The rate of a reaction has been defined in a basic level article as:

For a simple reaction, *A* → *B* at constant volume *V _{1}*, we can define its rate as

where *n _{B}* is the amount of

*B*.

Since the rate of formation of product must be the same as the rate of consumption of the reactant, and that the rate of a reaction is defined as a positive value, we have: .

Suppose we have another reaction, *C* → *D *at a different constant volume *V _{2}*, with a rate:

To have a meaningful comparison of the rates of these two reactions with different reacting volumes, we need a common basis, which can be obtained by dividing eq2 and eq3 by their respective volumes. Therefore, a better definition of the rate of a reaction is:

Consider an experiment where *HCl* is added to *CaCO _{3}* to produce

*CO*in a rigid vessel with a pressure gauge. The data collected are presented in a

_{2 }*CO*pressure versus time plot. If we use a constant mass of

_{2}*CaCO*and 3 different concentrations of 20 cm

_{3}^{3}of

*HCl*, we have 3 different gradients at the origin, which represent 3 initial rates of reaction (see diagram below) with

*R*corresponding to the reaction with the highest concentration of

_{A}*HCl*and

*R*relating to the reaction with the lowest concentration of

_{C}*HCl*.

The above results imply that the rate of reaction is proportional to the concentration of the reactant used. Therefore, in addition to eq4, we can also expressed the rate of a reaction in the form of an equation known as a ** rate law**:

where [*H _{3}O^{+}*] is the concentration of hydroxonium ions from

*HCl*and

*k*is a proportionality constant called the

**. Combining eq4 and eq5,**

*rate constant*Eq6 means that an increase in concentration of hydroxonium ions increases the rate of change in concentration of *CO _{2}*, which is measured by the rate of change in pressure of

*CO*, since .

_{2}