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
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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.