Kinetics (as used by chemists) refers to the study of rates of chemical reaction. This is very important as in organic chemistry you can often get several different products from the same mixture of reagents, but by controlling the reaction conditions the rate of formation of each product may be changed. Knowledge of the kinetics of a reaction helps to produce larger quantities of desired products while reducing the presence of unwanted byproducts.

Getting back to the point, rates of reaction are generally determined by two types of variable: temperature, and the concentration of each reagent.

You can see how this dependence works for a simple reaction if you think about it just a little. A group of molecules can only react if they are all together and if they have energy to form an activated state. Dealing with bulk quantities of materials, you need the probabilities that each type of molecule is in the same place and the probability that enough energy is available. The amount of energy is related to the temperature by a function familiar to anyone who has ever used thermodynamics or statistical mechanics. The probablility of a molecule being in a given place can be measured by the concentration of the molecule.

Needless to say, if you need N molecules of substance A in the same place, the probability of that happening is proportional to the concentration of A raised to the Nth power. The reaction is said to be Nth order with respect to A.

This also gives a 1st order differential equation which allows you to predict the rate of reaction for a simple reaction, but often reactions are more complex. Often you have intermediates, which act as products in one step of the reaction and reagents in the next step. This gives you a set of differential equations which must be solved simultaneously. This set of equations can always be reduced to a single differential equation of higher order, but may be nonlinear if the same reagent is involved in multiple steps of the reaction.

To make things more complicated, you usually don't really know the concentrations of the intermediates, and often don't have any way of measuring them. This is where the steady state approximation is useful.

Basically, for the steady state approximation you pretend the concentrations of the intermediates is constant over the entire duration of the reaction, even if you don't know what it is. Looking at the set of first order equations, you take each equation where a given intermediate is produced or consumed to get an equation describing the concentration of that particular intermediate in terms of concentrations of reagents, products, and other intermediates. Then you go back and substitute that equation for the intermediate concentration in the original equations, and repeat the process with another intermediate, until all you have left are concentrations of product and reagents.

The rate of a chemical reaction, r, quoted in units of mol dm-3 s-1, is given by the expression

r = k [A]0 [B]1 [C]2

where [A], [B] and [C] are the concentrations of the reactants in mol dm-3 and k is the rate constant.

For example, for the reaction between propanone (acetone) and iodine in the presence of hydrogen ions, the equation is

r = k [CH3COCH3]1 [I2]0 [H+] 1

If the concentration of a reactant is raised to the power 0 (making it equal to 1), it has no effect on the rate - it is zero order. It does not really have to be included in the equation at all. If it raised to the power 1 (i.e. not changed at all), it affects the rate and is first order. If it raised to the power 2 (i.e. squared), it has an even stronger effect on the rate and is second order.

Zero order reactants decrease in concentration at a constant rate throughout a reaction. First order reactants decrease at a rate that continually slows, with a constant half-life. Second order reactants decrease at a rate that continually slows and has an increasing half-life.

The order of a reactant is related to its role in the actual chemical process. Reactions take place in steps, forming intermediate compounds. The slowest step determines the overall rate of the reaction and is indeed known as the rate determining step. Only those reactants which are involved in or before this step will appear in the rate equation.

Reactant-orders cannot be predicted theoretically; we must determine them by experiment.

The rate constant, k, depends on temperature and on the orientation of the molecules involved and, for many reactions, is approximately given by the Arrhenius equation:

k = A e-EA/RT

where A is a pre-exponential factor relating to molecular orientation, EA is the activation energy of the reaction, R is the universal gas constant and T is the temperature. e, of course, is just e. As the temperature increases, EA/RT decreases, so e is being raised to a less negative number. Thus k increases, so the reaction goes faster.

With thanks to the editors for their patience.

  • The study of reaction rates.
  • The study of the relationship between force and mass.

  • From the BioTech Dictionary at For further information see the BioTech homenode.

    Ki*net"ics (?), n. Physics

    See Dynamics.


    © Webster 1913.
