Rate laws describe how the speed of a chemical reaction depends on reactant concentrations. They matter because chemists use them to predict how fast products form, compare reaction conditions, and design safer industrial processes. A rate law is found experimentally, not by simply reading the balanced chemical equation.
The central idea is that changing concentration can change collision frequency and sometimes the reaction pathway itself.
For a reaction involving reactants A and B, a common rate law is Rate = k[A]^m[B]^n, where m and n are reaction orders found from data. The overall order is m + n, and it tells how strongly the rate responds when all relevant concentrations change. Initial-rate experiments compare trials where one concentration changes while others stay constant.
Integrated rate laws and concentration versus time graphs help identify zero, first, and second order behavior.
Key Facts
- General rate law: Rate = k[A]^m[B]^n
- Overall reaction order = m + n
- If [A] doubles and rate is unchanged, the reaction is zero order in A: m = 0
- If [A] doubles and rate doubles, the reaction is first order in A: m = 1
- If [A] doubles and rate quadruples, the reaction is second order in A: m = 2
- Units of k depend on overall order: zero order M/s, first order 1/s, second order 1/(M s)
Vocabulary
- Rate law
- An equation that relates reaction rate to reactant concentrations and a rate constant.
- Reaction order
- The exponent on a concentration term in a rate law, showing how that reactant affects the rate.
- Rate constant
- The proportionality constant k in a rate law, whose value depends on temperature and the reaction mechanism.
- Initial rate
- The reaction rate measured near the start of a reaction before concentrations have changed much.
- Integrated rate law
- An equation that relates reactant concentration to time for a specific reaction order.
Common Mistakes to Avoid
- Using coefficients from the balanced equation as reaction orders. This is wrong for most reactions because rate-law exponents must be determined experimentally unless the step is an elementary reaction.
- Changing two reactant concentrations at once when comparing initial-rate trials. This is wrong because it becomes unclear which reactant caused the rate change.
- Assuming k is always unitless. This is wrong because the units of k change with the overall reaction order so that rate has units of concentration per time.
- Confusing overall order with molecularity. Overall order comes from the experimental rate law, while molecularity describes the number of particles in a single elementary step.
Practice Questions
- 1 For the rate law Rate = k[A]^2[B], what is the overall reaction order, and what happens to the rate if [A] is doubled while [B] stays constant?
- 2 Initial-rate data show that when [A] doubles and [B] stays constant, the rate doubles. When [B] doubles and [A] stays constant, the rate quadruples. Write the rate law in the form Rate = k[A]^m[B]^n.
- 3 A student claims that the reaction 2NO + O2 -> 2NO2 must have the rate law Rate = k[NO]^2[O2]. Explain why this conclusion may not be valid without experimental data.