dissolving because of their long hydrophobic carbon chains. Amides are typically less soluble than amines and carboxylic acids because they can both donate and accept hydrogen bonds.

Preparation

Carboxylic acids can be used to prepare amides by reaction with solid ammonium carbonate in acid to form a ammonium salt. Upon heating, this salt dehydrates to produce and amide and water. Acyl chlorides (acid chlorides) will react violently ammonia to create ammonium chloride and an amide of the acyl chloride. Acid anhydrides will also react with ammonia to produce amides such as the reaction of ethanoic anhydride with ammonia producing ethanamide and ammonium ethanoate.

Synthesis of Amides

Amides are derived from a reaction between an amine and a carboxylic acid. Between these two molecules we have two competing nucleophiles, the oxygen of the alcohol group in the carboxylic acid molecule and the nitrogen of the amine. A nucleophile is a chemical species that donates a pair of electrons to an electrophile to create a chemical bond in a reaction. With nitrogen lying to the left of the oxygen on the periodic table, nitrogen serves as a better base and better nucleophile than the alcohol. The reaction between an amine and a carboxylic acid is based on an addition and elimination reaction. Although this is a simple and easy reaction, it is not the most effective and efficient way of producing amides. The reaction between the two also contains a competing acid-base reaction, which produces a salt. Therefore, with competing products, the addition-elimination reaction between the two is not the most effective way to isolate an amide. A better procedure would be the reaction between an acyl halide, an activated carboxylic acid derivative, and an amine. The replacement of the hydroxyl group in the carboxylic acid with a halide produces a reactive molecule called an acyl halide. With halogens being the most electronegative atoms, the presence of it in the molecule pulls the electrons away from the carbon of the carbonyl atom, creating an electrophilic site. With an electrophilic site present, the nucleophilic nitrogen of the amine will easily react with the acyl halide to form an amide.

938 × 212px

Step 1: The nucleophilic nitrogen attacks the carbon of the carbonyl, pushing the electrons of the double bond of the carbonyl to the oxygen. Formation of a zwitterion occurs (negative charge on the oxygen, positive on the nitrogen).

Step 2: The favored and more stable carbonyl is reformed, kicking out the halide.

Step 3: The positive charge is quenched as the halide comes back and removes a hydrogen from the nitrogen, forming an amide and a hydrogen halide.

Reactions

Hydrolysis

Hydrolysis of amides can occur under both acidic and basic conditions. Under acidic conditions, amides catalyzed by dilute acid react with water to form carboxylic acids and ammonium chlorides. An example would be heating ethanamide in dilute hydrochloric acid to form ethanoic acid and ammonium chloride. If heated under basic conditions such as sodium hydroxide solution, ethanamide will form ammonium gas and sodium ethanoate salt.

Dehydration and Hofmann Degradation

Amides can be dehydrated by reaction with phosphorous (V) oxide such as heating ethanamide with phosphorous oxide to produce ethanenitrile with the loss of water. The Hofmann Degradation reaction involves reaction of an amide with a mixture of bromine and sodium hydroxide resulting in the loss of the carbonyl group such as degradation of ethanamide into methylamine.

Practical use of polyamides

Chemical Structure of Kevlar

Polyamides are in general polymers held together by amide links. Nylon consists of repeating chains of carbons held together by amide chains while kevlar is made up of chains of benzene instead of carbon. Nylon is formed from the loss of water between a reaction of hexanedioic acid and diaminohexane while kevlar is formed from the reaction of benzene dicarboxylic acid and diaminobenzene. Nylon is used commercially for clothing, carpets, ropes, and tires while the high strength to weigh ratio of kevlar makes it practical for use in bullet proof vests and other lightweight sturdy needs.

General Overview

A sulfhydryl is a functional group consisting of a sulfur bonded to a hydrogen atom. The sulfhydryl group, also called a thiol, is indicated in chemistry nomenclature by "-thiol" as a suffix and "mercapto-" or "sulfanyl" as a prefix. Thiols have great affinity for soft metals. Sulfhydryls play an important role in biochemistry, as disulfide bonds connect necessary amino acids together for functional purpose in secondary, tertiary, or quaternary proteins structures.

Several Important Roles in Biochemistry

Sulfhydryl groups can be found in the amino acid cysteine. When two cysteine residues are in close proximity to each other, they can form a disulfide bridge also called cystine. The formation of a disulfide bond is an example of a post translational modification. It can be helpful to the structure of proteins, but it can make it difficult to accurately determine the sequence of a protein through the technique of Edman Sequencing. Disulfide bridges often play important structural and functional roles in proteins. Their hold are crucial in the formation of many tertiary structures of proteins. Extracellular proteins usually contain disulfide bonds whereas intracellular proteins usually lack disulfide bonds. These disulfide bonds can be broken with the addition of beta mercaptoethanol. Beta mercaptoethanol reduces the disulfide bonds back into their sulfhydryl form. It was found by Christian Anfinsen that proteins denatured by beta mercaptoethanol in urea will spontaneously reform its disulfide bonds if trace amounts of beta mercaptoethanol are present. This reformed protein was found to be fully functional. If the protein was reformed in the absence of urea, it was found that the function was greatly reduced. This is due to the fact that the wrong disulfide bonds were formed and the amino acid became "scrambled."

Structure

The C-S bond in thiols are 180 picometers in length, and the C-S-H bonds are at a 90° angle, significantly more acute than alcohols. Their main cohesive forces are van der Waals interactions, and sulfur is less electronegative than the oxygen in hydroxyl groups, making them less polar.

Reactions

Thiols can be easily oxidized, and thiolates act as potent neucleophiles.

S-Alkylation

The conjugate bases of thiols can be alkylated to give thiolethers.

RSH + R'Br + base → RSR' + [Hbase]Br

Redox

Thiols, in the presence of a base, are oxidized to give an organic disulfide.

2 R–SH + Br₂ → R–S–S–R + 2 HBr

Thiols can be oxidized by more powerful reagents such as sodium hypochlorite to yield sulfonic acids.

R–SH + 3H₂O₂ → RSO₃H + 3H₂O

Thiols can also undergo thiol-disulfide exchange.

RS–SR + 2 R'SH → 2 RSH + R'S–SR'

Organolithium compounds can react with sulfur to make thiols.

RLi + S → RSLi RSLi + HCl → RSH + LiCl

References

1. Berg, Jeremy M. (2007). Biochemistry, 6th Ed., Sara Tenney. ISBN0-7167-8724-5. 2. Campbell, Neil A. Biology. 7th ed. San Francisco, 2005. 3. E. Jones and I. M. Moodie (1990), "2-Thiophenethiol", Org. Synth.; Coll. Vol. 6: 979

Introduction

A phosphate group consists of a phosphorus atom bonded to four oxygen atoms. It is usually ionized and attached to the carbon skeleton by one of its oxygen atoms. Compounds with phosphate groups are called organic phosphates and they are frequently involved in energy transfer reactions. Phosphates can also be found in the backbone of DNA forming phosphodiester bonds (two esters with a common posphonyl). The negative charges on phosphates are part of the reason why the backbone of DNA is on the outside. A good example of transferring energy is the three phosphate group which is found in the energy compound ATP(adenosine triphosphate).

Phosphorylation of proteins is important for several reasons. Phosphate has a charge of negative two, which is important for disrupting or forming electrostatic interactions that alter the structure and function of proteins. Phosphate is also capable of forming three hydrogen bonds. The free energy of phosphorylation is large allowing it to change the equilibrium between conformational states. Phosphorylation can also take place very quickly or very slowly making it very flexible and versatile in meeting the needs of the body. Phosphorylation also results in highly amplified effects where hundreds of target proteins are phosphorylated by one activated kinase in a short time span. Phosphorylation is prevalently seen with the application of ATP, which aids in the regulation of metabolism.

Properties

A phosphate group is a hypervalent molecule containing five bonds which makes ten electrons. Phosphate salts can form through ionic bonds between a cation and one of the oxygen anions. Molecules with a phosphate group are usually not very soluble in water. Phosphate can exist in four different forms. In acidic conditions, it can take the form of phosphoric acid while in the more basic conditions it can be fully deprotonated to the phosphate ion. The forms phosphoric acid, dihydrogen phosphate ion, hydrogen phosphate ion, and phosphate ion all act and behave as individual weak acids. Each pK value differs greater than 4. This creates three pH regions of very acidic, moderately acidic to moderately basic, and very basic, depending on which deprotonation stage. These characteristics give the phosphate group the functionality of being a good buffer. Polymeric ions such as pyrophosphate and metaphosphates can be formed. When a pyrophosphate binds to a calcium ion, calcium phosphate is formed. Phosphate is and important component to the strength and solidity of animal teeth and bones. The exoskeleton of crustaceans and insects contain calcium phosphate as well.

Phosphate Applications

Phosphate can be used chemically as polish metal alloys. Phosphoric acid polishers work by reacting with various metal ions and generating hydrogen gas thus releasing some of the metal ions. By doing so, the metal is "phosphatized" with a phosphate coating that prevents corrosion and increases paint adhesion. This also plays a role in producing and maintaining a system of drinkable water because water usually runs through pipes and other water distribution sources that can easily erode. Because phosphate acts as a good buffer, it ensures that the pH of drinking water is stable and resists dramatic changes.

Introduction

Organic Synthesis is the part of Organic Chemistry that deals with the creation of compounds from other available compounds. There are lots of different ways to make substances and synthesis studies this. There are four very basic reaction schemes that can help in seeing what synthesis consists of: S_N1 (Unimolecular Nucleophilic Substitution), S_N2 (Bimolecular Nucleophilic Substitution), E1 (Unimolecular Elimination), and E2 (Bimolecular Elimination).

S_N2

As the name implies, S_N2 reactions deal with the substitution of some group for another. There are certain things the group coming in and the group leaving need to have in order for the reaction to happen because, for example if the leaving group cannot leave or is a group that does not want to leave, the reaction will not happen. The group coming in is referred to as the nucleophile. The nucleophile has electrons that can be donated to some atom lacking electron density, the electrophile. in an substitution reaction, the electrophile has the leaving group attached. To be considered a good nucleophile, the group needs to have lots of electron density usually being charged, and are not able to balance these extra electrons. There are several ways a molecule can balance its electron density. Three of these ways are: resonance, atomic radius, inductance. Of these, resonance is the biggest factor although atomic radius also plays a big role. inductance could be important but it is a balancing force that loses its power rather quickly. Resonance refers to the phenomenon where a molecule oscillates between two different states, usually by moving around double bonds. This makes it so that charges move around in a molecule. This moving of the charge is more of an oversimplified way of seeing this. In reality, the molecule exist in a hybrid state of all the different resonance forms. This hybrid state actually spreads the electron density all around the molecule.

The resonance of benzene and the hybrid state it actually exists in

The resonance of acetate. The charge is spread between the two oxygens and the connecting carbon and the charge is better balanced

Atomic radius plays a role in stabilizing charge. For example, all the halogens have a -1 ionic form. The charge is the same for all of them. What changes is the atomic radius. a larger radius means the electron density can be dispersed more, thus balancing the charge better.

atomic (gray) and ionic (red and blue) radii of some elements

Inductance is the ability of electronegative atoms to pull electron density towards themselves and can stabilize charge in that way. This force dies out rather quickly though as the electronegative moves farther away from the charge.

A molecule that can donate electrons is called a Lewis base and indeed most nucleophiles are basic, although a good base is not necessarily a good nucleophile. A molecule that can withstand a negative charge (by using what was mentioned above) could be a good leaving group.

In the molecule of bromomethane, the leaving group is the bromo group and the electrophile is the carbon. this bromo group is very electron withdrawing and because it has a big atomic radius it can dissipate the charge, a term called polarizability. A typical S_N2 reaction involves this bromomethane and NaOH

After the reaction has occurred the stereochemistry of the new product will be the exact opposite than before. For example if it was an (R)-iodopropane reacting with bromine it would because (S)-bromopropane. The reason for this inversion of the stereocenter is because of the backside attack the nucleophile does in Sn2 reactions.

Bromomethane molecule. Bromine serves as the leaving group

Bimolecular Nucleophilic Substitution of Bromomethane

Ways to increase the rate of the reaction
1. Increasing the concentration of either the nucleophile or substrate will increase the rate of the reaction. This is because it is a bimolecular reaction, where rate = k[nucleophile][substrate]: the rate is dependent on the concentrations of nucleophile and substrate. Increasing the concentration of nucleophile will increase the rate, since rate is proportional to the concentration of nucleophile and concentration of substrate. The same applies to the substrate.

2. Using a good leaving group.

3. Increasing the temperature of the system.

4. Anything that will not sterically hinder the nucleophile will increase the rate of the reaction.

5. Using a polar aprotic solvent, meaning the solvent does not contain an acidic proton.

S_N1

The reason for the 2 in and S_N2 reaction is the fact that the leaving group and the nucleophile are both in the rate law of the reaction. This means that the concentration of both bromomethane and NaOH affect the rate of the reaction. This is because the reaction happens in one step with one intermediate where the leaving group-electrophile bond is breaking and the nucleophile-electrophile bond is forming. In an S_N1 reaction there are two steps and two intermediates. The first is the formation of a carbocation where the leaving group leaves and the carbon is left with a positive charge. The second is the formation of the nucleophile-electrophile interaction. Of these two The first is the slowest and called the rate determining step. This means that no matter how much nucleophile there is the rate of reaction will not changed until the concentration of the substrate changes. The first step is very slow because forming a carbocation is not very favorable. The only way for this to happen is if the leaving group is a really good one and for there to be a way to stabilize the carbocation. The first one is simple, find a molecule with a very good leaving group. By making the carbocation a secondary or tertiary carbocation, the charge is stabilized better. This is done by resonance or hyperconjugation. The electrons in the methyl group(s) interact with the carbocation and give it some stabilization.

The nucleophile can attack the carbocation from both faces, the alpha or beta face of the substrate. As a result one will generally see a racemic mixture, or 50:50 mixture of a R and S conformation of the resulting compound.

Formation of carbocation stabilized by the ethyl groups

In order for the slow first step to happen, the second faster step needs to be slowed down. The way to do this is to have not have a very good nucleophile. instead of the hydroxide ion, water its conjugate acid can be used.

The first step in the unimolecular substitution reaction is the departure of the leaving group. the second is the attack of a weak nucleophile

Ways to increase the rate of an Sn1 reaction
1. Another way to increase the rate of the reaction for Sn1 reactions is to use a polar protic solvent. The polar protic solvent will increase the rate of the reaction because it will help stabilize the leaving group. It will stabilize the leaving group because it will solvate the charge. Anything that can help stabilize the leaving group will make it more likely to leave because it will be going to a lower energy state.

2. Another way to increase the rate of the reaction for Sn1 reactions is to add a salt. The salt will help stabilize the leaving group through ionic interactions, and because the leaving group would be more stable it is more likely to leave and form the carbocation that much faster.

3. Another way to increase the rate of the reaction for Sn1 reaction is to increase the polarity of the solvent. The polarity of the solvent will help stabilize the carbocation by hydrogen bonding. Therefore if one wants to increase the rate of reaction of an Sn1 reaction he or she simply needs to increase the polarity of the solvent used.

4. Another way to increase the rate of the reaction is to increase the temperature of the system.

5. Using a polar, protic solvent, meaning the solvent has acidic hydrogens.

Beta branching effect on the rate of the reaction
Beta branching means that there is a carbon chain on the carbon next to the carbon that is bonded to the leaving group. . For SN2 reactions beta branching will slow down the reaction because of steric hindrance, thus the less beta branching the alkyl halide has the faster the reaction will be for SN2. Based on the experiment results in SN1 reactions more beta branching will slow down the Sn1 reactions because of steric hindrance that hinders the nucleophilic attack on the carbocation.

Ring effect
Experiments has shown that the bromocyclopentane reacted instaneously while the bromocyclohexane required some time for the reaction to proceed showing that 5 membered ring reacts faster than a 6 member ring. The reason behind this is because of transannular strain, the unfavorable interactions between ring substituents on non-adjacent carbon, the less transannular strain there are the slower the reaction will be for SN1 and SN2. A 5-member ring is the only exception to this rule, but in general 3, 4, 6 member ring will react slowly because it is more stable and low in energy.

Aromatic Ring Effect
Aromatic rings will enable primary alkyl halide to proceed SN1 reactions because the aromatic ring can stabilize the carbocation through resonance. The position of the aromatic ring has to be one carbon away from the leaving group so the electrons can resonate and stabilize the carbocation. Electon-donating groups will contribute to the electron density so it will “activate” the ring and make it more susceptible to electrophilic attack, increasing the rate of reactions, while electron-withdrawing group will decrease the rate of reactions because it removes electron density from the aromatic ring, as shown with experiments the methoxybenzyl chloride, an electron donating group, reacted, while the nitro, an electron withdrawing group, did not react at all.

Reference

1. Vollhardt, K. Peter C., and Neil Eric Schore. Organic Chemistry: Structure and Function. New York: W.H. Freeman, 2011. Print.

Macromolecules

They are polymers made of preexisting monomers. They can form functional units by connecting together. Forming these functional unites requires energy. Examples of macromolecules include proteins, nucleic acid, lipids and polysaccharides.

References

1 ^ "http://en.wikipedia.org/wiki/Macromolecule," 2009.

2 ^ Nelson, D.L., "Lehninger Principles of Biochemistry," 2008.

   Macromolecules consist of proteins, nucleic acids, carbohydrates and lipids. 

1) Proteins are made of smaller building blocks called amino acids. There are 20 different amino acids. The chains of amino acids fold up in complex ways, giving each protein a unique 3D shape. 2) Nucleic acids allow organisms to transfer genetic information from one generation to the next. 3) Carbohydrates have several roles in living organisms. Carbohydrate derivates are actively involved in fertilization, immune systems, the development of disease, blood clotting and development. 4) Lipids include fats, phospholipids and steroid. Lipids also have function in the body.

Background Information

Carbohydrates consist of numerous functions that are important to living organisms. They are also known as saccharides, or sugar if they exist in small quantities; these names are used interchangeably to describe the same thing. The simplest carbohydrates are the monosaccharides, also known as simple sugars. Disaccharides are double sugars, consisting of two monosaccharides joined by a covalent bond. Carbohydrates also include polysaccharides, which are polymers composed of many sugar building blocks. The name "carbohydrate" is derived from 'hydrates of carbon', and they arise from photosynthesis, where they exist as products.

Carbohydrates are the most abundant aldehyde compounds found in living organisms. They provide storage, transport starch and glycogen that provide energy to bodies, and contain structural components such as cellulose in plants and chitin in animals. Additionally, they contribute to the immune system, fertilization, pathogenesis, blood clotting, and development.

Structure of Carbohydrates

Examples of Carbohydrates.

The common chemical formula for carbohydrates is C_n(H₂O)_n, where the ratios are usually 1 Carbon: 2 Hydrogens: 1 Oxygen. Trioses, pentoses, and hexoses are found most commonly among monosaccharides.Their structure is composed of the functional groups, aldehyde and ketone, which are attached with various amount of hydroxylgroups. The hydroxyl groups are usually attached to the carbons not a part of the aldehyde or ketone functional groups, to form aldoses and ketosesStructural Biochemistry/Carbohydrates/Ketoses, respectively. The most elementary carbohydrates comprise polyhydroxyaldehydes (an aldehyde moiety) or polyhydroxyketones (a ketone moiety). There are four general classes of carbohydrates: monosaccharides, disaccharides, oligosaccharides, and polysaccharides.

Classification and Nomenclature

The most important carbohydrate is glucose. In general, monosaccharides have one carbonyl group (aldehyde, ketone, or acid), and the remaining carbons each bear one hydroxyl group. Monosaccharides can be linked together via ether and/or acetal bonds to form very large polymers called polysaccharides. A disaccharide consists of 2 linked monosaccharides and so on. Almost all saccharides in nature have at least one chiral carbon and they occur in nature as a single enantiomer. Glucose has 4 chiral carbons and has 15 other stereoisomers for a total of 16 possible stereoisomers of this gross structural formula.

The suffix –ose is often used in describing and naming carbohydrates. For example:

• A carbohydrate with 6 carbons is called a hexose
• A carbohydrate with 5 carbons is called a pentose
• A carbohydrate with an aldehyde as its carbonyl unit is called an aldose
• A carbohydrate with a ketone as its carbonyl unit is called a ketose

Looking at glyceraldehyde:

Enantiomers behave identically whether its a D or L conformation they both have the same boiling point, melting point and solubility. A different conformation seen in carbohydrates are diastereomers. Diasteriomers have the same chemical formula but different connectivity. A monosaccharide that has diastereomer conformation have different chemical and physical property. the D/l configuration applies to the highest numbered stereocenter (in most cases the highest stereocenter is the second to last hydroxyl group in a sugar molecule). The D/L configuration, like chiral molecules, means that the molecules will gradually rotate under polarization direction of linearly polarized light as it passes through, even under solution. D/L configuration also can be used in nomenclature to distinguish carbohydrate stereoisomers (meaning they have the same physical and chemical properties as well as the same formula, but rotate differently in polarized light) with the same formula. Naturally-occurring glyceralde has an R-configured chiral carbon. This is called D-glyceraldehyde. The opposite enantiomer is called L-glyceraldehyde. The D and the L refer to the configuration of the highest numbered chiral carbon when viewed in Fischer configuration. D-has the hydroxyl on the right side and the L has the hydroxyl on the left side. Almost all naturally occurring carbohydrates are of D-configuration.

Monosaccharides

The most basic carbohydrate is the monosaccharides (e.g. glucose, fructose, and galactose with the structural formula C₆H₁₂O₆), and it consists of 3 carbons or more within the molecule. Monosaccharides may appear to be linear molecules, but when they come in contact with aqueous solutions, they tend to form 5-carbon ring structures, which makes the molecules more stable. Monosaccharides are classified by their placement of the carbonyl group and its chirality. If the carbonyl group is a ketone, the monosaccharide is referred to as a ketose. However, if the carbonyl is an aldehyde, the monosaccharide is called a aldose. There are more possible configurations of the aldose form than ketose due to the presence of more chiral carbons found in aldoses. Carbons that have a hydroxyl group (-OH), disregarding both the first and last carbons are asymmetric. The asymmetric carbons lead to the two possible forms (R and S)which corresponds to the D- and L- configurations. Combining two monosaccharides together would result in a disaccharide, linked via a glycosidic bond; and condensation reaction is the process that fuses two monosaccharides together. Such reactions form a disaccharide by removing a hydroxyl group from one monosaccharideMonosaccharides and a proton from the other.

Ring Structure of Monosaccharides After hemiacetals and hemiketals form, the carbohydrate will form a ring structure. For example, in Glucose the Hydroxyl group from the Carbon 6 will attack the carbonyl Carbon from Carbon 1, with a hemiacetal intermediate. The Carbon that will then change from a carbonyl carbon to a carbon with a hydroxy will then be called an anomeric carbon. An anomeric carbon is the hemiacetal or hemiketal that is bounded by an alcohol group and an ester bond. Anomers are another form of isomers that differ in the hemiacetals or hemiketals.The most common structure of a glucose ring is the alpha conformation in contrast with the sterically hindered Beta conformation. The alpha conformation is structure in which the hydroxyl in Carbon 1 is facing the opposite of the plane as carbon 6, while the beta conformation is facing in the same plane as the Carbon 6 thus causing steric hinderence. For this reason of steric hinderance the alpha conformation is much more stable than beta.

Monosaccharide Streoisomer

Monosaccharide stereoisomers

Modified Monsaccharide

Modified Monosaccharide example

Disaccharides and Polysaccharides

Monosaccharides not only form disaccharides, but polysaccharides as well. It is called an oligosaccharides if monosaccharides are linked by O-glycosidc bonds. The only difference between disaccharides and polysaccharides is that there are more monosaccharides combined together into a long chain, whereas disaccharides only consist of 2 combined monosaccharides. Polysaccharides are frequently long chains of glucose monomers bonded together. There are two types of polysaccharides: Homo-polysaccharides and Hetero-polysaccharides. Homo-polysaccharides are chains of one type of monosaccharides, while Hetero-polysaccharides consist of multiple types of monosaccharides. Both types of polysaccharides can exist in either branched or unbranched forms. Some crucial homo-polysaccharides, like starch, glycogen and dextrans, play a role in energy storage, while other homo-polysaccharides like cellulose and chitin have more structure-based roles. The diversity of these branched and unbranched carbohydrates is due to the number of hydroxyl groups present in the sugar. Any one of these hydroxyl groups can act as the alcohol in the formation of the glycosidic linkage. One sugar acts as the alcohol while the other has the anomeric carbon that ca for the O-glycosidic bond. This reaction can occur 1,4 or 1,6 depending on the alpha (pointing down) and beta (pointing up) orientations of the hydroxyl groups. These orientations organize the sugars into different structures. Alpha and Beta sheets form with sugars arranged with the least amount of steric hindrance.

Monosaccharides and disaccharides seem to be stable compounds, but that is not necessarily true, since they are hemiacetals with reactive carbonyls in their structure. Hemiacetals, with the general formula R1R'1C(OH)OR2 where R2 is not a hydrogen, and is formed by the reaction of carbonyl compounds with alcohols, and the carbonyl group is fairly reactive. Because the carbonyl groups are very reactive, they can oxidize to products in a short period of time. However, some carbohydrates are acetals and, as such, prevent carbohydrates from oxidizing. This occurs because the anomeric carbon is fixed in a glycosidic linkage. Because of this resistance against oxidation, acetals are known as the non-reducing sugars, like sucrose, while the hemiacetals are known as the reducing sugars, such as glucose, maltose, and lactose.

Two or more monosaccharides are linked by a glycosidic bond. The links can be alpha or beta depending on the position of the bond formed. Monosaccharides can also form bonds with amino acids to form glycoproteins. A sugar linked to an asparagine is N-linked, or linked from the sugar to the nitrogen of ASN. A sugar attached to Serine or Threonine is O-linked, or linked from the carbon on the sugar to the oxygen in those two amino acids. An Asparagine residue can only accept an oligosaccharides if the residue is a part of an Asn-X-Ser or Asn-x-Thr sequence, where x can be anytype of amino acid. Therefore, potential site can be detected within amino acid sequences.

Glucogen Metabolism Glucose metabolism and various forms of it in the process is described by the process below. Glucose-containing compounds are digested and taken up by the body in the intestines, including starch, glycogen, disaccharides and as monosaccharide. Glucose is stored in mainly the liver and muscles as glycogen. It is distributed and utilized in tissues as free glucose.

Examples of glucogen metabolism.

Hemiacetal and Hemiketal

An aldehyde or a ketone can react with an alcohol to yield a hemiacetal or a hemiketal.Hemiacetals and hemiketals are compounds that are derived from aldehydes and ketones respectively. The Greek word hèmi means half. These compounds are formed by formal addition of an alcohol to the carbonyl group. When the alcohol group is replaced by a second alkoxy group, an acetal or a ketal, respectively, is formed. For example, the acetal formation ends with having two ethers on the target carbon, whereas the hemiacetal and hemiketal has both an alcohol and ether group.

Example of Hemiacetal and Hemiketal Formation

1- Hemiacetal Formation.

2- Hemiketal formation.

In Carbohydrates

The same way aldehydes and ketones react with alcohols to form hemiacetals and hemiketals, respectively, carbohydrates react intermolecularly to form rings. When forming a ring 5 or 6 membered ring is most favorable and will only be formed. The Carbon 1 will be attacked by either the Carbon 5 or Carbon 6 hydroxyl group to form a 5 or 6 membered (respectively)carbohydrate ring.

The carbohydrates are a major source of metabolic energy, both for plants and for animals that depend on plants for food. Aside from the sugars and starches that meet this vital nutritional role, carbohydrates also serve as a structural material (cellulose), a component of the energy transport compound ATP, recognition sites on cell surfaces, and one of three essential components of DNA and RNA. Carbohydrates are called saccharides or, if they are relatively small, sugars.

References

Berg, Biochemistry, 6th Edition

http://themedicalbiochemistrypage.org/carbohydrates.html

http://faculty.clintoncc.suny.edu/faculty/michael.gregory/files/bio%20101/bio%20101%20lectures/biochemistry/biochemi.htm

Viadiu, Hector. "Carbohydrates." Chem 114A. UCSD, La Jolla. 19 Nov. 2012. Lecture.

Structure and Terminology of Nucleic Acids

Deoxyribonucleic acid (DNA) and Ribonucleic acid (RNA) are the two types of nucleic acids. The structure of DNA differs from RNA only because the DNA molecule does not have a hydroxyl group on the 2' carbon atom of the sugar ring. The missing hydroxyl group keeps DNA from being hydrolyzed, making it more stable than the ribose sugar of RNA.

The sugars in nucleic acids are bonded together by a phosphodiester bond between the 3' carbon on one sugar and the 5' carbon on another sugar.

Deoxyribose sugar

Ribose Sugar

The phosphodiester bond is negatively charged to repel nucleophilic attack by hydroxide ions and is less susceptible to hydrolytic attack. This negative charge, as well as the lack of a hydroxide group on the 2' carbon of DNA, makes the DNA molecule more stable than the RNA molecule, which could explain why DNA is the hereditary information carrier of all cells.

The chain of sugars made by the phosphodiester bonds form the backbone of the nucleic acids. The backbone of every nucleic acid is the same, but each sugar can contain one of four different bases. Two of the bases, (A)adenine and guanine (G), are derivatives of purine. The other two bases, cytosine (C) and thymine (T) are derivatives of pyrimidine.

Ribose, unlike deoxyribose, contains a base called uracil (U) instead of the base thymine (T). Thymine has an extra methyl group on the 5-carbon atom, making it slightly different than Uracil.

A nucleoside is a combination of a sugar and a base.

The nucleosides in DNA are called deoxyadenosine, deoxyguanosine, deoxycytidine, and thymidine, and the nucleosides in RNA are called adenosine, guanosine, cytidine, and uridine. If the base is a purine, then the N-9 (nitrogen) is bonded to the C-1' (carbon) of the sugar. If the base is a pyrimidine, then the N-1 is bonded to the C-1' of the sugar.

A nucleotide is a nucleoside joined to one or more phosphate groups by ester linkages. Therefore, it consists of the sugar, base, and phosphate(s). Nucleotides are the monomers that form DNA and RNA strands.

Most nucleotides are formed by the attachment of a phosphate group to the C-5' of the sugar. However, some nucleotides are also formed by the attachment of the phosphate group to the C-3' of the sugar. The four nucleotides that compose DNA are called deoxyadenylate, deoxyguanylate, deoxycytidylate, and thymidylate.

A DNA chain has polarity. One end of the chain will have a free 5' OH group. The other end will have a free 3' OH group. By convention, the base sequence of a DNA strand is written from the 5' to the 3' end.

The three-dimensional structure of DNA is a double helix. The two helices are coiled around a common axis and they run in opposite directions. The hydrophilic backbone of the helices make the outside of the helix, while the hydrophobic bases make up the inside of the helix. The helical structure of DNA is stabilized by the hydrogen bonding between complementary base pairs. Purines are larger than pyrimidines, so in order to have a regular structure, base pairs are formed between one purine and one pyrimidine. The base pairs of DNA are adenine (A) with thymine (T) and guanine (G) with cytosine (C). In RNA the thymine is replaced with uracil (U). The DNA structure is further stabilized by van der Waals forces between stacks of base pairs and the rigid five-membered ring structure of the backbone sugars.

RNA

RNA is nucleic acid, and its single-stranded, helical structure is constructed by nucleotides of nitrogenous bases, a ribose sugar, and phosphate group(s); the bases that make up RNA are adenine, guanine, cytosine, and uracil, for which, 1’ nitrogen of pyrimidine base and 9’ nitrogen of purines base are bonded to 1’carbon of pentose sugar by glycosidic bond; base pairs of adenine and uracil are hydrogen bonded to cytosine and guanine, respectively; the ribose is a pentose sugar of carbon numbered from 1’ to 5’ and has a hydroxyl group on the 2’ carbon; the 3’ and 5’ carbons of ribose sugar are bonded to phosphate group by phosphodiester bond; more importantly, the structure is of A-form geometry, which is constructed as of vast and thin major groove and of flat and broad minor groove, the structure can fold on itself to form secondary structure, such as tRNA and rRNA, and the secondary structure that are stabilized by hydrogen bonds, domains of loops, and metal ions, such as Mg 2+, form specific tertiary form.

List of Nitrogenous bases found in DNA and RNA

Name	Abbreviation	Structure	Classification	Found in
Cytosine	C		Pyrimidine	DNA, RNA
Thymine	T		Pyrimidine	DNA
Uracil	U		Pyrimidine	RNA
Adenine	A		Purine	DNA, RNA
Guanine	G		Purine	DNA, RNA

References

Berg, Jeremy; Tymoczko, John; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. 2007.

Lipids Overview

Lipids are naturally occurring (organic) compounds that are insoluble in polar solvents such as water . Their insolubility can be attributed solely to their long hydrophobic hydrocarbon chains. These hydrophobic chains may be saturated or unsaturated. Unsaturated chains contain double or triple covalent bonds between adjacent carbons while saturated chains consist of all single bonds. Lipids are composed of a glycerol molecule bonded to long hydrocarbon chain(s) (can be single or multiple) and, depending on the lipid, to other molecules—such as a phosphate group (phospholipids).

Some examples of the types of lipids are: neutral, saturated, (poly/mono) unsaturated fats and oils (monoglycerides, diglycerides, triglycerides), phospholipids, sterols (steroid alcohols), zoosterols (cholesterol), waxes, and fat-soluble vitamins (vitamins A, D, E, and K). Lipids have many different biological functions such as fuel molecules, structural building blocks for phospholipids and glycolipids, covalent attachments to guide molecules to specific membrane locations, and intracellular messengers.

"Chemical structure of the saccharolipid lipid A as found in E. Coli."

There are three common types of Membrane Lipids. They are phospholipids, glycolipids, and cholesterol. [Structural Biochemistry].

Fats Fats consists of glycerol and 3 fatty acids. Fats are created via 3 condensation reactions creating ester linkages that link the fatty acid carboxyl groups to the hydroxyl groups in glycerol. There are two different types of fatty acids, saturated and unsaturated. In a saturated fatty acid, it has the maximum number of hydrogen atoms possible, thus there are no double bonds. There are only single bonds. Since saturated fatty acids are only single bonds, it can pack more tightly together at room temperature and this makes it a solid at room temperature. An example of a saturated fatty acid is butter. An unsaturated fatty acid has one more double bonds. These double bonds create a kink in the hydrocarbon tail, which in return results in looser packing. At room temperature, it is a liquid. An example of this is oil.

Phospholipids They are found in biological membranes. The components of phospholipids include a hydrophobic tail and hydrophilic head. The hydrophobic tail consists of two hydrocarbon chains. The hydrophilic head consists of choline, phosphate, and glycerol. The fatty acids give a hydrophobic barrier, whereas the remainder of the molecule has hydrophilic properties. Phospholipids spontaneously form lipid bilayers due to amphipathic nature of lipid molecules. Phospholipids are found in all cell membranes.

Cholesterol Cholesterol is a steroid and they are built from 4 fused hydrocarbon rings. The hydrocarbon tail is connected to the steroid at one end, and a hydroxyl group is connected to the other end. Cholesterol is a steroid important in cell membranes and acts as a precursor to some sex hormones. However, prokaryotes do not have cholesterol.

Triglycerides

Neutral fats (triglycerides) are composed of fatty acid hydrocarbon chains bonded to a single glycerol molecule. Fatty acids consist of long hydrocarbon chains with a carboxyl group while glycerol consists of 3 carbons and 3 hydroxyl groups. Fatty acids are the building blocks of fat molecules. The method by which the three fatty acid chains in a triglyceride attach to a single glycerol molecule is called dehydration synthesis. Dehydration synthesis is also used in various other reactions, including the joining of two monosaccharides to form a disaccharide. Triglycerides function primarily in energy storage, as a form of insulation, and to protect and cushion cells and organs.

There is an image of a triglyceride molecule with three neutral fatty acid chains and a glycerol group

Saturated fatty acids contain single bonds between the carbons of the hydrophobic chain. Saturated fatty acids originate from animals and are found as component chains in a triglyceride molecule. Saturated fatty acids exist in the solid state at room temperature. Unsaturated fatty acids however contain one (monounsaturated) or more (polyunsaturated) double bond(s) between the carbons of the hydrocarbon chain, which causes the molecule to bend. Triglycerides with too many bends cannot be packed as closely together as neutral fatty acids and therefore are less dense. Below is an example of a saturated fatty acid

A Saturated Fatty Acid

Triglycerides composed of many fatty acids that melt at lower temperatures than those triglycerides with saturated fatty acids. These unsaturated fatty acids do not bind at their maximum number of hydrogen’s because of double bonding between the carbons of the chain. Unsaturated fatty acids originate from plants and are found as component chains to triglyceride molecules. Unsaturated fatty acids exist in the liquid state at room temperature.

(E)-4-oxohexadec-2-enoic acid, An Unsaturated Fatty Acid

These images depict a saturated fatty acid chain (contain single carbon bonds) and an unsaturated fatty acid chain (contain double carbon bonds).

Phospholipids

Phospholipid structure.

Phospholipids are modified triglycerides with one of the fatty acid chains replaced with a phosphate group. They are made by four distinguished groups: fatty acid chains, a platform, a phosphate group, and an alcohol attached to the phosphate. The fatty acid chains are hydrocarbon chains that are typically 14-24 carbons in length. The platform is either glycerol or sphingosine, which is an amino alcohol with a hydrocarbon chain. Phospholipids have a very characteristic non-polar fatty acid chain portion and a polar phosphate portion.The amphipathic character of phospholipids contribute in its' crucial role in phospholipid bilayers. The polar phosphate group is capable of interacting with water molecules and spontaneously forms a bilayer in an aquatic environment. Phospholipids orientate themselves so that the polar heads are facing the water molecules and the hydrophobic fatty acids are oriented toward the inside of the bi-layer. The bi-layer environment enables the non-polar fatty acid chains to stay together, avoiding the water while the hydrophilic phosphate group is oriented toward the water. Phospholipids participate in the formation of the cell membrane by the coming together of two layers of phospholipids. The phospholipids are responsible for the membrane's semi-permeability and fluidity.

The structure describe a phospholipid:

1-Acylglycerin-3-phosphoethanolamin

This image illustrates the components and orientation of the hydrophilic phosphate group and the hydrophobic fatty acid chains that form the lipid bi-layer.

Phospholipid is the most common group of lipids. In fact, cell membranes as well as organized cellular compartments are all made up of these phospholipids. They can form structures called micelles, in which when phosopholipids congregate, the hydrophobic fatty acid tails join together in the center of the sphere away from the aqueous environment and the polar heads are exposed to the outside. Structures such as liposomes can also be artificially formed from these lipids: using high frequency sound waves to sonicate the sample containing phospholipids and molecules of interest to create phospholipid vesicles that contain the molecules of interest. This is often used to deliver drugs to the cells and study how drugs pass through the membrane.

The plasma membrane is made up of the phospholipid bilayer. The membrane is an amphipathic sheet-like structure that is fluid and electrically polarized. The membrane itself has little functions, but the proteins that are integrally and peripherally integrated to the membrane help mediate many of the functions that we contribute to membrane. The membrane is asymmetric in that the proteins are randomly distributed across the membrane, some are attached inside the cell, some outside, and others integrated within membrane. Also, rapid lateral diffusion and slow transverse diffusion contribute both to the membrane's asymmetric characteristic and fluidity. In transverse diffusion, phospholipids are flipped inside-out or outside-in, and this flipping is regulated by flipases. However, the longer the fatty acid chains are, the less likely for transverse diffusion to occur. Longer chains also decreases the fluidity of the plasma membrane. There are other factors that may affect plasma membrane's fluidity. For example, the better arranged the fatty acids chains are, the less fluid the membrane is. On top of that, the more unsaturated the fatty acids are, the more fluid the membrane is. This is because the double bonds bend the chains that allows sloppy arrangements. The interruption of cholesterol within the membrane also causes more fluidity since the polar hydroxy group in cholesterol disrupts the hydrophobic environment within the phospholipid bilayer.

Glycolipids

Glycolipids are sugar(glyco-)containing lipids. They are derived from sphingosine instead of a form of phospholipids that derives from glycerol (phospholipids exist in both derivatives from glycerol and sphingomyelin platform). Another difference from phospholipids is that glycolipids contain a sugar unit (can be glucose or galactose) instead of a phosphate group.

Examples: Glycolipid molecules exist from the most basic molecule, cerebroside which contains 1 fatty acid unit, a sphingosine backbone, and 1 sugar unit (glucose or galactose), to the most complex molecules containing branched chains of multiple sugar residues (up to seven residues in gangliosides).

Properties: When glycolipids exist in membranes, their sugar residue terminal always face the extracellular side.

Chemical structure of Glycolipids

Structure of Glycolipids'

Cholesterol

Cholesterol

Cholesterol is a form of lipids that differs from the rest of its relatives. It is relatively medium molecule that contains 4 adjacent cyclic hydrocarbon molecules with three six-member rings and one five-member ring that has a hydroxyl and a saturated hydrocarbon chain terminals.

The molecule functions as a bufferor a temperature stabilizer for the membrane in which it can make up of 25% of the membrane. When exist in membranes, the 4 cyclic molecules in the cholesterol molecule lay parallel to the fatty acid chains of the phospholipids, meanwhile the hydroxyl terminal points in the direction with the polar phospholipid heads in which it interact with.

Cholesterol molecules exist primarily in nerve cells. The molecule binds to the myelin sheath membrane which provides an outer coating that protects the nerve cell from its surroundings.

It is an essential predecessor to sex hormones that exists in males (testosterone) and females (oestradiol). Also an essential component in vitamin D that enables the body to utilize calcium to form bones.

Animals acquire very little cholesterol from the food they eat; they make cholesterol within the body. Although cholesterol is essential for many processes and structural function, it can be detrimental to have excess cholesterol. Too much cholesterol in the blood will cause blockages in the arteries which can result in heart disease, high blood pressure, and stroke. Only 0.25% of human beings retain High Cholesterol disease from heredity, however people are gaining high cholesterol in their blood from the food they eat (especially people in America).

Membrane Properties

The cell membrane has a set of properties that are contributed to the presences of the lipids as well as proteins. 1. Structures are lik sheets. 2. They are formed by lipids, proteins, and carbohydrates. 3. The membrane is amphipathic (contain hydrophobic and hydrophilic regions). 4. Each protein allows for the function of the membrane. 5. Membranes are held on by weak, non-covalent bonds. 6. The structure is asymmetric. 7. It is high in fluidity. 8. The membrane is polarized.

Fluidity

Lipid unsaturation effect

The presence of the lipids in the membrane structure of a cell is vital for the cell especially affecting its fluidity. As addressed in the section below, this is necessary to allow things to flow in and out of the cell. One factor that plays a big role in this is cholesterol as shown below. Another one is the presence of double bonds. The more the double bonds, the greater the amount of kinks or curve in the lipid and therefore more free space. The length of the lipids also plays a role. Lipids move in two distinct ways. Most commonly they interact in lateral diffusion where they switch places with the lipid to the left or right of them. Other times, they go through transverse diffusion where they flip flop with the ones whose tails they are facing. This is due to the weak, Van der Waals interaction of the lipid molecules. The longer the lipid is, the stronger this interaction is, therefore decreasing the mobility of the lipids. Decrease lipid mobility yields in decreasing the fluidity of the membrane.

Cholesterol and fluidity

This shows the cholesterol molecules submerged in the lipid bilayer

Cholesterol is an important factor in membrane permeability, that is, how much can flow through the cell. Cholesterol acts as a 'buffer' to prevent against any extremes. Obviously, the membranes permeability cannot be too fluid as to allow anything inside the cell(i.e. harmful agents),but at the same time, the membranes permeability must be fluid enough as to let out/let in important agents that need to enter the cell.

Cholesterol is a hydrocarbon steroid with one single alcohol group, which leads to its amphiphatic nature.

Generally,

1.AT LOW TEMPERATURES:Cholesterol in a membrane leads to a more fluid membrane.

2.AT HIGH TEMPERATURES:Cholesterol in a membrane leads to a less fluid membrane.

Ether Lipids with Branched Chains

An ether lipid.

Two major factor that separates Archea(bacteria) from Bacteria is the Archea's cell membranes phospholipid consists of ether linkages and the fatty acid hydrocarbon chains are completely saturated and branched with a methyl group every 5 carbons. These simple structural differences provides Archeabacteria drastic difference from bacteria in terms of their habitat's harsh environments. These 2 factors contribute to the chemical properties that their membrane is more resistant to hydrolysis (ether versus ester linkages) and resistant to oxidation (branched saturated hydrocarbon chains).

Phospholipids and Glycolipids Readily Form Bimolecular Sheets in Aqueous Media

Phospholipids and glycolipids have amphipathic characteristics which enables them to form a micelle or a "lipid bilayer". Due to the hydrophobic hydrocarbon tail and the hydrophillic polar head group, the lipids arrange in a form where the polar groups face water while the tail is away from water. One formation is the micelle where the lipids arrange themselves in a circle with the head groups making the circumference while the tails are inside. A more favorable formation is the lipid bilayer or the bimolecular sheet. This arrangement has the lipids form a barrier where the polar head groups face the aqueous media and the hydrophobic tail face inside away from water. This type of formation is favorable to cell membranes for it forms a barrier from the extracellular fluid and protects the cytoplasm within the cell. Integral and peripheral proteins may be present in the lipid bilayer to allow certain functions to occur such as transportation of ions or acting as pumps.

References

Berg, Jeremy; Tymoczko, John; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. 2007. Berg, Jeremy M., Tymoczko, John, L., Stryer, Lubert. Biochemistry. Seventh Edition.

The Organics of Biochemistry:[12]

"Lipid bilayer." Wikipedia. 7 December 2012. 7 December 2012 <http://en.wikipedia.org/wiki/Lipid_bilayer>.

Viadiu, Hector. "Lipids and Cell Membranes." UCSD, 19 November 2012. Protein molecules contain polypeptide chains made from sequences of the 20 amino acids. These amino acids are linked together by a peptide bond that is formed by condensation of two amino acids with the elimination of the elements of water. Protein function is dependent on its tertiary structure. Proteins tend to fold into three- dimensional structures because of the sequence of amino acids. Proteins also contain functional groups from each amino acid. These groups are reactive and also contribute to protein function. Proteins also interact with one another and with other macromolecules. Proteins can be rigid or flexible. This allows certain proteins to be found in different parts of the cell such as the cytoskeleton or in soft tissue.

Importance of Proteins

Enzymes are proteins that catalyze chemical reactions. Enzymes speed up the reactions in biological systems by lowering the activation barrier needed to start that reaction.

Hormones are proteins that are chemical messengers in the body. These proteins are sent to different parts of the body to send or receive messages. Hormones are very important in regulating the human body and keeping the body in a state of homeostasis. Some protein hormones include insulin, growth hormone, Luteinizing hormone (LH), follicle-stimulating hormone (FSH), and thyroid-stimulating hormone (TSH). These proteins are part of the glycoprotein hormones.

Transport Proteins are also used to transportation. For example, hemoglobin is a metalloprotein(protein that contain a metal as cofactor) that transports an oxygen in the red blood cells with the help of iron.

Motor Proteins help convert chemical energy to mechanical energy which relates to muscular motion in an organism. Examples are actin and myosin.

Protective Proteins protect cells by releasing, making antibodies, fighting and destroying foreign objects. Antibodies are gamma globulin proteins.

Structural Proteins help maintain the structure of a variety of biological components like cells and tissues in an organism. Collagen, elastin, α-keratin, sklerotin, and fibroin are all examples of proteins that contribute to the formation of an organism's body.

Storage Proteins that contain energy and can be digested during metabolism of the organism. Examples are egg ovalbumin and milk casein.

Membrane Proteins include receptors and membrane transport. The receptors in the membrane of cells allow ions to pass through. These prevent unwanted objects from coming into the cell. These receptors also determine if the cell is excited to create an action potential or not. Membrane transport is important because it allows ions, proteins, and other macromolecules to pass through the cell membrane.

Classification of proteins by location

External proteins-proteins outside of cells and are found in multicell organisms.

Internal proteins-proteins that are inside cells and perform functions for intercellular needs.

Membrane proteins-proteins that are embedded in the bilayer of the membrane of on the edges of the membrane helping with intracellular interactions.

Virus proteins-usually the coat for viruses

Enzymes

Enzymes are proteins that speed up the reaction rate. Many reactions cannot occur without the use of an enzyme.

This graph shows how the presence of an enzyme lowers the activation energy and therefore speeds up the reaction. Reactions with enzymes can speed up to 10 billion times faster than those without an enzyme. The rate at which the enzyme works is affected by the substrate and enzyme concentration, temperature, and pH.

Amino Acids

There are 20 amino acids that make up proteins. The main chain has an N terminal which is an amino group (NH₂) and a C terminal which is a carboxyl group (COOH). The side chains make each amino acid unique.

The 20 different amino acids can be classified into six different classes based on their side chain(R group).

1. Aliphatic - carbon side chain. The longer the aliphatic chain, the more hydrophobic.

Glycine, Alanine, Valine, Leucine, Isoleucine

2. Hydroxyl or Sulfur containing- The OH is reactive, hydrophilic (water loving), polar and uncharged. Sulfur is very reactive.

Serine, Threonine, Cysteine, Methionine

3. Cyclic- Proline

4. Aromatic-

Phenylalanine - purely hydrophobic

Tyrosine - OH reactive

Tryptophan - less hydrophobic due to its NH groups

5. Basic- hydrophilic and positively charged.

Lysine, Arginine, Histidine

6. Acidic and their amide- negatively charged.

Aspartate, Glutamate, Asparagine, Glutamine

This is the general structure of an amino acid in its unionized form.

This table shows all the amino acids and the their side chains. From left to right they are: glycine(Gly), alanine(Ala), valine(Val), Leucine(Leu), isoleucine(ILe), methionine(Met), phenylalanine(Phe), proline(Pro), aspartic acid(Asp), glutamic acid(Glu),serine(Ser), threonine(Thr), cystine(Cys), tyrosine(Tyr), asparagine(Asn), glutamine(Gln), tryptophan(Trp), lysine(Lys), arginine(Arg), and histidine (His).

Peptide Bonds

Proteins are made from many amino acids. They are connected together by peptide bond. Peptide bonds are formed by condensation, the loss of awater molecule and between the carboxyl group with the aminogroup.

The reaction above shows how two alanine are linked together by a peptide bond. The bond is formed between the n terminal amino group and the c terminal carboxyl group. Two hydrogens and an oxygen come out in this reaction which produces water. The peptide bond acts almost like a double bond due to the resonance of the carbonyl. Because of this, there is no rotation about this bond so therefore conformation is limited. This limits stereochemistry in that almost all peptide bonds in proteins are trans isomers to limit steric hindrance between the R groups. Only proline can be either cis or trans because the energy levels these two isomers show are about the same(the side chain of proline has similar distance to the adjacent R group in either isomer). The reason for that is because proline's side chain form a ring with the alpha-amino group. Proline is the only amino acid whose side chain form a ring with the alpha-amino group.

Structures

Proteins can fold into four different structures. These structures determine protein function and characteristics.

Primary Structure- The primary structure of a polypeptide is its amino acid sequence, from beginning to end. The primary structures of polypeptides are determined by genes. Genes carry the information to make polypeptides with a defined amino acid sequence. For the protein to function correctly, each amino acid needs to be in order as the genes assigned. Even a little change in the amino acid sequence would affect the shape of the protein and its ability to function. An average polypeptide is about 300 amino acids in length, and some genes encode polypeptides that are a few thousand amino acids long.

Secondary Structures- The amino acid sequence of a polypeptide, together with the laws of chemistry and physics, cause a polypeptide to fold into a more compact structure. Amino acids can rotate around bonds within a protein. This is the reason proteins are flexible and can fold into a member of shapes. Folding can be irregular or certain regions can gave a repeating folding pattern. Such repeating patterns are called secondary structures. The two types are the α-helix and β-sheet. In an α-helix, the polypeptide backbone forms a repeating helical structure that is stabilized by hydrogen bonds. These hydrogen bonds occur at regular intervals and cause the polypeptide backbone to form a helix. In a β-sheet, regions of the polypeptide backbone come to lie parallel to each other. When these regions form hydrogen bonds, the polypeptide backbone form a repeating zigzag shape called a β-sheet.

Tertiary Structure- As the secondary structure becomes established due to the primary structure, a polypeptide folds and refolds upon itself to assume a complex three-dimensional shape called the protein tertiary structure. The tertiary structure is the three-dimensional shape of a single polypeptide. It's usually a result of interactions among the R groups of the amino acids that make up the polypeptide. For some proteins, such as ribonuclease, the tertiary structure is the final structure of a functional protein. Other proteins are composed of two or more polypeptides and adopt a quaternary structure. Tertiary structure is important in regards to enzymatic activity.

Quaternary Structure- Most functional proteins are composed of two or polypeptide that each adopt a tertiary structure and then assemble with each other. The individual polypeptides are called protein subunits. Subunits may be identical polypeptides or they may be different. Each subunit has a non-protein component that is necessary for the protein to function correctly. These components are called heme. When proteins consist of more than one polypeptide chain, they are said to have quaternary structure and are also known as multimeric proteins, meaning many parts.

Factors that influence protein structure

Several factors determine the way that polypeptides adopt their secondary, tertiary and quaternary structures. The amino acid sequences of polypeptides are the defining features that distinguish the structure of one protein from another. As polypeptides are synthesized in a cell, they fold into secondary and tertiary structures, which assemble into quaternary structures for most proteins. As mentioned, the laws of Chemistry and physics, together with amino acid sequence, govern this process. Five factors are critical for protein folding and stability:

1. Hydrogen bonds

2.