What you'll learn
The rate and extent of chemical change asks two questions about every reaction: how fast does it go, and how far does it go? This guide covers the explanations behind both. The first half uses collision theory to account for why concentration, pressure, surface area, temperature and catalysts change the rate — and the explanations are not interchangeable, since temperature works differently from the others. The second half deals with reversible reactions and equilibrium, where a reaction settles into a balance that can be deliberately shifted. By the end of this guide you should be able to state collision theory precisely, explain the effect of each factor in the terms examiners require, describe how catalysts work and why they are used industrially, explain dynamic equilibrium, and predict how changing concentration, temperature or pressure shifts the position of equilibrium.
Key terms and definitions
Rate of reaction — how quickly reactants are used up or products formed
Collision theory — the model stating that reactions occur only when particles collide with sufficient energy
Successful collision — a collision with energy equal to or greater than the activation energy, which results in reaction
Activation energy — the minimum energy colliding particles must have for a reaction to occur
Catalyst — a substance that increases the rate of reaction by providing an alternative pathway of lower activation energy, without being used up
Enzyme — a biological catalyst
Surface area to volume ratio — the relationship that makes smaller pieces of a solid react faster than larger pieces of the same mass
Reversible reaction — a reaction in which the products can react to reform the reactants
Dynamic equilibrium — the state in a closed system where forward and backward reactions occur at equal rates and concentrations stay constant
Closed system — a system from which no reactants or products can escape
Position of equilibrium — whether the equilibrium mixture contains more reactants or more products
Core concepts
Collision theory
Chemical reactions occur only when reacting particles collide with each other, and only when they collide with sufficient energy. The minimum energy needed is the activation energy, and a collision meeting it is a successful collision.
It follows that there are exactly two ways to increase the rate of any reaction: increase the frequency of collisions, or increase the proportion of collisions that are successful.
Every factor in this topic works through one of these two routes, and identifying which route is what earns the explanation mark. Writing that a factor makes particles "collide more" without specifying frequency or success is too vague to score.
Concentration
Increasing the concentration of a solution means there are more particles of the reactant in the same volume. The particles are therefore closer together and collide more frequently, so more successful collisions occur each second and the rate increases.
Note that concentration changes the frequency of collisions, not the proportion that succeed. The average energy of the particles is unchanged.
Pressure of gases
Increasing the pressure of a gas compresses the same number of particles into a smaller volume. As with concentration, the particles are closer together and collide more frequently, so the rate increases.
Pressure and concentration are therefore the same explanation applied to gases and solutions respectively, and either can be described as increasing the number of particles per unit volume.
Surface area
Breaking a solid into smaller pieces, or grinding it into a powder, increases the total surface area exposed to the other reactant while the mass stays the same.
Reaction can only occur at the surface of a solid, so a larger surface area means more particles of the solid are available to be collided with. Collisions are therefore more frequent and the rate increases.
The phrase examiners look for is surface area to volume ratio: a powder has a much higher ratio than a single lump of the same mass. This also explains an industrial hazard, since fine combustible powders such as flour or sawdust can ignite explosively when dispersed in air.
Temperature
Temperature is the factor that works differently, and a full-mark answer needs both of its effects.
Increasing the temperature gives the particles more kinetic energy, so they move faster and collide more frequently. This is the smaller of the two effects.
More importantly, a greater proportion of the particles now have energy equal to or greater than the activation energy, so a greater proportion of collisions are successful.
Answers that mention only the first effect typically score half the available marks. The activation energy point is the one that matters, and it is also what explains why a modest temperature rise can produce a large increase in rate.
Catalysts
A catalyst increases the rate of reaction by providing a different reaction pathway with a lower activation energy. Because the barrier is lower, a greater proportion of collisions have sufficient energy to react.
Three further properties are routinely examined. A catalyst is not used up in the reaction, so it can be recovered unchanged and reused. It does not appear in the equation for the reaction, because it is neither a reactant nor a product. And different reactions require different catalysts, so a catalyst is specific rather than general.
Enzymes act as catalysts in biological systems.
A catalyst does not change the products of a reaction, nor the yield of a reversible reaction, nor the overall energy change. It affects the rate only, which is precisely why it is shown on a reaction profile as a lower hump between unchanged reactant and product levels.
Industrially, catalysts matter economically. They allow an acceptable rate to be achieved at a lower temperature than would otherwise be needed, which reduces energy costs and the associated carbon dioxide emissions.
Reversible reactions
In some reactions the products can react to reform the original reactants. Such reactions are written with a double arrow rather than a single one.
The standard demonstration is the effect of heat on hydrated copper sulfate. Heating the blue hydrated crystals drives off water, leaving white anhydrous copper sulfate; adding water to the white solid turns it blue again. The forward reaction here is endothermic and the reverse exothermic.
This illustrates a general rule: if the forward reaction is exothermic, the backward reaction is endothermic, and exactly the same quantity of energy is transferred in each direction.
Dynamic equilibrium
When a reversible reaction occurs in a closed system, equilibrium is reached when the forward and backward reactions occur at exactly the same rate.
The word dynamic matters and is frequently the subject of a mark. At equilibrium both reactions continue to occur; nothing has stopped. Because they proceed at equal rates, the concentrations of reactants and products remain constant even though individual molecules continue to react in both directions.
A closed system is essential, since if a product escapes the backward reaction cannot keep pace and equilibrium is never established.
Equilibrium does not mean equal amounts of reactants and products. The position of equilibrium may lie well towards either side, and it depends on the conditions.
Changing the conditions
When a system at equilibrium experiences a change in conditions, the position of equilibrium shifts so as to counteract that change. Three changes are examinable.
Changing concentration: increasing the concentration of a reactant shifts the position towards the products, as the system responds by using up the added substance. Removing a product has the same effect.
Changing temperature: increasing the temperature shifts the position in the endothermic direction, because that direction absorbs the added energy. Decreasing it shifts the position in the exothermic direction. Identifying which direction is exothermic is therefore always the first step.
Changing pressure, for reactions involving gases: increasing the pressure shifts the position towards the side with fewer molecules of gas, since fewer molecules exert less pressure. Counting the gas molecules on each side of the equation is the whole of the working, and if both sides have equal numbers, pressure has no effect on the position.
These principles are applied directly in the Haber process, where the compromise conditions follow from exactly this reasoning.
Worked examples
Example 1: Explaining the effect of surface area (3 marks)
Explain why powdered calcium carbonate reacts faster with dilute acid than a single lump of the same mass.
Breaking the solid into a powder greatly increases its total surface area, and therefore its surface area to volume ratio, while the mass remains the same. Reaction can occur only where acid particles meet the surface of the solid, so a larger surface area means more solid particles are exposed and available for collision. Collisions between the acid particles and the solid are therefore more frequent, so the rate of reaction increases.
Example 2: Explaining the effect of temperature (4 marks)
Explain why increasing the temperature from 25 to 45 degrees Celsius substantially increases the rate of a reaction.
At the higher temperature the particles have greater kinetic energy, so they move faster and collide more frequently, which increases the rate to some extent.
More significantly, a greater proportion of the particles now have energy equal to or greater than the activation energy, so a greater proportion of the collisions that occur are successful. Because both the frequency of collisions and the proportion that succeed have increased, the number of successful collisions per second rises substantially and the rate increases sharply.
Example 3: Predicting equilibrium shifts (4 marks)
In a reversible reaction, the forward reaction is endothermic, and there are two molecules of gas on the left and three on the right. Predict the effect on the yield of product of increasing the temperature, and of increasing the pressure.
Increasing the temperature shifts the position of equilibrium in the endothermic direction. Here the forward reaction is endothermic, so the position shifts to the right and the yield of product increases.
Increasing the pressure shifts the position of equilibrium towards the side with fewer molecules of gas. Here the left has two molecules and the right has three, so the position shifts to the left and the yield of product decreases.
Common mistakes and how to avoid them
The most frequent error is writing that particles collide more, without saying whether more frequently or more successfully. The mechanism is the mark.
Students very often explain temperature by kinetic energy alone. The activation energy effect is the more important half and must be included.
Describing a catalyst as lowering the activation energy of the reaction is a subtle but penalised error. It provides an alternative pathway with a lower activation energy; the original pathway is unchanged.
In equilibrium questions, many answers state that the reaction stops once equilibrium is reached. Both reactions continue at equal rates, which is what dynamic means.
Another common slip is assuming equilibrium means equal concentrations. It means equal rates.
Finally, candidates sometimes apply the pressure rule without counting the gas molecules, or count all molecules including solids and liquids. Only gases count.
Exam technique for "The rate and extent of chemical change"
Answer every rate explanation in two parts: what happens to the particles, then the effect on collisions. For example, at higher concentration there are more particles in the same volume, so collisions are more frequent.
Reserve the phrase greater proportion of collisions have energy above the activation energy for temperature and catalysts only. Using it for concentration or surface area is wrong and is penalised.
Before answering any equilibrium question, write down which direction is exothermic and count the gas molecules on each side. That preparation answers every part that follows.
When a question asks about yield rather than rate, check whether the factor affects the position of equilibrium at all — a catalyst does not, and pressure does not when the number of gas molecules is equal on both sides.
Quick revision summary
Collision theory states that reactions occur when particles collide with at least the activation energy, so rate increases either through more frequent collisions or through a greater proportion being successful. Higher concentration and higher gas pressure place more particles in the same volume, increasing collision frequency; greater surface area exposes more of a solid, also increasing frequency. Higher temperature increases both frequency and, more importantly, the proportion of collisions with enough energy. A catalyst provides an alternative pathway of lower activation energy, is not used up, does not appear in the equation, is specific to a reaction, and changes rate but never yield or overall energy change. Reversible reactions run both ways, with the reverse direction having the opposite energy change of equal magnitude. In a closed system, dynamic equilibrium is reached when forward and backward rates are equal, with both reactions continuing and concentrations constant but not necessarily equal. Increasing a reactant's concentration shifts the position towards products, raising temperature shifts it in the endothermic direction, and raising pressure shifts it towards the side with fewer gas molecules.