Quick Simpler Start
In one sentence
You learn what makes a reaction faster or slower, how catalysts work, and how to read reaction graphs in a sensible way.
Exam habit
When you revise this chapter, ask yourself: "Can I explain this to a friend in one easy paragraph?" If yes, you are in a good place.
Interactive Simulator
This small tool gives you a quick visual check before you move deeper into the chapter notes.
Simulator
Rate Explorer
Choose a change and the tool tells you whether the reaction gets faster or slower.
Choose a change to see the effect on rate.
Chapter 6: Rates of Reaction
1. Objectives
- Define rate of reaction and explain how it can be measured.
- Describe the effect of particle size, concentration, temperature, catalyst and nature of reactants on reaction rate.
- Use collision theory to explain changes in rate.
- Explain activation energy using energy profile diagrams.
- Distinguish homogeneous and heterogeneous catalysis.
- Describe autocatalysis and recognise its graph shape.
- Explain surface adsorption and intermediate compound formation as catalytic mechanisms.
- Describe the purpose and chemistry of catalytic converters.
2. Key Definitions
| Term | Exam-ready meaning |
|---|---|
| Rate of reaction | The change in concentration of a reactant or product per unit time. |
| Collision theory | Reactions occur when particles collide with enough energy and correct orientation. |
| Activation energy | The minimum energy that reacting particles must have for a successful collision. |
| Catalyst | A substance that changes the rate of a reaction without being used up in the overall reaction. |
| Heterogeneous catalyst | A catalyst in a different physical phase from the reactants. |
| Homogeneous catalyst | A catalyst in the same physical phase as the reactants. |
| Autocatalysis | A reaction in which one of the products acts as a catalyst for the reaction. |
3. Visual Learning Zone
A. Collision theory: what makes a collision successful?
For a reaction to occur, particles must collide with enough energy and suitable orientation.
B. Rate graphs: product formed and reactant used up
At the start, rate is usually fastest because reactant concentration is highest. The graph levels off when the reaction is complete.
C. Energy profile: catalyst lowers activation energy
A catalyst speeds up the reaction by providing an alternative route with lower activation energy. It does not change the overall energy difference.
D. Factors affecting rate
All rate factors are explained using collision theory: they increase successful collisions per second.
E. Heterogeneous catalysis: surface adsorption
In heterogeneous catalysis, reactants adsorb onto a catalyst surface, react, and then leave the surface.
4. Core Notes
6.1 Reaction kinetics
The rate of a reaction tells us how quickly reactants are used up or products are formed. It can be measured by following changes such as mass loss, gas volume produced, precipitate formation, colour change, pH change or conductivity change.
Rate = change in amount or concentration ÷ time
Average rate = Δquantity / Δtime
Average rate = Δquantity / Δtime
| Factor | Effect on rate | Collision theory explanation |
|---|---|---|
| Smaller particle size | Faster reaction | Larger surface area means more collisions at the surface per second. |
| Higher concentration | Faster reaction | More particles in the same volume, so collisions happen more often. |
| Higher temperature | Faster reaction | Particles move faster and more have energy greater than activation energy. |
| Catalyst | Faster reaction | Provides an alternative pathway with lower activation energy. |
| Nature of reactants | Varies | Some bonds are easier to break or particles react more readily. |
6.2 Collision theory and activation energy
Not every collision causes a reaction. A collision is successful only if particles collide with enough energy and the correct orientation. The minimum required energy is called activation energy.
Temperature has a very strong effect because a small rise in temperature can greatly increase the number of particles with energy above the activation energy.
6.3 Catalysis
Catalysts change the rate of reaction without being consumed. They are not used up in the overall reaction, although they may form temporary intermediate species during the mechanism.
| Type | Meaning | Example idea |
|---|---|---|
| Homogeneous catalysis | Catalyst and reactants are in the same phase | All in aqueous solution or all gases |
| Heterogeneous catalysis | Catalyst and reactants are in different phases | Gases reacting on a solid metal surface |
| Autocatalysis | A product of the reaction acts as a catalyst | Rate starts slow, then speeds up, then slows as reactants are used up |
HL: Mechanisms of catalysis
- Surface adsorption: reactants attach to the surface of a solid catalyst, bonds weaken and products form.
- Intermediate compound formation: catalyst temporarily reacts to form an intermediate, then is regenerated.
- Autocatalysis graph: initially slow, then steep, then levels off.
Catalytic converters
Catalytic converters reduce harmful gases from car exhausts. They use solid catalysts such as platinum, palladium and rhodium to convert carbon monoxide, unburned hydrocarbons and nitrogen oxides into less harmful gases.
2CO + O₂ → 2CO₂
2NO → N₂ + O₂
Hydrocarbons + oxygen → CO₂ + H₂O
2NO → N₂ + O₂
Hydrocarbons + oxygen → CO₂ + H₂O
5. Worked Examples
Example 1: Calculating average rate
In an experiment, 60 cm³ of gas is produced in 120 seconds. Calculate the average rate of gas production.
Step 1: Rate = volume ÷ time
Step 2: Rate = 60 ÷ 120
Answer: 0.50 cm³ s⁻¹
In an experiment, 60 cm³ of gas is produced in 120 seconds. Calculate the average rate of gas production.
Step 1: Rate = volume ÷ time
Step 2: Rate = 60 ÷ 120
Answer: 0.50 cm³ s⁻¹
Example 2: Explaining the effect of temperature
Question: Explain why increasing temperature increases the rate of reaction. [3]
Answer: Particles gain kinetic energy [1]. They move faster and collide more often [1]. A greater fraction of particles have energy equal to or greater than the activation energy [1].
Question: Explain why increasing temperature increases the rate of reaction. [3]
Answer: Particles gain kinetic energy [1]. They move faster and collide more often [1]. A greater fraction of particles have energy equal to or greater than the activation energy [1].
Example 3: Catalyst explanation
Question: Explain how a catalyst increases reaction rate. [3]
Answer: It provides an alternative reaction pathway [1] with lower activation energy [1], so more particles have enough energy for successful collisions per second [1].
Question: Explain how a catalyst increases reaction rate. [3]
Answer: It provides an alternative reaction pathway [1] with lower activation energy [1], so more particles have enough energy for successful collisions per second [1].
6. Practical Skills
Experiment: Reaction of marble chips with hydrochloric acid
Safety: wear eye protection; handle acids carefully.
- Place marble chips in a conical flask.
- Add hydrochloric acid and quickly attach a gas syringe or place the flask on a balance.
- Measure volume of carbon dioxide produced over time, or mass lost over time.
- Repeat using different chip sizes, acid concentrations or temperatures.
- Plot gas volume against time or mass against time.
Safety: wear eye protection; handle acids carefully.
F. Measuring gas production with a gas syringe
The slope of a gas-volume/time graph gives the reaction rate. A steeper slope means a faster reaction.
7. Examiner Tips
- When asked about temperature, mention both more frequent collisions and more particles exceeding activation energy.
- When asked about concentration, focus on more particles per unit volume.
- When asked about particle size, focus on surface area.
- When asked about catalysts, always say alternative pathway with lower activation energy.
- For graphs, the gradient/slope represents reaction rate.
8. Common Mistakes
- Saying a catalyst gives particles more energy. It does not; it lowers activation energy.
- Writing that catalysts are used up. They are regenerated and not consumed overall.
- Confusing surface area with particle size: smaller particles have larger total surface area.
- Forgetting that rate decreases during many reactions because reactants are being used up.
- Saying every collision is successful. Only collisions with enough energy and correct orientation are successful.
9. Examiner Traps
If a question says “powder instead of lumps,” the answer is not “more particles.” The correct exam phrase is: greater surface area exposed to the acid, so more frequent successful collisions occur.
If a question says “catalyst remains unchanged,” that means unchanged at the end of the reaction, not necessarily unchanged during every step.
If a question says “catalyst remains unchanged,” that means unchanged at the end of the reaction, not necessarily unchanged during every step.
10. Exam Practice Questions
- Define rate of reaction. [2]
- Explain why powdered calcium carbonate reacts faster with hydrochloric acid than large marble chips. [3]
- Explain why increasing the concentration of acid increases the rate of reaction. [3]
- Explain why increasing temperature usually increases rate more strongly than increasing concentration. [4]
- Describe how a catalyst increases the rate of reaction. [3]
- Draw and label an energy profile diagram showing catalysed and uncatalysed pathways. [5]
- Distinguish between homogeneous and heterogeneous catalysis. [4]
- HL: Explain surface adsorption in heterogeneous catalysis. [4]
- HL: Explain what is meant by autocatalysis and describe the shape of its rate graph. [4]
- State one function of a catalytic converter in a car exhaust system. [2]
11. MCQs with Explanations
| Question | Answer & explanation |
|---|---|
| 1. Which change increases rate by increasing surface area? A Higher temperature B Powdering a solid C Adding water D Lowering concentration | B. Powdering a solid creates smaller particles and a larger exposed surface area. |
| 2. A catalyst increases rate because it: A increases product energy B is used up C lowers activation energy D increases temperature | C. A catalyst provides an alternative pathway with lower activation energy. |
| 3. The gradient of a product-volume/time graph represents: A yield B rate C activation energy D concentration | B. Gradient shows change in product volume per unit time. |
| 4. Which statement is true about successful collisions? A All collisions are successful B Particles need enough energy only C Correct orientation only is enough D Enough energy and correct orientation are needed | D. Both enough energy and correct orientation are required. |
| 5. In a catalytic converter, catalysts are usually: A solid metals B liquids C gases D salts dissolved in water | A. Platinum, palladium and rhodium are solid metal catalysts. |
12. HL Extension: Autocatalysis and Catalytic Mechanisms
Autocatalysis: A reaction product acts as a catalyst. The reaction may begin slowly because little catalyst is present. As more product forms, the rate increases. Later, the rate slows as reactants are used up.
Surface adsorption: reactants attach to the catalyst surface. This weakens bonds and brings particles close together in the correct orientation.
Intermediate formation: catalyst forms a temporary intermediate compound and is regenerated in a later step.
Surface adsorption: reactants attach to the catalyst surface. This weakens bonds and brings particles close together in the correct orientation.
Intermediate formation: catalyst forms a temporary intermediate compound and is regenerated in a later step.
13. Last-Minute Revision Sheet
- Rate = change in amount/concentration ÷ time.
- Successful collisions need enough energy and correct orientation.
- Activation energy = minimum energy needed for reaction.
- Higher concentration = more particles per volume = more collisions.
- Higher temperature = faster particles + more above activation energy.
- Smaller particles = greater surface area.
- Catalyst = alternative pathway with lower activation energy.
- Heterogeneous catalyst = different phase; homogeneous catalyst = same phase.
- Catalytic converters reduce harmful exhaust gases.
14. Self-Assessment Checklist
- I can define rate of reaction.
- I can explain rate changes using collision theory.
- I can identify factors affecting reaction rate.
- I can explain activation energy from an energy profile diagram.
- I can describe how catalysts work.
- I can compare homogeneous and heterogeneous catalysis.
- I can explain catalytic converters.
- I can describe autocatalysis at Higher Level.
15. Answers / Mark Scheme
Q1 [2] change in concentration/amount of reactant or product [1] per unit time [1].
Q2 [3] powder has smaller particles [1]; larger surface area [1]; more frequent successful collisions with acid [1].
Q3 [3] more acid particles per unit volume [1]; particles are closer together [1]; more frequent successful collisions [1].
Q4 [4] particles have more kinetic energy [1]; move faster and collide more often [1]; greater fraction exceed activation energy [1]; more successful collisions per second [1].
Q5 [3] alternative pathway [1]; lower activation energy [1]; more successful collisions per second [1].
Q6 [5] labelled axes energy/progress [1]; reactants/products shown [1]; uncatalysed higher peak [1]; catalysed lower peak [1]; activation energy labelled [1].
Q7 [4] homogeneous same phase [1] with example [1]; heterogeneous different phase [1] with example [1].
Q8 HL [4] reactants adsorb onto surface [1]; bonds weaken [1]; reactants held in suitable orientation [1]; products desorb and catalyst remains [1].
Q9 HL [4] product acts as catalyst [1]; rate begins slow [1]; rate increases as catalyst product forms [1]; later levels/slows as reactants are used up [1].
Q10 [2] converts carbon monoxide/nitrogen oxides/unburned hydrocarbons [1] into less harmful gases such as carbon dioxide, nitrogen and water [1].
Q2 [3] powder has smaller particles [1]; larger surface area [1]; more frequent successful collisions with acid [1].
Q3 [3] more acid particles per unit volume [1]; particles are closer together [1]; more frequent successful collisions [1].
Q4 [4] particles have more kinetic energy [1]; move faster and collide more often [1]; greater fraction exceed activation energy [1]; more successful collisions per second [1].
Q5 [3] alternative pathway [1]; lower activation energy [1]; more successful collisions per second [1].
Q6 [5] labelled axes energy/progress [1]; reactants/products shown [1]; uncatalysed higher peak [1]; catalysed lower peak [1]; activation energy labelled [1].
Q7 [4] homogeneous same phase [1] with example [1]; heterogeneous different phase [1] with example [1].
Q8 HL [4] reactants adsorb onto surface [1]; bonds weaken [1]; reactants held in suitable orientation [1]; products desorb and catalyst remains [1].
Q9 HL [4] product acts as catalyst [1]; rate begins slow [1]; rate increases as catalyst product forms [1]; later levels/slows as reactants are used up [1].
Q10 [2] converts carbon monoxide/nitrogen oxides/unburned hydrocarbons [1] into less harmful gases such as carbon dioxide, nitrogen and water [1].
16. HL Syllabus Patch: Missing Core Theory
Average Rate and Instantaneous Rate
Average rate is the overall change divided by the total time taken. Instantaneous rate is the rate at one exact moment and is found from the slope of a tangent to the graph.
average rate = change / time
HL Exam Tip: If the question says “instantaneous rate”, draw a tangent to the curve at the chosen point and calculate its gradient. Do not use the slope of the whole curve.
| Type of rate | How to find it | When used |
|---|---|---|
| Average rate | Use two points over a time interval | Overall rate during a section of the reaction |
| Instantaneous rate | Use gradient of tangent at a point | Exact rate at one moment |
Nature of Reactants
The rate also depends on the chemical nature of the reacting substances. Ionic reactions in solution are often fast because ions are already free to react. Covalent reactions may be slower because bonds must first be broken.
Example: Silver nitrate reacting with sodium chloride is very fast in solution, while the reaction between marble chips and acid is slower because it occurs at a solid surface.
Dust Explosions
Powdered solids can react far more rapidly than lumps of the same substance because the surface area is much greater. Fine flour, coal dust or metal dust mixed with air can burn so quickly that an explosion occurs.
Examiner Trap: A dust explosion is not caused by “more heat” in the powder. It is caused mainly by the huge increase in surface area, which gives many more successful collisions per second.
17. Mandatory Experiments and Required Demonstrations
Mandatory Experiment 6.1: Rate of decomposition of hydrogen peroxide using manganese dioxide
Aim: To investigate the effect of a catalyst on the rate of decomposition of hydrogen peroxide.
| Item | Details |
|---|---|
| Reaction | 2H2O2(aq) → 2H2O(l) + O2(g) |
| Catalyst | Manganese dioxide, MnO2 |
| Measurement | Collect oxygen over time using a gas syringe or inverted measuring cylinder |
| Observation | Rapid bubbling; oxygen relights a glowing splint |
Key idea: MnO2 speeds up the reaction but is not used up overall.
Mandatory Experiment 6.2: Effect of concentration using sodium thiosulfate and hydrochloric acid
Aim: To study how concentration affects reaction rate using the disappearing cross method.
| Reaction | Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + SO2(g) + S(s) + H2O(l) |
|---|
| What changes? | What stays the same? | What is timed? |
|---|---|---|
| Concentration of sodium thiosulfate | Volume of acid, total volume, temperature | Time for the cross to disappear |
Conclusion: A higher concentration gives a faster reaction because there are more particles per unit volume and therefore more frequent successful collisions.
Required demonstration: oxidation of methanol using hot platinum or nichrome wire
Hot platinum or nichrome wire catalyses the oxidation of methanol by oxygen in air. The wire remains glowing because the oxidation is exothermic.
Why this matters: This demonstration shows catalytic oxidation and how a catalyst can provide an easier pathway for reaction.
Required demonstration: oxidation of potassium sodium tartrate by hydrogen peroxide catalysed by cobalt(II) salts
This reaction is used to show homogeneous catalysis. The cobalt(II) ions catalyse the reaction while being regenerated during the process.
HL Link: This is a good example of catalysis in solution and is useful when comparing homogeneous and heterogeneous catalysis.
18. Enzymes and Catalyst Poisons
Enzymes as Biological Catalysts
Enzymes are protein catalysts found in living systems. They work best under suitable temperature and pH conditions.
| Enzyme | Substrate | Action |
|---|---|---|
| Catalase | Hydrogen peroxide | Breaks it down into water and oxygen |
| Amylase | Starch | Breaks starch into smaller sugar molecules |
Common Mistake: Enzymes are not “used up” by reactions, but they can be denatured by unsuitable temperature or pH.
Catalyst Poisons
A catalyst poison is a substance that reduces or destroys the activity of a catalyst.
| Catalyst | Poison | Effect |
|---|---|---|
| Car catalyst metals | Lead compounds | Block active sites and reduce efficiency |
| Industrial catalysts | Sulfur impurities or other contaminants | Lower reaction rate by preventing adsorption |
Examiner Trap: A catalyst poison does not “reverse the catalyst”. It blocks the catalyst surface or interferes with the catalytic cycle.
19. Quick Final Checklist for HL Chapter 6
- I can distinguish average rate from instantaneous rate.
- I can calculate an instantaneous rate from the tangent to a graph.
- I can explain how concentration, particle size, temperature, nature of reactants and catalysts affect rate using collision theory.
- I can describe the two mandatory experiments for this chapter.
- I can explain dust explosions using surface area.
- I can name two enzymes and describe them as catalysts.
- I can explain what is meant by a catalyst poison.