A rate law is an experimentally determined equation (rate = k[A]^m[B]^n) showing how a reaction's rate depends on reactant concentrations, where k is the rate constant and the exponents are the reaction orders. Unlike equilibrium expressions, rate laws cannot be read off the overall balanced equation.
A rate law is the mathematical relationship between a reaction's rate and the concentrations of its reactants, written as rate = k[A]^m[B]^n. The k is the rate constant (which changes with temperature), and the exponents m and n are the reaction orders, which tell you how sensitive the rate is to each concentration. Double a first-order reactant and the rate doubles. Double a second-order reactant and the rate quadruples. Here's the rule that trips everyone up: for the overall reaction, those exponents come from experimental data, NOT from the coefficients in the balanced equation.
There's exactly one exception, and it's the bridge to reaction mechanisms (Topic 5.7). For an elementary step, the rate law CAN be written directly from the step's coefficients, because an elementary step describes a single actual molecular collision. That's why rate laws are the main evidence chemists use to support or reject a proposed mechanism. If the experimental rate law doesn't match what the mechanism predicts, the mechanism is wrong.
Rate laws live in Unit 5 (Kinetics), but their reach extends across the course. In Topic 5.7, learning objective AP Chem 5.7.A asks you to identify the components of a reaction mechanism, and the rate law is your tool for testing whether a series of elementary steps actually adds up to reality (EK 5.7.A.1, 5.7.A.2). In Unit 9, Topic 9.4 (AP Chem 9.4.A) asks you to explain why a thermodynamically favored reaction might not happen at a measurable rate. The answer is kinetics. A reaction under kinetic control has a high activation energy, which shows up as a tiny rate constant k, which makes the rate crawl no matter how favorable ฮG is (EK 9.4.A.2). And in Unit 7, equilibrium is defined as the dynamic state where forward and reverse rates are equal (AP Chem 7.1.A), so rate laws are the kinetic machinery underneath every equilibrium expression you write. Rate laws are how AP Chem connects "how fast" to "how far."
Keep studying AP Chemistry Unit 5
Reaction Mechanisms (Unit 5)
For an elementary step, the rate law comes straight from the step's coefficients because that step represents one real collision. Comparing the experimental rate law to the rate law predicted by the slow step is how you confirm or reject a proposed mechanism, alongside detecting intermediates (EK 5.7.A.4).
Rate Constant and Order of Reaction (Unit 5)
The rate law is built from these two pieces. The order tells you how rate responds to concentration changes, and the rate constant k carries the temperature and activation-energy dependence. Heating a reaction doesn't change the orders, it changes k.
Thermodynamic and Kinetic Control (Unit 9)
A reaction can be thermodynamically favored and still go nowhere, like diamond converting to graphite. High activation energy means a tiny k, so the rate law spits out a rate near zero. That's kinetic control, and it's why "favored" never means "fast" (AP Chem 9.4.A).
Chemical Equilibrium (Unit 7)
Equilibrium is dynamic, not frozen. It's the point where the forward rate law and the reverse rate law produce equal rates, so concentrations stop changing even though both reactions keep running (EK 7.1.A.3).
Rate laws are a multiple-choice and FRQ staple. MCQ stems give you a table of initial concentrations and initial rates and ask you to find the order with respect to each reactant, or they give you a mechanism and ask for the rate law of an elementary step. Practice questions hit exactly these skills, like "How do you determine the rate law from an elementary step?" and identifying what makes a reaction elementary. On FRQs, expect to determine a rate law from experimental data, calculate k with correct units, or evaluate whether a proposed mechanism is consistent with an observed rate law. You may also need to explain in writing why a thermodynamically favored reaction is slow, which means connecting a high activation energy to a small rate constant. One classic trap to dodge: writing the overall rate law from the balanced equation's coefficients. That only works for elementary steps.
Both look like concentrations raised to powers, but they come from completely different places. The equilibrium expression K is always written from the coefficients of the balanced overall equation and tells you how far a reaction goes. The rate law must be found experimentally (except for elementary steps) and tells you how fast it goes. A reaction with a huge K can still have a tiny rate, which is the whole point of kinetic control in Topic 9.4.
A rate law has the form rate = k[A]^m[B]^n, where k is the rate constant and the exponents are the reaction orders.
For an overall reaction, the orders in the rate law must come from experimental data, never from the balanced equation's coefficients.
For an elementary step only, you can write the rate law directly from that step's coefficients, because the step represents a single molecular event.
Matching an experimental rate law to a mechanism's predicted rate law is key evidence for or against that mechanism.
A thermodynamically favored reaction can still be extremely slow if activation energy is high and k is tiny, which is called kinetic control.
At equilibrium, the forward and reverse rates are equal, so rate laws are the kinetic foundation underneath the equilibrium state.
A rate law is an equation, rate = k[A]^m[B]^n, that shows how a reaction's rate depends on reactant concentrations. The rate constant k and the orders m and n are found from experiments, usually by comparing initial rates across trials.
No, not for an overall reaction. The orders must be determined experimentally. The only exception is an elementary step in a mechanism, where the coefficients of that single step do give the rate law.
A rate law describes how fast a reaction goes and must be found experimentally, while the equilibrium expression K describes how far it goes and is always written from the balanced equation's coefficients. A reaction can have a large K but a very slow rate.
No. If activation energy is high, the rate constant k is tiny and the reaction barely proceeds, even with a negative ฮG. AP Chem calls this kinetic control (AP Chem 9.4.A), and diamond turning into graphite is the classic example.
Each elementary step has a predictable rate law from its own coefficients, and the slow step controls the overall rate. If the mechanism's predicted rate law matches the experimental one, and any intermediates can be detected, the mechanism is supported (EK 5.7.A.4).