The reaction rate or rate of reaction is the speed at which a chemical reaction takes place, defined as proportional to the increase in the concentration of a product per unit time or to the decrease in the concentration of a reactant per unit time. Reaction rates can vary dramatically. For example, the oxidative rusting of iron under Earth's atmosphere is a slow reaction that can take many years, but the combustion of cellulose in fire is a reaction that takes place in fractions of a second. For most reactions, the rate decreases as the reaction proceeds. A reaction's rate can be determined by measuring the changes in concentration over time. Chemical kinetics is the part of physical chemistry that concerns how rates of chemical reactions are measured and predicted, and how reaction-rate data can be used to deduce probable reaction mechanisms. The concepts of chemical kinetics are applied in many disciplines, such as chemical engineering, enzymology and environmental engineering.
Formal definition Consider a typical balanced chemical reaction:
a A
+ b B ⟶ p P
+ q Q {\displaystyle {\ce {{\mathit {a}}\,A{}+{\mathit {b}}\,B->{\mathit {p}}\,P{}+{\mathit {q}}\,Q}}}
The lowercase letters (a, b, p, and q) represent stoichiometric coefficients, while the capital letters represent the reactants (A and B) and the products (P and Q). According to IUPAC's Gold Book definition the reaction rate ν {\displaystyle \nu } for a chemical reaction occurring in a closed system at constant volume, without a build-up of reaction intermediates, is defined as
ν = − 1 a d [ A ] d t = − 1 b d [ B ] d t = 1 p d [ P ] d t = 1 q d [ Q ] d t , {\displaystyle \nu =-{\frac {1}{a}}{\frac {d[\mathrm {A} ]}{dt}}=-{\frac {1}{b}}{\frac {d[\mathrm {B} ]}{dt}}={\frac {1}{p}}{\frac {d[\mathrm {P} ]}{dt}}={\frac {1}{q}}{\frac {d[\mathrm {Q} ]}{dt}},}
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