Chemical kinetics (kinetics): The study of the rates of chemical reactions and the steps by which they occur.
Rate: Rapidity of change, a property involving time.
Some chemical reactions are very fast and appear instantaneous to the eye on mixing the reactants, while others are very slow and the products only form over minutes, hours or even longer.Reaction rate: The change in concentration of a reactant or product divided by the time it takes for the change to occur. Common units are mol L−1 s−1.
In this example the numerical value of the rate depends on which reactant or product concentration is being measured. To overcome this it is possible to express the rate of reaction defined by its chemical equation.
Most chemical reactions take place in a series of elementary reactions (steps).
Unimolecular reaction: An elementary reaction involving only one particle Termolecular reaction: An elementary reaction involving only three particles.
Termolecular reaction: An elementary reaction involving only three particles.
X + C P + Q
The overall reaction is A + B + C P + Q
The species X is called a reaction intermediate.
Reaction intermediate: A species produced in one step of a chemical reaction but which is consumed in a later step. It is not a product of the overall reaction, and hence does not appear in the overall equation.
Rate equation (law): A mathematical equation relating the instantaneous rate of the reaction to the concentrations of reactants, products, or any other permanent species (such as catalysts) taking part in the reaction at that instance
For a reaction between two species to occur they must collide, and most reactions take place in a series of bimolecular collisions, i.e. by a series of bimolecular elementary steps. If the step involves two different species, A and B, the rate of collision, and therefore the rate of reaction, is proportional to the concentrations of A and B. Thus for the elementary bimolecular reaction
As bonds are often broken in elementary reactions not all collisions between reactant species lead to reaction, only those which collide with sufficient energy to overcome the activation energy barrier.
Activation energy: Symbol Ea. The minimum energy needed for colliding species to react, or cross the barrier. The magnitude of k depends on the height of the barrier, i.e. on the activation energy. The greater the Ea, the smaller the k, and the slower the reaction. The activation energy, Ea, is determined from an Arrhenius plot.
Arrhenius plot: A plot of ln k or ln rate against 1/T where T is the absolute temperature (temperature in kelvin). The slope of an Arrhenius plot equals -Ea/R where R is the universal gas constant. This follows from the Arrhenius equation.
Arrhenius equation: k = Aexp(−Ea/RT) where A is called the Arrhenius pre-exponential factor.
The kinetic energy of molecules, and hence the fraction of collisions with energy in excess of Ea increases with increasing temperature. Therefore the rate of a reaction, and the magnitude of k, increase with temperature. The greater the Ea, the more sensitive the rate and rate constant to changes in temperature.
Unlike equilibrium constant expressions, the rate law for a reaction cannot be derived from inspection of the overall chemical equation but must be determined experimentally. Only for elementary reactions (i.e. one step reactions) can the rate law be deduced directly from the equation (as explained above).
Information on reaction mechanisms can be obtained from experimentally obtained rate laws. While a rate law for a reaction can be derived for a postulated mechanism (because the rate law for each elementary step can be written), agreement between the experimental rate law and postulated mechanism does not prove the mechanism. Different mechanisms can lead to the same rate law.
Integrated rate equation: A mathematical equation which relates the concentration of a species at a given time to its initial concentration and the time, and which involves the rate constant
Half-life: Symbol t½, the time for the amount of a reactant in limiting amount (i.e. not in excess) to fall to half its original value. Commonly applied to radioactive isotopes
Mean life-time: Symbol τ, the time for the amount of a reactant in limiting amount (i.e. not in excess) to fall to 0.368 (1/e) of its original value. This is the average time a molecule must wait before undergoing reaction.
Chain reaction: A reaction in which an intermediate formed in an initiation step produces further reactive intermediates in the propagation steps until removed in a termination step.
Catalyst: A substance which increases the rate of a reaction without being consumed in the overall reaction. It provides an alternative pathway or mechanism which increases the rate of the overall reaction. The catalyst is involved in the new pathway but if it is consumed in one step it is regenerated in a later step.
Homogeneous catalysis: The catalyst is in the same phase as the reactants.
Heterogeneous catalysis: The catalyst is in a different phase from the reactants. A solid catalyst adsorbs reacting species onto its surface where the reaction takes place.