Rate Law
Rate = k[A]^m[B]^n
The rate law expresses reaction rate as a function of reactant concentrations. The orders (m, n) must be determined experimentally and are not necessarily equal to stoichiometric coefficients.
Variables
Rate of reaction in M/s
Constant specific to reaction and temperature
Molar concentrations of reactants
Exponents determined experimentally
Example Calculation
Scenario
For 2NO + O2 -> 2NO2, experiments show Rate = k[NO]^2[O2]. If [NO] = 0.010 M, [O2] = 0.020 M, and k = 7.1 x 10^3 M^-2s^-1, find the rate.
Given Data
Calculation
Rate = k[NO]^2[O2] = (7.1 x 10^3)(0.010)^2(0.020)
Result
Rate = 1.4 x 10^-2 M/s
Interpretation
The reaction proceeds at 0.014 M/s under these conditions. Doubling [NO] would quadruple the rate; doubling [O2] would double it.
When to Use This Formula
- ✓Calculating reaction rates from concentrations
- ✓Determining reaction orders from data
- ✓Predicting how concentration changes affect rate
- ✓Finding rate constants experimentally
Common Mistakes
- ✗Assuming orders equal stoichiometric coefficients
- ✗Using wrong units for rate constant k
- ✗Not recognizing zero-order behavior
- ✗Confusing rate law with integrated rate law
FAQs
Common questions about this formula
Use the method of initial rates. Compare experiments where one concentration changes while others stay constant. If doubling [A] doubles rate, order is 1; if it quadruples rate, order is 2; if no change, order is 0.
Units of k depend on overall reaction order. For first order: s^-1. For second order: M^-1s^-1. For third order: M^-2s^-1. Units must make the rate come out in M/s.