Kinetics is a study of the progress of a chemical reaction, especially the rate of the reaction.
Rates of ReactionEdit
The rate of reaction is concerned with how quickly a reaction reaches a certain point. Technically, it is defined as the increase of product concentration over time, or, the decrease of reactant concentration over time. The SI unit of rate of reaction is mol/dm3, which means moles per dm3 or moles per liter. In short, the rate of a chemical reaction can be generalized as the following, when X is a reactant:
The rate of reaction can only be determined experimentally.
7.1.2 : Interpretation of rate graphs. Reaction rate graphs will generally be graphed with time on the x-axis and some measure of how far the reaction has gone (ie concentration, volume, mass loss etc) on the y-axis. This will generally produce a curve with, for example, the concentration of the reactants approaching zero.
The Collision TheoryEdit
The collision theory is the main theory that concerns kinetics. It tells us under which circumstances reactions occur and how different factors affect the rate of reaction. According to the collision theory, reactions can only occur when:
- Reactants collide in the correct orientation.
- Reactants have the enough energy to form its products.
The energy required for reactants to form its products is called the activation energy (denoted as Ea). The activation energy is essential in kinetics, as it allows us to determine whether if a reaction can take place. A reaction is said to be spontaneous when it able to produce reactants.
What Happens During a ReactionEdit
A chemical reaction is essentially involves two processes, bond breaking and bond forming. As mentioned in the Energetics chapter, particles tend to react to become more stable.
The Reaction CoordinateEdit
The reaction coordinate, also called the reaction pathway, describes the changes in energy in the system throughout the reaction. The reaction coordinate is usually presented graphically, in a diagram called the energy profile (see Fig. 1).
The activation energy is the energy change from the reactants to the highest energy state of the reaction.
ΔH, or enthalpy change, is the difference between the energy of the system when reactants are formed and the energy of the system when products are formed. Enthalpy is very similar to energy, but enthalpy is the energy of the system under a constant pressure. The enthalpy change is particularly important, as it determines whether if the reaction is exothermic, or endothermic. As you might have noticed on the graph, the energy of the A+B are lower than that of C+D. Where did the energy go? The energy is released outside the system, in other words, the energy is released to the environment. Reactions that result in energy released to the environment are called exothermic reactions. When the energy of reactants are lower than that of products, then energy is absorbed into the system, making the reaction exothermic.
Usually, the energy released to the environment in the form of heat energy, enabling us to detect a rise in temperature. However, the released energy can also be released in the form of light and sound energy. Similarly, endothermic reactions, which absorbs energy from the environment, cause an increase in energy in the system, usually resulting in a drop in temperature. This description leads to one of the most common misconceptions of endothermic reactions, as you may wonder why temperature decreases when energy is increased in the system. Energy is stored in bonds of the particles in the system in the form of chemical energy, rather than heat energy. Therefore, when a decrease in temperature is detected (using an instrument or through touch), heat energy is decreased in the environment.
7.2.1 : Collision theory -- reactions take place as a result of particles (atoms or molecules) colliding and then undergoing a reaction. Not all collisions cause reaction, however, even in a system where the reaction is spontaneous. The particles must have sufficient kinetic energy, and the correct orientation with respect to each other for the two to react. Even then, the transition state may revert to the reactant molecules instead of forming the product molecules.
7.2.2 : Higher temperature causes a greater average kinetic energy of the particles in a material. This leads to a faster reaction because there are more collisions, and each collision is more likely to succeed.
Higher concentrations cause more collisions and therefore a faster reaction.
Catalysts may provide an alternative pathway with lower activation energy and increase the probability of proper orientation. Each collision is more likely to succeed and this results in a faster reaction.
In heterogeneous reactions (where the reactants are in different states) the size of the particles of a solid may change reaction rate, since the surface is where the reaction takes place, and the surface area is increased when the particles are more finely divided; therefore smaller solid particles in a heterogeneous reaction cause a faster reaction.
7.2.3 : Most reactions involve several steps, which can be individually slow or fast, and which, all together, make up the complete reaction. The slowest of these steps is called the rate determining step, as is determines how fast the reaction will go. It is also not necessary that all the reactants are involved in ever step, and so the rate determining step may not involve all the reactants. as a result, increasing their concentration (for example) of a reactant which is not involved in the rate determining step will not change the overall reaction rate.
Topic 16 is the additional HL material for Topic 7.