(I once asked a class, "Why does a fire burn faster when you blow on it?" Most students said, "Because you are giving it more oxygen." That is part of it. But the deeper answer is about collisions — the particles of oxygen are colliding with the fuel more often, and each collision has more energy. That is what this chapter is about. It is not just knowing that things happen faster; it is knowing why. And once you know why, you can control reactions — speed them up when you want them to go, and slow them down when you do not.) Think about the reactions in your daily life. A match burns in seconds. Food cooks in minutes. Milk sours in hours. Iron rusts over years. A diamond takes millions of years to form. All of these are chemical reactions — but their rates are wildly different. Why? What makes one reaction fast and another slow? This is the question that rates of reaction answers. The rate of a reaction is how quickly reactants are turned into products. It depends on several factors: concentration, temperature, surface area, pressure, and the presence of a catalyst. Understanding these factors allows chemists to control reactions — speeding up the ones we want (like the reactions in a car engine or in our bodies) and slowing down the ones we do not (like food spoiling or metal rusting). In this chapter, we are going to look at the factors that affect reaction rates, the collision theory that explains them, and the experimental methods used to measure them. By the end, you will understand why some reactions are fast and some are slow — and how to control them. Let us begin.
The rate of a reaction is the change in concentration of a reactant or product per unit time. In simpler terms, it is how fast a reaction happens.
Ways to measure rate:
Units of rate: mol/dm³/s or g/s or cm³/s, depending on what is being measured.
(Think of reaction rate like driving a car. The speedometer tells you how fast you are going at any moment. In a reaction, the "speed" changes over time — it is fastest at the start when there are plenty of reactants, and it slows down as the reactants are used up. This is why we often measure the average rate over the first few seconds of a reaction.)
The collision theory explains how reactions happen at the particle level. For a reaction to occur:
If these conditions are met, the collision is successful and products are formed. If not, the particles simply bounce off each other.
Activation energy (Ea): The minimum energy required for a collision to result in a reaction. It is the energy barrier that reactants must overcome.
(Imagine trying to push a boulder over a hill. You need enough energy to get it to the top before it will roll down the other side. The activation energy is the height of that hill. Reactions with low activation energy happen easily; reactions with high activation energy need a lot of energy to get started. This is why a spark can start a fire — it provides the activation energy.)
Several factors affect the rate of a reaction. Each can be explained using collision theory.
Effect: Increasing the concentration of a reactant increases the rate of reaction.
Explanation: Higher concentration means more particles in a given volume. This leads to more frequent collisions, increasing the rate.
Example: Magnesium reacts faster with concentrated hydrochloric acid than with dilute hydrochloric acid.
Effect: Increasing the temperature increases the rate of reaction.
Explanation: Higher temperature means particles have more kinetic energy. They move faster, collide more frequently, and more importantly, more collisions have enough energy to overcome the activation energy. The effect is dramatic — a 10°C rise can double or triple the rate.
Example: Food cooks faster at higher temperatures. A glow stick glows brighter in warm water.
Effect: Increasing the surface area of a solid reactant increases the rate of reaction.
Explanation: A larger surface area means more particles are exposed to collisions with the other reactant. This increases the frequency of collisions.
Example: Powdered calcium carbonate reacts faster with acid than a lump of calcium carbonate. This is why we chew food — it increases the surface area for digestion.
Effect: Increasing the pressure of gaseous reactants increases the rate of reaction.
Explanation: Higher pressure means gas particles are squeezed closer together, increasing the frequency of collisions.
Example: Reactions in compressed gases occur faster than in gases at atmospheric pressure.
Effect: A catalyst increases the rate of reaction without being used up.
Explanation: A catalyst provides an alternative pathway with a lower activation energy. More collisions have enough energy to react, so the rate increases.
Example: Manganese(IV) oxide catalyses the decomposition of hydrogen peroxide. Enzymes in your body catalyse biochemical reactions.
| Factor | Effect on Rate | Explanation (Collision Theory) |
|---|---|---|
| Concentration | Increases | More particles → more frequent collisions |
| Temperature | Increases | More kinetic energy → more frequent and more energetic collisions |
| Surface area | Increases | More exposed particles → more frequent collisions |
| Pressure (gases) | Increases | Particles closer together → more frequent collisions |
| Catalyst | Increases | Lower activation energy → more collisions have enough energy |
Chemists use several methods to measure reaction rates. Each method measures a change in a property of the reaction mixture over time.
Suitable for reactions that produce a gas. The gas is collected in a gas syringe or measuring cylinder, and the volume is recorded at regular intervals.
Example: Magnesium + hydrochloric acid → hydrogen gas.
Suitable for reactions that produce a gas. The reaction flask is placed on a balance, and the mass is recorded at regular intervals. The loss in mass is due to the escaping gas.
Example: Calcium carbonate + hydrochloric acid → carbon dioxide.
Suitable for reactions involving a colour change. A colorimeter is used to measure the absorbance of light at regular intervals.
Example: The reaction between iodine and propanone.
Suitable for reactions that produce a precipitate. A cross is drawn on a piece of paper under the reaction vessel, and the time taken for the precipitate to obscure the cross is recorded.
Example: Sodium thiosulphate + hydrochloric acid → sulphur precipitate.
A rate graph plots the amount of reactant or product against time. The slope of the graph gives the rate of reaction.
| Feature | Explanation |
|---|---|
| Steep slope at start | Rate is fastest at the beginning because reactant concentration is highest. |
| Slope decreases over time | Rate slows down as reactants are used up. |
| Flat line at the end | Reaction has stopped — all reactants have been used up (or equilibrium has been reached). |
(A rate graph tells a story. The steep part at the beginning is the excitement of the reaction — everything is fresh and collisions are frequent. The curve gradually flattens as the reactants run out. The flat line at the end is the finish — either the reactants are gone, or the reaction has reached equilibrium. Learning to read rate graphs is like learning to read the pulse of a chemical reaction.)
A catalyst is a substance that increases the rate of a chemical reaction without being used up. It provides an alternative reaction pathway with a lower activation energy.
| Type of Catalyst | Description | Example |
|---|---|---|
| Homogeneous | Same physical state as the reactants. | Acid catalysts in esterification (all liquids). |
| Heterogeneous | Different physical state from the reactants. | Iron in the Haber process (solid catalyst, gaseous reactants). |
| Enzymes | Biological catalysts — proteins that speed up reactions in living things. | Amylase breaks down starch; catalase breaks down hydrogen peroxide. |
Important: A catalyst does not affect the position of equilibrium — it only speeds up the rate at which equilibrium is reached.
Catalyst poisoning: Some substances (called poisons) can bind to a catalyst and prevent it from working. This is why industrial catalysts must be protected from impurities.
Understanding and controlling reaction rates is essential in many areas of life and industry:
| Area | Importance |
|---|---|
| Industry | Chemical engineers optimise temperature, pressure, and catalysts to maximise the rate and yield of industrial reactions (e.g., Haber process, contact process). |
| Medicine | The rate at which drugs are absorbed and metabolised affects their effectiveness. Enzymes in the body control the rates of biochemical reactions. |
| Food preservation | Refrigeration slows down the rate of food spoilage. Canning and drying also slow down reactions. |
| Everyday life | Pressure cookers speed up cooking by increasing temperature and pressure. Chewing food increases surface area for digestion. |
| Environment | Catalytic converters in cars speed up the conversion of harmful gases into less harmful ones. |
Some factors that students sometimes think affect rate do not:
Question: Explain, using collision theory, why increasing the temperature increases the rate of a reaction.
Solution: Increasing the temperature gives the reactant particles more kinetic energy. They move faster and collide more frequently. More importantly, a greater proportion of the collisions have enough energy to overcome the activation energy. This means more successful collisions per unit time, so the rate of reaction increases.
Question: Explain why powdered calcium carbonate reacts faster with hydrochloric acid than the same mass of calcium carbonate in large lumps.
Solution: Powdered calcium carbonate has a much larger surface area than large lumps. This means more of the calcium carbonate particles are exposed to the hydrochloric acid. The frequency of collisions between the acid particles and the calcium carbonate particles increases, so the rate of reaction is faster.
Question: Explain how a catalyst increases the rate of a reaction.
Solution: A catalyst provides an alternative reaction pathway with a lower activation energy. This means that a greater proportion of collisions have enough energy to react. More successful collisions occur per unit time, so the rate of reaction increases. The catalyst is not used up in the reaction.
Question: A graph shows the volume of gas produced against time for a reaction. The curve is steep at first, then flattens. Explain the shape of the graph.
Solution: The curve is steep at first because the reaction is fastest at the beginning — the concentration of reactants is highest, so collisions are most frequent. As the reaction proceeds, the reactants are used up, so the concentration decreases and the rate slows down. The curve flattens as the reaction approaches completion — the reactants are nearly gone, and the rate approaches zero. The flat line at the end indicates that the reaction has stopped.
1. Define the rate of a reaction.
The rate of a reaction is the change in concentration of a reactant or product per unit time.
2. State the three conditions for a successful collision according to collision theory.
1. Particles must collide. 2. The collision must have enough energy (activation energy). 3. The particles must collide with the correct orientation.
3. Explain why increasing concentration increases the rate of reaction.
Increasing concentration means more particles in a given volume. This leads to more frequent collisions, increasing the rate of reaction.
4. Explain why increasing temperature increases the rate of reaction.
Higher temperature gives particles more kinetic energy. They collide more frequently, and more collisions have enough energy to overcome the activation energy. Both effects increase the rate.
5. Explain why powdered reactants react faster than large lumps.
Powdered reactants have a larger surface area, so more particles are exposed to collisions with the other reactant. This increases the frequency of collisions and the rate of reaction.
6. What is a catalyst? How does it increase the rate of a reaction?
A catalyst is a substance that increases the rate of a reaction without being used up. It provides an alternative pathway with a lower activation energy, so more collisions have enough energy to react.
7. Why does a rate graph flatten out over time?
The rate graph flattens because the reactants are being used up. As their concentration decreases, collisions become less frequent, and the rate slows down. Eventually, the reaction stops when the reactants are exhausted (or equilibrium is reached).
Question 1 [10 marks]
(a) Define the rate of a reaction and state three ways it can be measured. [4]
(b) Describe the collision theory and explain what is meant by activation energy. [4]
(c) Explain, using collision theory, why a reaction slows down over time. [2]
(a) Rate of reaction is the change in concentration of a reactant or product per unit time. Ways to measure it: measuring volume of gas produced, measuring loss in mass, measuring change in colour/absorbance, measuring time for precipitate to form. (4 marks: 2 for definition, 2 for any two methods.)
(b) Collision theory states that for a reaction to occur, particles must collide with enough energy and the correct orientation. Activation energy is the minimum energy required for a collision to result in a reaction. (4 marks: 2 for collision theory, 2 for activation energy.)
(c) As the reaction proceeds, the reactants are used up, so their concentration decreases. This means there are fewer particles to collide, so the frequency of successful collisions decreases, and the rate slows down. (2 marks: 1 for "reactants used up", 1 for "fewer collisions".)
Question 2 [10 marks]
(a) Describe an experiment to investigate the effect of temperature on the rate of reaction between sodium thiosulphate and hydrochloric acid. [5]
(b) Explain, using collision theory, the effect of temperature on the rate of reaction. [3]
(c) Sketch a rate graph for this reaction at two different temperatures and explain the difference. [2]
(a) Place a measured volume of sodium thiosulphate solution in a conical flask on a piece of paper with a cross marked on it. Add a measured volume of hydrochloric acid and start the timer. Record the time taken for the cross to disappear (due to the sulphur precipitate). Repeat at different temperatures (e.g., 20°C, 30°C, 40°C). (5 marks: 1 for apparatus, 1 for cross method, 1 for timing, 1 for varying temperature, 1 for repeat.)
(b) Increasing the temperature gives particles more kinetic energy. They move faster and collide more frequently. More importantly, more collisions have enough energy to overcome the activation energy. Both effects increase the rate of reaction. (3 marks: 1 for more collisions, 1 for more energy, 1 for "more successful collisions".)
(c) At higher temperatures, the curve is steeper and the reaction finishes sooner. At lower temperatures, the curve is less steep and takes longer. This is because higher temperatures increase the rate of reaction. (2 marks: 1 for steeper curve at higher temperature, 1 for explanation.)
Question 3 [10 marks]
(a) Describe what is meant by a catalyst and explain how it works. [3]
(b) Describe two industrial uses of catalysts. [4]
(c) Explain why a catalyst does not affect the position of equilibrium. [3]
(a) A catalyst is a substance that increases the rate of a reaction without being used up. It provides an alternative reaction pathway with a lower activation energy, so more collisions have enough energy to react. (3 marks: 1 for definition, 1 for "lower activation energy", 1 for "more successful collisions".)
(b) Iron in the Haber process (making ammonia from nitrogen and hydrogen). Vanadium(V) oxide in the contact process (making sulphur trioxide for sulphuric acid). (4 marks: 2 for each use with catalyst and process.)
(c) A catalyst does not affect the position of equilibrium because it speeds up the forward and reverse reactions equally. It only helps the system reach equilibrium faster. The equilibrium position depends on the thermodynamics of the reaction, not on the presence of a catalyst. (3 marks: 1 for "speeds up both directions equally", 1 for "equilibrium depends on thermodynamics", 1 for "reaches equilibrium faster".)
Challenge 1: Explain why increasing the pressure of a gaseous reaction increases the rate, but adding an inert gas at constant volume does not.
Increasing the pressure of a gaseous reaction increases the rate because it squeezes the gas particles closer together. The concentration of the reactant particles increases, so they collide more frequently, and the rate increases. Adding an inert gas at constant volume, however, does not change the partial pressures of the reactant gases. The reactant particles are not squeezed closer together — the inert gas just occupies space. The frequency of collisions between reactant particles remains unchanged, so the rate is unchanged. The key difference is that increasing pressure increases the concentration of the reactants, while adding an inert gas at constant volume does not.
Challenge 2: Explain why a solid reactant in excess does not increase the rate of a reaction when more is added.
If a solid reactant is already in excess, adding more of it does not increase the rate of reaction because the rate is limited by the other reactant, not by the solid. The rate depends on the frequency of collisions between the solid and the other reactant, and this is determined by the surface area of the solid that is exposed, not by how much of the solid is present. Once the solid is in excess, the surface area exposed to the other reactant is already more than enough to react with all the available molecules of the other reactant. Adding more solid does not increase the surface area in contact with the other reactant, so the rate does not change. The other reactant is the limiting factor, and its concentration determines the rate.
Challenge 3: A student says, "A catalyst increases the yield of a reaction." Is this statement accurate? Explain your reasoning.
The statement is not accurate. A catalyst increases the rate of a reaction, not the yield. It provides an alternative pathway with lower activation energy, so the reaction reaches equilibrium faster. However, it does not change the position of equilibrium. The yield of a reaction depends on the thermodynamics of the reaction — the relative energies of reactants and products, and the conditions (temperature and pressure) — not on the presence of a catalyst. For a reaction that does not reach completion, a catalyst speeds up the approach to equilibrium but does not change the equilibrium yield. For a reaction that goes to completion, the yield is determined by the limiting reactant, not by the catalyst. So a catalyst increases the rate, not the yield.
"Rate of reaction is the pulse of chemistry — how fast, how slow, and why. Now you know how to read it."
