What factors affect the rate of a chemical reaction?

Concentration, temperature, surface area, catalysts and light (for photochemical reactions). Each either increases the frequency of particle collisions or the proportion of collisions with sufficient energy.

What is chemical equilibrium?

Equilibrium is reached in a reversible reaction when the rate of the forward reaction equals the rate of the reverse reaction. Concentrations of reactants and products remain constant but the reaction has not stopped.

What is Le Chatelier's principle?

If a system at equilibrium is subjected to a change (concentration, pressure, temperature), it will shift in the direction that partially opposes that change.

What is activation energy and how do catalysts affect it?

Activation energy (Ea) is the minimum energy colliding particles must have for a reaction to occur. Catalysts provide an alternative pathway with a lower Ea, increasing the fraction of particles that can react.

GUIDE Chemistry Intermediate

Reaction Rates and Equilibrium

Why some reactions are explosive and others take centuries — and what determines where a reversible reaction settles when left to reach equilibrium.

Collision theory → Back to all guides

Primary sources: Khan Academy AP Chemistry · RSC Education Resources

01 · Collision theory

Reactions need the right collision.

For a reaction to occur, reacting particles must:

Collide — particles must come into contact.
Have sufficient energy — the combined kinetic energy must meet or exceed the activation energy (Ea).
Have the correct orientation — reactive parts of the molecules must face each other.

The Maxwell-Boltzmann distribution describes how kinetic energies are spread across a population of molecules at a given temperature. The area under the curve to the right of Ea represents the fraction of molecules that can react. A catalyst shifts Ea to the left, increasing that fraction. See our catalyst glossary entry.

02 · Factors affecting rate

Four variables, one explanation.

All four factors work through collision theory — they either change how often collisions occur or what fraction of them succeed.

[A]

Concentration

More particles per litre → more frequent collisions → faster rate.

Temperature

Higher T → greater average kinetic energy → more particles exceed Ea. Rate roughly doubles per 10 °C.

Surface area

Smaller particle size → greater exposed surface → more collisions per unit time.

Catalyst

Provides lower-energy reaction pathway → larger fraction of particles can react at the same temperature.

03 · Equilibrium

Reversible reactions reach a balance.

Many reactions are reversible (written with ⇌): the products can react to reform the reactants. As the reaction proceeds, the forward rate slows (reactants depleted) and the reverse rate increases (products build up) until they are equal. At this point — dynamic equilibrium — concentrations no longer change, but both reactions continue at equal rates.

The equilibrium constant K is defined as:

Kc = [products]^n / [reactants]^m

(where exponents are stoichiometric coefficients). K depends only on temperature — not on initial concentrations. A large K means products are favoured; a small K means reactants are favoured.

04 · Le Chatelier's principle

Systems oppose change — the Haber process.

Le Chatelier's principle states that a system at equilibrium shifts to oppose any imposed change in conditions. This guides the design of industrial chemical processes:

Haber process: N₂ + 3H₂ ⇌ 2NH₃ (ΔH = −92 kJ mol⁻¹, exothermic). High pressure (150–300 atm) shifts equilibrium right (fewer gas moles on product side). Low temperature would favour NH₃ but slow the rate — a compromise of ~450 °C with an iron catalyst is used industrially.
Contact process: 2SO₂ + O₂ ⇌ 2SO₃. High pressure and V₂O₅ catalyst (see catalyst glossary) at ~450 °C. Excess O₂ shifts equilibrium right.

Quantum computing may soon allow us to simulate reaction energy surfaces and design better catalysts computationally — see our quantum computing and chemistry article.

Questions about reaction rates and equilibrium

: Concentration — more particles per unit volume → more collisions → faster rate.
Temperature — higher T → particles move faster → more frequent and more energetic collisions. Rate roughly doubles per 10 °C rise (Q10 ≈ 2).
Surface area — smaller particles expose more surface → faster reaction.
Catalyst — lowers activation energy; see our catalyst glossary entry.
All explained by collision theory and the Maxwell-Boltzmann distribution.
: In a reversible reaction (⇌), both forward and reverse reactions occur simultaneously. Equilibrium is the point where the two rates are equal, so concentrations no longer change. It is a dynamic equilibrium — both reactions continue, but at equal rates. The equilibrium position can be expressed by the equilibrium constant K, which depends only on temperature.
: Le Chatelier's principle (1884): a system at equilibrium will respond to a disturbance by shifting to reduce the effect of that change. Increase reactant concentration → equilibrium shifts right (toward products). Increase pressure → equilibrium shifts toward fewer moles of gas. Increase temperature → equilibrium shifts in the endothermic direction. This underpins industrial processes like the Haber process (N₂ + 3H₂ ⇌ 2NH₃).
: The activation energy (Ea) is the energy barrier that must be overcome. At a given temperature, only the fraction of particles with energy ≥ Ea (the area under the Maxwell-Boltzmann curve to the right of Ea) can react. A catalyst lowers Ea, increasing this fraction — and therefore the rate — without being consumed. See our catalyst glossary entry and the enzymes guide for biological catalysts.

Catalysts: biological and industrial

Catalysts are central to reaction rates. The catalyst glossary entry compares enzymes with industrial catalysts like platinum and vanadium(V) oxide.

Catalyst glossary →