What you'll learn
The rate and extent of chemical change is the unit of AQA GCSE Combined Science: Trilogy that asks two questions about every reaction: how fast does it go, and how far does it go? The first half covers rates of reaction, explained through collision theory and the factors that change them. The second half covers reversible reactions and equilibrium, where a reaction does not go to completion but settles into a balance that can be shifted by changing the conditions. By the end you should be able to calculate a mean rate of reaction from a graph or from measurements, explain the effect of concentration, pressure, surface area, temperature and catalysts using collision theory, draw and interpret rate graphs including drawing a tangent, describe what dynamic equilibrium means, and predict how changing concentration, temperature or pressure shifts an equilibrium. This unit is assessed on Chemistry Paper 2 and includes the required practical on investigating rates.
Key terms and definitions
Rate of reaction — how quickly a reactant is used up or a product is formed, measured per unit time
Collision theory — the model stating that reactions occur only when particles collide with sufficient energy
Activation energy — the minimum amount of energy that colliding particles must have for a reaction to occur
Catalyst — a substance that speeds up a reaction without being used up, by providing a different pathway with a lower activation energy
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 original reactants
Equilibrium — the state reached in a closed system when the forward and backward reactions occur at exactly the same rate
Dynamic equilibrium — equilibrium in which both reactions continue but the concentrations of reactants and products remain constant
Closed system — a system in which no reactants or products can escape
Tangent — a straight line touching a curve at one point, used to find the rate at that moment
Core concepts
Measuring the rate of reaction
The mean rate of a reaction is calculated as the quantity of reactant used divided by the time taken, or the quantity of product formed divided by the time taken. The quantity may be a mass in grams, giving a rate in grams per second, or a volume of gas in cubic centimetres, giving cubic centimetres per second.
There are two standard methods of following a reaction. The first measures the mass lost as a gas escapes, by standing the reaction flask on a balance and recording the mass at intervals. The second measures the volume of gas produced, collected in a gas syringe or an inverted measuring cylinder over water. A third approach, used where a precipitate forms, times how long a cross marked on paper beneath the flask takes to disappear from view as the mixture turns cloudy.
Reading a rate graph
Plotting the quantity of product against time gives a curve that is steepest at the start and gradually levels off. The steepness at any point is the rate at that moment.
The curve is steepest at the beginning because the concentration of reactants is highest, so collisions are most frequent. As the reaction proceeds, reactants are used up, collisions become less frequent and the curve flattens. When the line becomes horizontal the reaction has stopped, because at least one reactant has been completely used up.
To find the rate at a particular time, draw a tangent to the curve at that point and calculate its gradient: the change in the vertical value divided by the change in the horizontal value. For a mean rate over a period, simply divide the total change by the total time.
Comparing two curves on the same axes is a common question. A steeper curve means a faster reaction. If both level off at the same height, the same amount of product was made, so the same quantity of reactant was used — only the speed differed.
Collision theory
Chemical reactions occur when reacting particles collide with each other and with sufficient energy. The minimum energy needed is the activation energy. Increasing the rate therefore requires either more frequent collisions or a greater proportion of collisions having enough energy.
Every factor in this unit works through one of those two routes, and saying which one is what earns the explanation mark.
Factors affecting the rate
Increasing the concentration of a solution means there are more particles in the same volume, so collisions are more frequent and the rate increases. Increasing the pressure of a gas does the same thing by the same reasoning: the particles are pushed into a smaller volume, so they are closer together and collide more often.
Increasing the surface area of a solid by breaking it into smaller pieces exposes more of the solid to the other reactant. The same mass of a powder has a far larger surface area than a single lump, so collisions with the surface are more frequent and the rate increases.
Increasing the temperature has two effects, and a full-mark answer usually needs both. Particles move faster, so collisions are more frequent; and more importantly, a greater proportion of particles have energy equal to or greater than the activation energy, so a greater proportion of collisions are successful.
A catalyst provides a different reaction pathway with a lower activation energy. This means a greater proportion of collisions now have enough energy to react. A catalyst is not used up and does not appear in the equation for the reaction, so it can be recovered unchanged at the end and reused. Different reactions need different catalysts, and enzymes act as catalysts in biological systems.
Reversible reactions
In some reactions the products can react to reform the original reactants. Such a reaction is written with a double arrow rather than a single one.
A standard example is the effect of heat on hydrated copper sulfate. Heating the blue hydrated crystals drives off water to leave white anhydrous copper sulfate, and adding water to the white solid turns it blue again. The forward reaction here is endothermic, taking in energy, and the reverse is exothermic.
This illustrates a general principle: if the forward reaction is exothermic, the backward reaction is endothermic, and exactly the same quantity of energy is transferred in each direction.
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. At equilibrium both reactions are still happening; it is not that everything has stopped. Because they proceed at equal rates, the concentrations of reactants and products stay constant, but individual molecules continue to react in both directions. A closed system is essential, because if a product can escape the reverse reaction cannot keep pace.
Equilibrium does not mean equal amounts. The position of equilibrium may lie towards the products or towards the reactants, and it depends on the conditions.
Changing the conditions
If a system at equilibrium is subjected to a change in conditions, the position of equilibrium shifts so as to counteract that change. Three changes are examinable.
Changing concentration: if the concentration of a reactant is increased, more products are formed until equilibrium is reached again, so the position shifts towards the products. Removing a product has the same effect, since the system responds by making more of it.
Changing temperature: if the temperature is increased, the position shifts in the endothermic direction, because that direction absorbs the added energy. If the temperature is decreased, it shifts in the exothermic direction. This is why knowing which direction is exothermic is always the first step in a temperature question.
Changing pressure, for reactions involving gases: increasing the pressure shifts the equilibrium towards the side with the smaller number of gas molecules, since fewer molecules exert less pressure. Counting the molecules on each side of the equation is therefore the whole of the working.
Worked examples
Example 1: Calculating a mean rate (3 marks)
A reaction produces 48 cubic centimetres of gas in 40 seconds. Calculate the mean rate of reaction.
The mean rate is the quantity of product formed divided by the time taken, which is 48 divided by 40. This gives 1.2 cubic centimetres per second. The unit must be included, and it comes directly from the two quantities used.
Example 2: Explaining the effect of temperature (4 marks)
Explain why increasing the temperature from 20 to 40 degrees Celsius increases the rate of a reaction.
At the higher temperature the particles have more kinetic energy, so they move faster and collide more frequently. 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 are successful. Both effects increase the number of successful collisions per second, which increases the rate.
Example 3: Predicting an equilibrium shift (4 marks)
In a reversible reaction, the forward reaction is exothermic and there are four molecules of gas on the left and two on the right. Predict the effect of increasing the temperature and of increasing the pressure.
Increasing the temperature shifts the position of equilibrium in the endothermic direction, which here is the backward reaction, so the yield of product decreases. Increasing the pressure shifts the position towards the side with fewer gas molecules, which is the right-hand side with two molecules, so the yield of product increases.
Common mistakes and how to avoid them
The most common error in rate explanations is writing that particles collide more, without saying more frequently or more successfully. The mark is for the mechanism, and more collisions is too vague to earn it.
Students very often say that raising the temperature gives particles more energy so they react faster. That misses the activation energy point, which is the more important half. Always mention the proportion of particles with energy above the activation energy.
Describing a catalyst as lowering the activation energy of the reaction is a subtle but penalised error. The catalyst provides an alternative pathway with a lower activation energy; the original pathway is unchanged.
In equilibrium questions, many answers state that the reaction stops. It does not; both reactions continue at equal rates, which is why the word dynamic is used.
Finally, students frequently assume equilibrium means equal concentrations of reactants and products. It means equal rates, not equal amounts.
Exam technique for "The Rate and Extent of Chemical Change"
For any rate explanation, use a two-part sentence: state what happens to the particles, then state the effect on collisions. For example, at higher concentration there are more particles in the same volume, so collisions are more frequent.
When drawing a tangent, make it long enough to read the gradient accurately, and choose two points far apart on the line for the calculation. A short tangent is the main cause of inaccurate answers.
For the required practical, be ready to identify the independent variable, usually concentration or temperature, the dependent variable, usually time taken or volume of gas, and the control variables, which include the volume and concentration of the other reactant and the mass and form of any solid.
For equilibrium questions, work in a fixed order: identify which direction is exothermic, count the gas molecules on each side, then answer whichever part is asked. Doing this preparation first makes every part of the question straightforward.
Quick revision summary
Rate is the quantity of reactant used or product formed divided by time, followed by mass loss, gas volume or a disappearing cross. Rate graphs are steepest at the start, flatten as reactants are used up, and become horizontal when the reaction stops; the gradient of a tangent gives the rate at a moment. Collision theory says particles must collide with at least the activation energy. Higher concentration, higher pressure and greater surface area all increase the frequency of collisions; higher temperature increases both frequency and the proportion of collisions with enough energy; a catalyst provides an alternative pathway with a lower activation energy and is not used up. Reversible reactions run both ways, and if the forward reaction is exothermic the reverse is endothermic by the same quantity. In a closed system, dynamic equilibrium is reached when forward and backward rates are equal. Raising temperature shifts the equilibrium in the endothermic direction, raising pressure shifts it towards fewer gas molecules, and increasing a reactant concentration shifts it towards the products.