R2.2
How fast? The rate of chemical change
9 SL · +6 HL (SL 2.2.1-2.2.5) (HL 2.2.6-2.2.13)
Unverified
← View on the mapParent topic: How much, how fast and how far?
Guiding question
Guiding questionHow can the rate of a reaction be controlled?
The one big idea
Reaction rate depends on collision frequency, energy and orientation. At HL, the rate equation, reaction order and the Arrhenius equation turn these qualitative ideas into quantitative predictions.
What this rests on
Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.
- EditorialS1.4 Counting particles by mass: The moleConcentration calculations for reaction rates are mole calculations.
- EditorialS3.1 The periodic table: Classification of elementsHL onlyTransition elements are useful as catalysts; their features come from periodic position (HL linking question R2.2.5).
- EditorialS1.1 Introduction to the particulate nature of matterCollision theory is kinetic molecular theory applied to reactions (explicit linking question).
- EditorialR1.1 Measuring enthalpy changesEnergy profiles and activation energy extend the enthalpy-change energy profile.
What rests on this
Core concepts that must be mastered
- Rate of reaction and collision theory(2.2.1, 2.2.2)Rate is the change in concentration of a reactant or product per unit time. You determine rates from concentration/volume/mass vs time graphs — the instantaneous rate at a point is the gradient of the tangent to the curve at that point. The rate decreases as the reaction proceeds because reactant concentrations fall. Concentration changes are not usually measured directly; methods include collecting gas volume over time, mass loss, colorimetry, pH change, and conductivity. Species react via collisions of sufficient energy and proper orientation. The relationship between particle kinetic energy and temperature (in kelvin) comes from kinetic molecular theory (S1.1): higher temperature means higher average kinetic energy. But not every collision leads to reaction — the colliding particles need enough energy to break bonds (the activation energy) and the right geometry. Temperature is proportional to average kinetic energy, but the distribution of energies matters more than the average for reaction rate — this is the Maxwell–Boltzmann insight.
- Factors affecting rate and activation energy(2.2.3, 2.2.4)Rate is influenced by pressure (gases), concentration (solutions), surface area (solids), temperature and catalyst. Increasing pressure or concentration increases the frequency of collisions; increasing surface area exposes more particles to reactant; increasing temperature increases both collision frequency and, more importantly, the fraction of collisions with energy ≥ Ea; a catalyst provides an alternative pathway with lower Ea. The activation energy Ea is the minimum energy colliding particles need to react. You must construct Maxwell–Boltzmann energy distribution curves: x-axis = kinetic energy, y-axis = number of particles. The curve starts at the origin (no particles have zero energy), rises to a peak, and tails off. The area under the curve beyond Ea represents the fraction of particles with enough energy to react. When temperature increases, the peak moves right and down (the average energy rises) and the curve broadens — the total area stays constant (same number of particles), but the area beyond Ea increases dramatically. This is why a small temperature increase can cause a large rate increase: the dominant effect is not that particles move slightly faster, but that the fraction with enough energy to react rises exponentially.
- Catalysts(2.2.5)Catalysts increase rate by providing an alternative pathway with lower Ea. A catalyst does not lower Ea of the uncatalysed pathway — the original pathway still exists. You must sketch and explain energy profiles with and without a catalyst: the catalysed profile has a smaller hill but the same start and end points. A catalyst does not change ΔH (a state function) or the equilibrium position — it speeds the approach to equilibrium in both directions equally. Biological catalysts are enzymes. Homogeneous vs heterogeneous catalyst mechanisms are not assessed.
- HL: Reaction mechanisms and the rate-determining step(2.2.6–2.2.8)Many reactions occur in elementary steps; the slowest step — the rate-determining step (RDS) — determines the rate. You must evaluate proposed mechanisms and recognise intermediates (species that form in one step and are consumed in a later step) and transition states (high-energy maxima between intermediates). The RDS can be any step, not necessarily the first. A mechanism must be consistent with both the kinetic data (the rate equation) and the stoichiometric data (the overall equation). Energy profiles of multistep reactions show multiple hills, with Ea and the transition state of the RDS visible. The molecularity of an elementary step is the number of reacting particles in it: unimolecular (one), bimolecular (two), termolecular (three — rare).
- HL: Rate equations, order and the rate constant(2.2.9–2.2.11)The rate equation depends on the mechanism and can only be determined experimentally — the order is not the coefficient in the balanced equation. The order with respect to a reactant is its exponent in the rate equation; the overall order is the sum of the individual orders. For rate = k[A]^m^[B]^n^, the order with respect to A is m, with respect to B is n, and the overall order is m + n. Only integer orders (0, 1, 2) are assessed. You must analyse concentration–time and rate–concentration graphs to identify zero, first and second order reactions. The rate constant k is temperature-dependent; its units follow from the overall order. For overall order n, the units of k are (mol dm⁻³)^(1−n)^ s⁻¹.
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HL: The Arrhenius equation (2.2.12, 2.2.13): The Arrhenius equation uses the temperature dependence of k to determine Ea: k = Ae^(−Ea/RT)^. The qualitative relationship: k increases with T. The linear form is ln k = (−Ea/R)(1/T) + ln A — a plot of ln k vs 1/T gives a straight line with gradient −Ea/R and intercept ln A. The Arrhenius factor A accounts for the frequency of collisions with proper orientation. You must determine Ea and A from experimental data, using both forms. Both forms are in the data booklet.
- How this sub-topic connectsKinetics is the "how fast" half of Reactivity 2, paired with equilibrium's "how far" (R2.3). Collision theory extends kinetic molecular theory (S1.1) into reactive encounters. The Maxwell–Boltzmann distribution is the key diagram. A catalyst appears in energy profiles (R1.1) and in equilibrium (R2.3: a catalyst does not change K). At HL, the rate equation connects to reaction mechanisms (R3.4: the SN1 and SN2 mechanisms of halogenoalkanes have different rate equations — SN1 is first order, SN2 is second order) and to the Nature of Science theme (mechanisms are "possible" models, not certainties). The Arrhenius equation's linear form connects to Tool 3 (linearising data for graphical analysis). Alkanes are kinetically stable (high Ea) but thermodynamically unstable — this explains why they do not spontaneously combust despite a negative ΔG (R3.3, R1.4).
Quantitative non-negotiables
- Determine rate from concentration/volume/mass vs time graphs, including by drawing tangents.
- Construct Maxwell–Boltzmann energy distribution curves and use them to explain the effect of temperature and catalyst on rate.
- Sketch energy profiles with and without a catalyst for endo- and exothermic reactions.
- At HL: deduce the rate equation from experimental data and determine the order with respect to each reactant.HL
- At HL: determine the units of k from the overall order.HL
- At HL: apply the Arrhenius equation (both forms, including the linear form) to determine E_a and A.HL
Common failure modes
M-32 — A catalyst changes ΔH.
Confidence: verified
Why it’s wrong: A catalyst lowers Ea by providing an alternative pathway; ΔH (a state function) is unchanged.
Correction: A catalyst does not change ΔH or the equilibrium position.
M-38 — The order of reaction is the coefficient in the balanced equation.
Confidence: verified
Why it’s wrong: Orders are determined only by experiment.
Correction: They reflect the mechanism up to and including the rate-determining step. They may be zero, fractional, or unrelated to the coefficients.
M-39 — Increasing temperature speeds a reaction mainly because particles move faster and collide more often.
Confidence: verified
Why it’s wrong: Collision frequency rises only modestly (roughly with √T).
Correction: The dominant effect is the exponential increase in the fraction of collisions with energy ≥ Ea, visible as the area under the tail of the Maxwell–Boltzmann curve.
M-40 — A catalyst lowers the activation energy of the reaction.
Confidence: verified
Why it’s wrong: A catalyst provides an alternative pathway with a lower Ea; the uncatalysed pathway still exists.
Correction: Say 'provides an alternative pathway with lower Ea', not 'lowers the activation energy'.
M-41 — The rate constant k changes when you change the concentration.
Confidence: verified
Why it’s wrong: k depends on temperature (and the catalyst/pathway) only.
Correction: Changing concentration changes the rate, not k.
M-42 — The Maxwell–Boltzmann curve shifts up as well as right when temperature increases.
Confidence: verified
Why it’s wrong: The total area (total number of particles) is constant.
Correction: The peak moves right and down, and the curve broadens.
M-43 — The rate-determining step is the one with the biggest activation energy in the profile, so it must be the first step.
Confidence: likely
Why it’s wrong: The RDS is the slowest step, which is the one with the highest activation energy barrier from its preceding intermediate.
Correction: It can be anywhere in the sequence.
What "HL standard" actually looks like
HLHL adds multi-step mechanisms with intermediates vs transition states, molecularity, experimentally-determined rate equations and orders (0, 1, 2 — integer only), the units of k, the Arrhenius equation (including the linear form and factor A), and deducing rate equations from the rate-determining step. +6 HL hours — the second-heaviest HL block in Reactivity.
Linking questions
Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.
- Official IB
What is the graphical distribution of kinetic energy values of particles in a sample at a fixed temperature?
- Official IB
What must happen to particles for a chemical reaction to occur?
- Official IBwording unverified against the guide
Graphs can be presented as sketches or as accurately plotted data points. What are the advantages and limitations of each representation?
- Official IB
How does the presence of double and triple bonds in molecules influence their reactivity?
- Official IB
Why is high activation energy often considered to be a useful property of a fuel?
- Official IB
What is the relationship between the kinetic molecular theory and collision theory?
- Official IB
What is the relative effect of a catalyst on the rate of the forward and backward reactions?
- Official IB
What are the features of transition elements that make them useful as catalysts?
- Official IB
Which mechanism in the hydrolysis of halogenoalkanes involves an intermediate?
- Official IB
What are the rate equations and units of k for the reactions of primary and tertiary halogenoalkanes with aqueous alkali?
- Official IB
Why do catalysts have no effect on the value of K or on the equilibrium composition?
- Official IB
Why are alkanes described as kinetically stable but thermodynamically unstable?
- Official IB
What differences would be expected between the energy profiles for SN1 and SN2 reactions?
- Official IB
What are the rate equations for these SN1 and SN2 reactions?
- Official IB
How useful are mechanistic models such as SN1 and SN2?
6 further official linking questions target Tool, Inquiry or Nature of Science strands and are not drawn as edges on the map:
Editorial The decision to surface these off-graph questions here is our editorial judgement — the IB does not prescribe where they should appear.
- Official IB
Concentration changes in reactions are not usually measured directly. What methods are used to provide data to determine the rate of reactions?
- Official IB
What experiments measuring reaction rates might use time as i) a dependent variable ii) an independent variable?
- Official IB
What variables must be controlled in studying the effect of a factor on the rate of a reaction?
- Official IB
How can graphs provide evidence of systematic and random error?
- Official IB
What measurements are needed to deduce the order of reaction for a specific reactant?
- Official IB
Why are reaction mechanisms only considered as "possible mechanisms"?
Bridge: GCSE → IB HL
Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.