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Chemical Kinetics & Reaction Rates cheat sheet - grade 11-12

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Chemistry Grade 11-12

Chemical Kinetics & Reaction Rates Cheat Sheet

A printable reference covering reaction rates, rate laws, integrated rate laws, half-life, activation energy, and catalysts for grades 11-12.

TopicsChemical ReactionsMoleculesEnergyTimeStudy GuideChemistry
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Chemical kinetics studies how fast reactions occur and what factors control their speeds. This cheat sheet helps students connect concentration changes, rate laws, graphs, and mechanisms in one reference. It is especially useful for solving lab data problems, interpreting energy diagrams, and comparing reaction conditions.

Grade 11-12 chemistry students need these tools to move from memorizing formulas to explaining why reactions speed up or slow down.

The most important ideas are rate, rate law, reaction order, integrated rate laws, and activation energy. A rate law such as rate=k[A]m[B]n\text{rate} = k[A]^m[B]^n shows how reactant concentrations affect speed, while integrated laws show how concentration changes over time. The Arrhenius equation, k=Ae−Ea/(RT)k = Ae^{-E_a/(RT)}, connects temperature to the rate constant.

Catalysts increase reaction rate by lowering EaE_a without being consumed or changing the overall reaction energy.

Key Facts

  • For aA+bB→pPaA + bB \rightarrow pP, the average reaction rate is rate=−1aΔ[A]Δt=1pΔ[P]Δt\text{rate} = -\frac{1}{a}\frac{\Delta[A]}{\Delta t} = \frac{1}{p}\frac{\Delta[P]}{\Delta t}.
  • A general rate law has the form rate=k[A]m[B]n\text{rate} = k[A]^m[B]^n, where mm and nn must be found from experimental data.
  • The overall reaction order is m+nm+n for the rate law rate=k[A]m[B]n\text{rate} = k[A]^m[B]^n.
  • A zero-order reaction follows [A]t=−kt+[A]0[A]_t = -kt + [A]_0 and has half-life t1/2=[A]02kt_{1/2} = \frac{[A]_0}{2k}.
  • A first-order reaction follows ln⁡[A]t=−kt+ln⁡[A]0\ln[A]_t = -kt + \ln[A]_0 and has half-life t1/2=0.693kt_{1/2} = \frac{0.693}{k}.
  • A second-order reaction in one reactant follows 1[A]t=kt+1[A]0\frac{1}{[A]_t} = kt + \frac{1}{[A]_0} and has half-life t1/2=1k[A]0t_{1/2} = \frac{1}{k[A]_0}.
  • The Arrhenius equation is k=Ae−Ea/(RT)k = Ae^{-E_a/(RT)}, where increasing TT usually increases kk by giving more particles enough energy to react.
  • A catalyst lowers activation energy EaE_a and increases rate, but it does not change ΔH\Delta H, ΔG\Delta G, or the equilibrium constant KK.

Vocabulary

Reaction rate
Reaction rate is the change in concentration of a reactant or product per unit time, often written as Δ[substance]Δt\frac{\Delta[\text{substance}]}{\Delta t}.
Rate law
A rate law is an experimentally determined equation, such as rate=k[A]m[B]n\text{rate} = k[A]^m[B]^n, that relates reaction rate to reactant concentrations.
Rate constant
The rate constant kk is the proportionality factor in a rate law, and its value depends on temperature and the reaction pathway.
Reaction order
Reaction order is the exponent on a concentration term in a rate law, and the overall order is the sum of all exponents.
Activation energy
Activation energy EaE_a is the minimum energy needed for reacting particles to form the activated complex and proceed to products.
Catalyst
A catalyst is a substance that speeds up a reaction by providing a lower-energy pathway without being consumed overall.

Common Mistakes to Avoid

  • Using stoichiometric coefficients as rate-law exponents, which is wrong because exponents in rate=k[A]m[B]n\text{rate} = k[A]^m[B]^n must come from experimental data unless the step is elementary.
  • Forgetting the negative sign for reactant disappearance, which is wrong because Δ[A]Δt\frac{\Delta[A]}{\Delta t} is negative when a reactant is being consumed.
  • Choosing the wrong integrated rate law, which is wrong because zero-order uses [A][A] vs. tt, first-order uses ln⁡[A]\ln[A] vs. tt, and second-order uses 1[A]\frac{1}{[A]} vs. tt.
  • Treating kk as unitless, which is wrong because the units of kk change with overall reaction order so that rate units remain M s−1\text{M}\,\text{s}^{-1}.
  • Saying a catalyst changes the amount of product at equilibrium, which is wrong because a catalyst changes the rate of reaching equilibrium but not KK or the final equilibrium composition.

Practice Questions

  1. 1 For 2NO2(g)→2NO(g)+O2(g)2NO_2(g) \rightarrow 2NO(g) + O_2(g), [NO2][NO_2] decreases from 0.100 M0.100\ \text{M} to 0.074 M0.074\ \text{M} in 50.0 s50.0\ \text{s}. What is the average reaction rate?
  2. 2 A first-order reaction has k=0.0462 s−1k = 0.0462\ \text{s}^{-1} and [A]0=0.800 M[A]_0 = 0.800\ \text{M}. Find [A]t[A]_t after 30.0 s30.0\ \text{s} and calculate t1/2t_{1/2}.
  3. 3 Initial-rate data show that doubling [A][A] doubles the rate, while tripling [B][B] increases the rate by a factor of 99. Write the rate law and identify the overall order.
  4. 4 Explain why raising temperature and adding a catalyst can both increase reaction rate, but only the catalyst creates a different reaction pathway.

Understanding Chemical Kinetics & Reaction Rates

At the particle level, a reaction happens only when particles meet in a useful way. They need enough energy to disturb old bonds, and their collision must have a suitable orientation. Many collisions fail for one of these reasons.

This explains why a reaction can contain plenty of reactants yet proceed slowly. Most reactions occur through several smaller steps, not one giant event.

One slow step can limit the speed of the whole sequence, much like a narrow doorway limits movement through a crowded room. A proposed mechanism must agree with the measured rate behavior.

Experimental data are used to determine how concentration affects speed. In an initial rates experiment, chemists compare trials where one reactant concentration changes while other conditions stay fixed. This makes it possible to see the separate effect of each reactant.

The powers in a rate law are not usually taken from the balanced chemical equation. They describe what the experiment reveals about the slow parts of the process. Graphs provide another useful route.

A straight concentration versus time graph, a straight natural log of concentration versus time graph, or a straight reciprocal concentration versus time graph points to a different reaction order. Students should check graph axes carefully before choosing a model.

Temperature changes reaction speed because particles have a spread of kinetic energies. Heating does not give every particle the same extra energy. Instead, it increases the fraction able to cross the energy barrier during collisions.

This can cause a large rate change from a modest temperature increase. A catalyst provides a different sequence of steps with a lower barrier. In living things, enzymes act as highly specific catalysts.

Their shape holds reacting molecules in positions that make reaction more likely. Temperature, acidity, or substances that block the active site can reduce enzyme activity. A catalyst can help a system reach equilibrium sooner, but it does not decide which side of an equilibrium is favored.

Reaction rates matter outside the classroom. Refrigeration slows chemical changes that spoil food. Rusting, battery discharge, medicine breakdown, combustion, and industrial production all depend on controlled rates.

In a lab, accurate timing and consistent mixing matter because concentration can change rapidly at the start. Temperature must be controlled, since even small differences can affect results. When solving problems, keep reaction order separate from stoichiometric coefficients.

Notice whether a question gives data from several trials, a concentration time graph, or an energy diagram. Each source gives different evidence.

Units are useful checks too. The units of a rate constant change with reaction order, so they can expose an incorrect answer.