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Glossary
- 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.
- enzyme.
- 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 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.
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