Structural Biochemistry/Chemical Bonding/Covalent bonds

Introduction edit

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 edit

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. [1] 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. [2]

Types of Covalent Bonds edit

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 edit

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 edit

Formation of a Pi-Bond from two p-orbitals.

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 edit


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 edit

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.[3]

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 edit

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 edit

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 ΔHrxn. 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. ΔH0rxn= ΔHreactant bonds broken + ΔHproduct bonds formed[4]

References edit

  1. Organic Chemistry by Vollhardt and Shore
  3. Berg, Jeremy; Tymoczko, John; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. 2007. (7)
  4. Silberberg, Martin S.(2010). Principles of General Chemistry (2nd Edition).McGraw Hill Publishing Company. ISBN978-0-07-351108-05

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