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How to measure and calculate a rate of reaction, what collision theory says about the factors that change it, and how catalysts work.
Specification references this topic covers: AQA 5.6.1; Edexcel C2.3, C4.4, C7.1, C7.3, C7.4, C7.5, C7.6, C7.7.
Some reactions are over in a flash, like a firework. Others take years, like iron rusting. The rate of a reaction tells us how fast it goes.
This matters in real life. Factories want products made quickly. Food companies want food that does not go off too fast. Understanding rate helps us control both.
By the end you will be able to calculate a mean rate and read a rate graph. You will explain what changes the rate using collisions between particles. You will also say what a catalyst does.
The rate of a reaction tells you how fast reactants are used up or products are made. It works like speed in a car, which is distance divided by time.
Mean rate = amount of reactant used or product formed ÷ time taken. Mean means average over the whole time.
Say 48 cm³ of gas is made in 60 seconds. The mean rate is 48 ÷ 60 = 0.80 cm³/s. The unit depends on what you measure: g/s for mass, cm³/s for gas volume and mol/s for amount.
To follow a reaction that makes a gas, collect the gas in a syringe. Record the volume at regular times. Or stand the flask on a balance and record the mass lost. For a cloudy mixture, time how long a cross under the flask takes to disappear.
Plot the results as product against time. The line is steepest at the start, when the reactants are at their highest concentration. It flattens as they are used up. A flat line means the reaction has finished.
The steeper curve is the faster reaction. If both end at the same height, the same amount of product was made.
Particles can only react if they collide. The collision also needs enough energy. A gentle bump does nothing, but a hard crash does.
The minimum energy needed is called the activation energy. A collision with at least this much energy is a successful collision.
So the rate depends on two things. One is how often the particles collide, which is the frequency of collisions. The other is how many of the collisions have enough energy. Anything that raises either one makes the reaction faster.
A higher concentration means more particles in the same volume. Think of a busy corridor between lessons. The more people packed in, the more often you bump into someone. The particles collide more often, so the rate goes up.
Pressure works the same way for gases. A higher pressure squashes the gas particles closer together. So they collide more often.
Breaking a solid into smaller pieces increases its surface area. More particles are exposed to collisions. Powdered marble reacts faster than marble chips of the same mass.
At a higher temperature the particles have more energy and move faster. They collide more often. A bigger fraction of the collisions also have enough energy to react. The second reason matters most.
A catalyst speeds up a reaction without being used up. It gives the reaction a different route with a lower activation energy. So more collisions are successful. The amount of product does not change.
Manganese dioxide is a catalyst for hydrogen peroxide breaking down: 2H₂O₂ → 2H₂O + O₂. Enzymes are biological catalysts.
Higher tier. On a curve the rate keeps changing. To find the rate at one moment, draw a tangent. A tangent is a straight line that just touches the curve at that point.
The rate is the gradient of the tangent. Gradient = change in y ÷ change in x. Think of how steep a hill is: how far it rises for each step along.
Say a tangent at 20 s passes through (0, 10) and (40, 40). The gradient is (40 − 10) ÷ (40 − 0) = 30 ÷ 40 = 0.75 cm³/s. Choose two points far apart to be more accurate.
Mean rate = amount ÷ time. A steeper line on a graph means a faster reaction, and a flat line means it has finished.
A reaction needs collisions with at least the activation energy. More frequent or more energetic collisions give a faster rate.
Concentration, pressure, surface area and temperature all change the rate. A catalyst lowers the activation energy and is not used up.
The rate at one moment is the gradient of a tangent.
45 cm³ of gas is collected in 30 seconds. Calculate the mean rate of reaction.
Answer: 1.5 cm³/s
Explain why powdered calcium carbonate reacts faster with hydrochloric acid than marble chips of the same mass.
Answer: Powder has a larger surface area, so more collisions happen each second.
The volume of gas collected is 40 cm³ at 20 s and 56 cm³ at 40 s. Calculate the mean rate between 20 s and 40 s and explain why it is lower than the mean rate over the first 20 s.
Answer: 0.80 cm³/s, which is lower because the reactants are being used up
Rates of reaction and catalysts practice questions for GCSE Combined Science Chemistry. Try each one, then open the answer. Easy questions come first.
A reaction makes 60 cm³ of gas in 30 seconds. Calculate the mean rate in cm³/s.
Answer: 2.0 cm³/s
What does a flat line at the end of a graph of product against time show?
Answer: the reaction has finished
Name the term for the minimum energy that particles need to react.
Answer: activation energy
Does a catalyst get used up in a reaction? Answer yes or no.
Answer: no
Does increasing the temperature make a reaction faster or slower?
Answer: faster
Which has the larger surface area, a lump of marble or the same mass of marble powder?
Answer: the powder
A balance shows a mass loss of 0.90 g in 180 s. Calculate the mean rate in g/s.
Answer: 0.0050 g/s
The volume of gas at 10 s is 24 cm³ and at 30 s is 50 cm³. Calculate the mean rate between 10 s and 30 s.
Answer: 1.3 cm³/s
Longer questions in the style of a written paper, with the marking points. These are Yumio practice questions.
A student reacts magnesium with dilute hydrochloric acid and measures the volume of hydrogen made. 36 cm³ of gas is collected in the first 20 seconds. (a) Calculate the mean rate of reaction in cm³/s over the first 20 s. [2 marks] (b) The experiment is repeated with the same mass of magnesium powder. Explain why the reaction is faster. [2 marks]
Answer: (a) 1.8 cm³/s (b) The powder has a larger surface area, so there are more collisions each second.
Explain, in terms of particles, how the rate of reaction between zinc and hydrochloric acid changes when (i) the acid is made more concentrated, and (ii) the temperature is increased. [6 marks]
Answer: (i) A more concentrated acid has more acid particles in the same volume, so the particles collide with zinc more often. There are more collisions per second, so a faster rate. (ii) A higher temperature gives the particles more energy, so they move faster and collide more often. More importantly, a greater proportion of the collisions have energy equal to or greater than the activation energy, so more collisions are successful and the rate is faster.
Saying a catalyst is used up
A catalyst is not used up. It lowers the activation energy.
Saying a higher temperature lowers the activation energy
The activation energy stays the same. More particles have enough energy.
Saying more concentrated means the particles move faster
Concentration changes how close the particles are, not their speed.
Dividing time by amount to find the rate
Rate = amount ÷ time.
Rates of reaction and catalysts is on both Foundation and Higher papers for GCSE Combined Science Chemistry, aimed at around grade 5.
Saying a catalyst is used up. A catalyst is not used up. It lowers the activation energy.
Open Show answer under each question on this page. The questions are Yumio practice questions, not past exam questions.
See how the papers are set: AQA Combined Science papers and Edexcel Combined Science papers.
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