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Glossary: Reaction rates

activated complex. transition state.
An intermediate structure formed in the conversion of reactants to products. The activated complex is the structure at the maximum energy point along the reaction path; the activation energy* is the difference between the energies of the activated complex and the reactants.

activation energy. (Ea)
The minimum energy required to convert reactants into products; the difference between the energies of the activated complex* and the reactants.

Arrhenius equation.
In 1889, Svante Arrhenius explained the variation of rate constants* with temperature for several elementary reactions* using the relationship

k = A exp(-Ea/RT)

where the rate constant k is the total frequency of collisions* between reaction molecules A times the fraction of collisions exp(-Ea/RT) that have an energy that exceeds a threshold activation energy* Ea at a temperature of T (in kelvins). R is the universal gas constant*.

catalyst. catalyze; catalysis.
A substance that increases the rate of a chemical reaction, without being consumed or produced by the reaction. Catalysts speed both the forward and reverse reactions, without changing the position of equilibrium*. Enzymes* are catalysts for many biochemical reactions.

collision frequency. collision frequencies; frequency of collision.
The average number of collisions that a molecule undergoes each second.

collision theory. collision model.
A theory that explains reaction rates* in terms of collisions between reactant* molecules.

elementary reaction. Compare with net chemical reaction*.
A reaction that occurs in a single step. Equations for elementary reactions show the actual molecules, atoms, and ions that react on a molecular level.

Protein* or protein-based molecules that speed up chemical reactions occurring in living things. Enzymes act as catalysts* for a single reaction, converting a specific set of reactants (called substrates*) into specific products. Without enzymes life as we know it would be impossible.

first order reaction. Compare with zero order reaction* and second order reaction*.
The sum of concentration exponents in the rate law for a first order reaction is one. Many radioactive decays are first order reactions.

half life.
The half life of a reaction is the time required for the amount of reactant to drop to one half its initial value.

integrated rate law.
Rate laws like d[A]/dt = -k[A] give instantaneous concentration changes. To find the change in concentration over time, the instantaneous changes must by added (integrated) over the desired time interval. The rate law d[A]/dt = -k[A] can be integrated from time zero to time t to obtain the integrated rate law ln([A]/[A]o = -kt, where [A]o is the initial concentration of A.

intermediate. reactive intermediate; reaction intermediate.
A highly reactive substance that forms and then reacts further during the conversion of reactants* to products* in a chemical reaction. Intermediates never appear as products in the chemical equation* for a net chemical reaction*.

order. order of reaction; reaction order.
The order of a reaction is the sum of concentration exponents in the rate law for the reaction. For example, a reaction with rate law d[C]/dt = k[A]2[B] would be a third order reaction. Noninteger orders are possible.

rate constant. (k)
A rate constant is a proportionality constant that appears in a rate law*. For example, k is the rate constant in the rate law d[A]/dt = k[A]. Rate constants are independent of concentration but depend on other factors, most notably temperature.

rate law.
A rate law or rate equation relates reaction rate* with the concentrations of reactants, catalysts, and inhibitors. For example, the rate law for the one-step reaction A + B rightarrow C is d[C]/dt = k[A][B].

reaction mechanism. mechanism.
A list of all elementary reactions* that occur in the course of an overall chemical reaction*.

reaction rate.
A reaction rate is the speed at which reactants are converted into products in a chemical reaction. The reaction rate is given as the instantaneous rate of change for any reactant or product, and is usually written as a derivative (e. g. d[A]/dt) with units of concentration per unit time (e. g. mol L-1 s-1).

second order reaction. Compare with zero order reaction* and first order reaction*.
A reaction with a rate law that is proportional to either the concentration of a reactant squared, or the product of concentrations of two reactants.

unimolecular reaction.
A reaction that involves isomerization* or decomposition* of a single molecule.

zero order reaction. Compare with first order reaction* and second order reaction*.
A reaction with a reaction rate* that does not change when reactant concentrations change.

General Chemistry Online! Reaction rates

Copyright © 1997-2010 by Fred Senese
Comments & questions to fsenese@frostburg.edu
Last Revised 08/17/15.URL: http://antoine.frostburg.edu/chem/senese/101/kinetics/glossary.shtml