What you'll learn
This topic covers how chemists measure and control the speed of reactions and understand reversible processes. You'll explore factors affecting reaction rates, collision theory, catalysts, and dynamic equilibrium. These principles underpin industrial chemical processes like the Haber process and Contact process, which are essential for manufacturing fertilisers and acids globally.
Key terms and definitions
Rate of reaction — the speed at which reactants are converted into products, measured as the change in concentration (or amount) per unit time
Activation energy — the minimum energy that colliding particles must have for a reaction to occur
Catalyst — a substance that increases the rate of a reaction without being used up, by providing an alternative pathway with lower activation energy
Reversible reaction — a chemical reaction where the products can react to reform the original reactants
Dynamic equilibrium — the state in a reversible reaction where the forward and reverse reactions occur at the same rate, and concentrations of reactants and products remain constant
Le Chatelier's principle — if a system at equilibrium is subjected to a change in conditions, the position of equilibrium shifts to counteract that change
Yield — the amount of product obtained from a reaction, often expressed as a percentage of the theoretical maximum
Exothermic reaction — a reaction that transfers energy to the surroundings, usually as heat
Core concepts
Measuring rates of reaction
The rate of a chemical reaction can be measured by monitoring either:
- The decrease in concentration of a reactant over time
- The increase in concentration of a product over time
Common experimental methods include:
Measuring gas volume produced: When a gas is produced (e.g., CO₂ from calcium carbonate and hydrochloric acid), collect it in a gas syringe or measure displacement of water. Record volume at regular time intervals. Calculate rate using the formula:
Rate = volume of gas produced ÷ time
Measuring mass loss: Place the reaction vessel on a balance. If a gas escapes, the mass decreases. Record mass at intervals. More rapid mass loss indicates a faster reaction.
Measuring time for a precipitate to form: In reactions producing cloudy precipitates (e.g., sodium thiosulfate with acid), measure how long it takes for a mark beneath the flask to disappear. Shorter time means faster rate.
Monitoring colour change: Use a colorimeter to measure light absorption through coloured solutions. Useful for reactions involving transition metal ions.
The gradient of a concentration-time or volume-time graph gives the rate at any point. Steeper gradients indicate faster rates. The gradient is steepest at the start when concentration is highest.
Factors affecting reaction rates
Five key factors influence how fast reactions proceed:
Temperature: Increasing temperature raises the kinetic energy of particles. They move faster, collide more frequently, and more collisions exceed the activation energy. A 10°C rise typically doubles or triples reaction rate.
Concentration of solutions: Higher concentration means more particles in the same volume. This increases collision frequency, leading to more successful collisions per unit time. Rate is directly proportional to concentration for many reactions.
Pressure of gases: Increasing pressure compresses gas particles into a smaller volume. More particles per unit volume means more frequent collisions and faster reactions.
Surface area of solids: Larger surface area (smaller particle size) exposes more particles to collision. Powdered solids react much faster than lumps. Breaking a solid into smaller pieces increases the area available for reaction without changing the mass.
Catalysts: These provide an alternative reaction pathway with lower activation energy. More collisions now have sufficient energy to react, dramatically increasing rate without raising temperature.
Collision theory and activation energy
For a reaction to occur between particles, they must:
- Collide with each other
- Have energy equal to or greater than the activation energy
- Collide with the correct orientation (geometry matters for complex molecules)
Not all collisions lead to reactions. Only successful collisions — those meeting all three criteria — result in product formation.
Activation energy acts as an energy barrier. Particles with kinetic energy below this threshold bounce apart unchanged, even if they collide. Raising temperature or adding a catalyst increases the proportion of particles exceeding this barrier.
Energy profile diagrams show:
- Energy on the y-axis
- Reaction progress on the x-axis
- Reactants' energy level
- Products' energy level
- An activation energy "hill" between them
For exothermic reactions, products sit lower than reactants (energy released). For endothermic reactions, products sit higher (energy absorbed).
Catalysts reduce the activation energy, creating a lower "hill" on the energy profile. The same energy change occurs between reactants and products, but more particles can now overcome the reduced barrier.
Catalysts and their importance
Catalysts are crucial in industrial chemistry because they:
- Increase reaction rates at lower temperatures, saving energy costs
- Reduce energy consumption, lowering carbon emissions
- Improve economic efficiency of processes
- Remain chemically unchanged, so can be reused repeatedly
Different reactions require specific catalysts:
- Iron catalyses ammonia production in the Haber process
- Vanadium(V) oxide catalyses sulfuric acid manufacture in the Contact process
- Platinum and palladium in catalytic converters reduce vehicle emissions
- Enzymes (biological catalysts) are used in food processing and medicine manufacture
Catalysts work by:
- Adsorbing reactant molecules onto their surface
- Weakening bonds within reactant molecules
- Allowing products to form more easily
- Releasing products and freeing the catalyst surface for reuse
Reversible reactions and equilibrium
In a closed system (no substances can enter or leave), many reactions are reversible. The symbol ⇌ represents this:
A + B ⇌ C + D
Initially, only the forward reaction (A + B → C + D) occurs. As C and D form, the reverse reaction (C + D → A + B) begins. Eventually, both reactions occur at equal rates — this is dynamic equilibrium.
At equilibrium:
- Both forward and reverse reactions continue constantly
- Concentrations of all substances remain constant (not necessarily equal)
- The system appears static macroscopically but is active at the molecular level
The position of equilibrium describes whether reactants or products dominate at equilibrium:
- Equilibrium lies to the right: more products than reactants
- Equilibrium lies to the left: more reactants than products
Le Chatelier's principle and predicting equilibrium shifts
When conditions change, equilibrium shifts to oppose that change. This is Le Chatelier's principle.
Effect of concentration changes:
- Increase reactant concentration → equilibrium shifts right (more products form)
- Increase product concentration → equilibrium shifts left (more reactants form)
- Decrease concentration → equilibrium shifts to replace what was removed
Effect of temperature changes:
- For exothermic forward reactions (releasing heat):
- Increase temperature → equilibrium shifts left (favouring the reverse endothermic direction)
- Decrease temperature → equilibrium shifts right (favouring forward reaction)
- For endothermic forward reactions (absorbing heat):
- Increase temperature → equilibrium shifts right (favouring forward reaction)
- Decrease temperature → equilibrium shifts left
Effect of pressure changes (gases only):
- Increase pressure → equilibrium shifts toward the side with fewer gas molecules
- Decrease pressure → equilibrium shifts toward the side with more gas molecules
- If both sides have equal moles of gas, pressure changes have no effect
Catalysts and equilibrium: Catalysts increase the rate at which equilibrium is reached but do not change the position of equilibrium. Both forward and reverse rates increase equally.
Industrial applications: The Haber process
The Haber process manufactures ammonia (NH₃) for fertiliser production:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol (exothermic)
Conditions used:
- Temperature: 450°C
- Pressure: 200 atmospheres
- Catalyst: iron
Applying Le Chatelier's principle:
The forward reaction is exothermic and produces fewer gas molecules (4 moles → 2 moles).
Low temperature would shift equilibrium right, increasing yield. However, rates would be too slow. The compromise of 450°C gives reasonable yield at practical rates.
High pressure shifts equilibrium right (fewer molecules). 200 atmospheres is used because higher pressures require stronger, more expensive equipment. Safety concerns also limit pressure.
Iron catalyst speeds up the reaction without affecting yield. This allows acceptable rates at moderate temperature.
Economic considerations:
- Unreacted N₂ and H₂ are recycled to improve atom economy
- Energy costs must balance against yield
- Equipment costs increase significantly with pressure
Industrial applications: The Contact process
This process produces sulfuric acid (H₂SO₄), vital for manufacturing fertilisers, detergents, and paints:
2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = -197 kJ/mol (exothermic)
Conditions used:
- Temperature: 450°C
- Pressure: 2 atmospheres
- Catalyst: vanadium(V) oxide (V₂O₅)
Applying Le Chatelier's principle:
The forward reaction is exothermic and produces fewer molecules (3 moles → 2 moles).
Moderate temperature (450°C) balances yield and rate. Lower temperatures increase yield but slow the reaction unacceptably.
Low pressure (2 atmospheres) is sufficient because the yield is already high (~99%) at this pressure. Higher pressures would be unnecessary and expensive.
Vanadium(V) oxide catalyst provides economical rates at moderate temperature.
The SO₃ is then dissolved in concentrated H₂SO₄ (not water, which is too exothermic) to form oleum, which is diluted to produce sulfuric acid.
Worked examples
Example 1: A student investigates how concentration affects reaction rate using magnesium ribbon and hydrochloric acid. They measure the volume of hydrogen gas produced.
(a) State two variables that must be controlled. [2 marks]
Answer:
- Temperature (1 mark)
- Mass/length of magnesium ribbon (1 mark)
[Also acceptable: surface area of magnesium, volume of acid used]
(b) Explain why increasing acid concentration increases the rate of reaction. [3 marks]
Answer:
- More acid particles per unit volume / increased particle concentration (1 mark)
- More frequent collisions between particles (1 mark)
- More successful collisions per unit time (1 mark)
Example 2: The Haber process produces ammonia:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol
(a) Explain why increasing pressure increases the yield of ammonia. [3 marks]
Answer:
- Equilibrium shifts to oppose the change / in the direction with fewer molecules (1 mark)
- The forward reaction produces fewer moles of gas (4 moles → 2 moles) (1 mark)
- So equilibrium shifts right, producing more ammonia (1 mark)
(b) Explain why a temperature of 450°C is used rather than a lower temperature. [2 marks]
Answer:
- Lower temperature would increase yield (1 mark)
- But would make the rate too slow / 450°C is a compromise between rate and yield (1 mark)
Example 3: A student measures the rate of reaction between sodium thiosulfate solution and hydrochloric acid at different temperatures. They record the time for a cross beneath the flask to disappear.
| Temperature (°C) | Time (s) |
|---|---|
| 20 | 180 |
| 30 | 90 |
| 40 | 45 |
| 50 | 22 |
(a) Describe the relationship between temperature and time. [1 mark]
Answer: As temperature increases, time decreases / they are inversely proportional (1 mark)
(b) Explain this relationship using collision theory. [3 marks]
Answer:
- Higher temperature gives particles more kinetic energy (1 mark)
- Particles collide more frequently (1 mark)
- More collisions have energy greater than the activation energy / more successful collisions (1 mark)
Common mistakes and how to avoid them
Confusing "rate" with "yield": Rate is how fast a reaction occurs; yield is how much product forms. Catalysts increase rate but don't affect equilibrium yield. Temperature compromises balance both.
Saying particles "have more energy to react": Be specific. Say particles have "kinetic energy greater than the activation energy" or "more particles exceed the activation energy threshold."
Claiming equilibrium means concentrations are equal: At equilibrium, concentrations remain constant, not equal. The position of equilibrium determines the ratio of products to reactants.
Thinking catalysts shift equilibrium position: Catalysts speed up both forward and reverse reactions equally. They reduce time to reach equilibrium but don't change the final concentrations.
Forgetting to mention collision frequency AND energy: Both are essential. More frequent collisions alone don't guarantee faster reactions if energy remains below activation energy.
Vague explanations of Le Chatelier's principle: Always state which direction equilibrium shifts AND explain why (more/fewer molecules, exothermic/endothermic, counteracting the change).
Exam technique for "C5: Monitoring and Controlling Chemical Reactions"
"Explain" questions require reasons, not just descriptions. For rate questions, mention collision frequency AND successful collisions with sufficient energy. Link your explanation to collision theory explicitly.
Industrial process questions test understanding of compromises. Explain both the ideal conditions (from Le Chatelier) and why practical conditions differ (economics, safety, rate considerations). Typically worth 4-6 marks.
Graph questions may ask you to calculate rate from gradient or identify where rate is fastest. Rate = change in y ÷ change in x. Steepest gradient = fastest rate, usually at the start.
Command words matter: "State" = no explanation needed (1 mark each). "Describe" = say what happens. "Explain" = say why it happens using scientific principles. Match your answer length to marks available.
Quick revision summary
Rate of reaction depends on temperature, concentration, pressure, surface area, and catalysts. Collision theory explains that particles must collide with sufficient energy (activation energy) for reaction. Catalysts provide alternative pathways with lower activation energy. In reversible reactions, dynamic equilibrium occurs when forward and reverse rates equal. Le Chatelier's principle predicts equilibrium shifts: systems oppose changes in concentration, temperature, or pressure. Industrial processes like Haber (ammonia) and Contact (sulfuric acid) use compromise conditions balancing yield, rate, and economics.