What you'll learn
This revision guide covers C3: Chemical Reactions from the OCR GCSE Chemistry specification. You'll master energy changes in reactions, understand reaction rates and how to manipulate them, explore equilibria, and calculate quantities in chemical reactions. These concepts form the foundation for understanding how and why chemical processes occur in industry and everyday life.
Key terms and definitions
Exothermic reaction — a reaction that transfers energy to the surroundings, resulting in a temperature increase of the surroundings
Endothermic reaction — a reaction that takes in energy from the surroundings, resulting in a temperature decrease of the surroundings
Activation energy — the minimum energy that colliding particles must have for a reaction to occur
Catalyst — a substance that increases the rate of a chemical reaction without being used up in the process
Rate of reaction — the speed at which reactants are converted into products, measured as the change in concentration (or amount) per unit time
Dynamic equilibrium — in a reversible reaction, the state where the forward and reverse reactions occur at the same rate, so concentrations remain constant
Le Chatelier's principle — if a system at equilibrium is subjected to a change, the position of equilibrium shifts to counteract that change
Collision theory — the theory that reactions occur when particles collide with sufficient energy (activation energy) and correct orientation
Core concepts
Energy changes in reactions
All chemical reactions involve energy changes. Breaking bonds requires energy (endothermic process), while making bonds releases energy (exothermic process). The overall energy change determines whether a reaction is exothermic or endothermic.
Exothermic reactions:
- Transfer energy to surroundings
- Temperature of surroundings increases
- Energy level of products is lower than reactants
- Examples: combustion, oxidation, neutralisation, many displacement reactions
- Everyday applications: hand warmers, self-heating cans
Endothermic reactions:
- Take in energy from surroundings
- Temperature of surroundings decreases
- Energy level of products is higher than reactants
- Examples: thermal decomposition, photosynthesis, dissolving some salts
- Everyday applications: sports injury cold packs
Reaction profiles:
Reaction profile diagrams show the energy changes during reactions. The vertical axis represents energy, the horizontal axis represents reaction progress.
For exothermic reactions:
- Reactants start at a higher energy level
- Products end at a lower energy level
- The activation energy is shown as the peak from reactants
- Overall energy change (ΔH) is negative
For endothermic reactions:
- Reactants start at a lower energy level
- Products end at a higher energy level
- The activation energy is shown as the peak from reactants
- Overall energy change (ΔH) is positive
The activation energy appears as a "hump" that must be overcome for both reaction types.
Rates of reaction
The rate of reaction measures how quickly reactants form products. You can measure this by monitoring:
- Volume of gas produced per unit time
- Mass loss per unit time
- Change in colour or turbidity
- Change in pH
Factors affecting reaction rates:
1. Concentration (solutions) or pressure (gases)
- Higher concentration/pressure means more particles in same volume
- Particles are closer together
- More frequent successful collisions per unit time
- Faster reaction rate
2. Temperature
- Higher temperature gives particles more kinetic energy
- Particles move faster and collide more frequently
- More importantly, a greater proportion of collisions exceed activation energy
- Faster reaction rate
- Even a small temperature increase significantly increases rate
3. Surface area (solid reactants)
- Smaller pieces/powder have larger surface area to volume ratio
- More particles exposed for collisions
- More frequent successful collisions
- Faster reaction rate
4. Catalysts
- Provide alternative reaction pathway with lower activation energy
- More collisions now have sufficient energy to react
- Faster reaction rate without being consumed
- Examples: iron in Haber process, platinum in catalytic converters, enzymes in biological systems
Collision theory explanation:
For a reaction to occur, particles must:
- Collide with each other
- Have energy equal to or greater than activation energy
- Have correct orientation (though this isn't always emphasised at GCSE)
Increasing concentration, temperature, or surface area, or adding a catalyst, all increase the frequency of successful collisions or the proportion of collisions with sufficient energy.
Measuring rates:
Calculate mean rate of reaction using:
Rate = amount of reactant used or product formed / time
Units depend on what you measure:
- cm³/s (volume of gas)
- g/s (mass change)
- mol/s (concentration change)
The gradient of a concentration-time or amount-time graph gives the rate at that moment.
Reversible reactions and equilibrium
Some reactions are reversible — products can react to reform reactants. The ⇌ symbol shows reversibility.
Example: Hydrated copper sulfate ⇌ anhydrous copper sulfate + water
When a reversible reaction occurs in a closed system, it reaches dynamic equilibrium where:
- Forward and reverse reactions continue
- They occur at equal rates
- Concentrations of all reactants and products remain constant
- Macroscopic properties (colour, pH, pressure) remain constant
Le Chatelier's principle:
When conditions change, the equilibrium position shifts to counteract the change.
Effect of concentration:
- Increasing concentration of reactant shifts equilibrium right (toward products)
- Increasing concentration of product shifts equilibrium left (toward reactants)
- System tries to remove the added substance
Effect of temperature:
- For exothermic forward reaction: increasing temperature shifts equilibrium left (favours endothermic reverse reaction)
- For endothermic forward reaction: increasing temperature shifts equilibrium right (favours endothermic forward reaction)
- System moves in direction that opposes temperature change
Effect of pressure (gases only):
- Increasing pressure shifts equilibrium toward side with fewer gas molecules
- Decreasing pressure shifts equilibrium toward side with more gas molecules
- System tries to reduce pressure increase
Catalysts and equilibrium:
- Catalysts speed up both forward and reverse reactions equally
- Equilibrium position unchanged
- Equilibrium reached more quickly
Industrial applications:
The Haber process (manufacturing ammonia):
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol (exothermic)
Optimal conditions balance yield with rate:
- Pressure: 200 atmospheres (high pressure favours fewer molecules, but very high pressure is expensive and dangerous)
- Temperature: 450°C (compromise — lower temperature increases yield but slows reaction; higher temperature needed for acceptable rate)
- Catalyst: iron (speeds up reaction without affecting position)
Calculating quantities in reactions
Using balanced equations:
Balanced symbol equations show the ratio of moles (or particles) reacting. This stoichiometry allows quantitative predictions.
Example: 2Mg + O₂ → 2MgO
This shows:
- 2 moles of magnesium react with 1 mole of oxygen
- 2 atoms of magnesium react with 1 molecule of oxygen
- The ratio is always 2:1:2
Mass calculations:
Use relative atomic masses (Ar) or relative formula masses (Mr) to calculate reacting masses.
Steps:
- Write balanced equation
- Identify known and unknown substances
- Calculate moles of known substance (mass ÷ Mr)
- Use equation ratio to find moles of unknown
- Calculate mass of unknown (moles × Mr)
Limiting reactants:
When reactants aren't in exact stoichiometric ratio, one runs out first. This is the limiting reactant — it determines the maximum amount of product formed. The other reactant is in excess.
Energy calculations
Using bond energies:
Energy change = total energy required to break bonds - total energy released making bonds
For exothermic reactions: more energy released making bonds than used breaking bonds (negative overall)
For endothermic reactions: more energy used breaking bonds than released making bonds (positive overall)
Steps:
- Draw out all bonds in reactants and products (or use displayed formulae)
- Sum all bond energies in reactants (energy in)
- Sum all bond energies in products (energy out)
- Calculate: Energy in - Energy out
Remember: breaking bonds requires energy (positive), making bonds releases energy (negative in overall calculation)
Worked examples
Example 1: Le Chatelier's principle
The Contact process produces sulfur trioxide: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = -197 kJ/mol
(a) Predict and explain the effect of increasing pressure on the yield of sulfur trioxide. [3 marks]
Answer:
- Yield increases [1 mark]
- Because increasing pressure shifts equilibrium to the side with fewer gas molecules [1 mark]
- There are 3 molecules on left (2SO₂ + O₂) and 2 molecules on right (2SO₃) [1 mark]
(b) Predict and explain the effect of increasing temperature on the yield of sulfur trioxide. [3 marks]
Answer:
- Yield decreases [1 mark]
- Because increasing temperature shifts equilibrium in the endothermic direction [1 mark]
- The forward reaction is exothermic, so reverse (endothermic) reaction is favoured [1 mark]
Example 2: Rate of reaction calculation
A student reacts magnesium ribbon with hydrochloric acid. They collect 48 cm³ of hydrogen gas in 30 seconds.
Calculate the mean rate of reaction. Give your answer in cm³/s. [2 marks]
Answer:
- Rate = volume / time = 48 / 30 [1 mark]
- = 1.6 cm³/s [1 mark]
Example 3: Mass calculation
Calculate the mass of magnesium oxide produced when 6 g of magnesium reacts completely with oxygen. (Ar: Mg = 24, O = 16)
2Mg + O₂ → 2MgO [2 marks]
Answer:
- Moles of Mg = 6 / 24 = 0.25 mol
- From equation: 2 mol Mg produces 2 mol MgO, so ratio is 1:1
- Moles of MgO = 0.25 mol [1 mark]
- Mr of MgO = 24 + 16 = 40
- Mass = 0.25 × 40 = 10 g [1 mark]
Common mistakes and how to avoid them
Confusing exothermic and endothermic: Remember EXothermic = EXits (energy exits to surroundings). In exothermic reactions, the surroundings get hotter, not the reactants.
Thinking catalysts shift equilibrium position: Catalysts only speed up the rate at which equilibrium is reached. They have no effect on the position of equilibrium or the amounts of products at equilibrium.
Forgetting that equilibrium is dynamic: At equilibrium, reactions haven't stopped — both forward and reverse reactions continue at equal rates. Concentrations are constant, not zero.
Incorrect stoichiometry in calculations: Always check the ratio from the balanced equation. If the equation shows 2:1, don't assume 1:1. Write out the ratio explicitly.
Confusing bond breaking and making: Bond breaking always requires energy (endothermic), bond making always releases energy (exothermic). Whether the overall reaction is exothermic or endothermic depends on which process releases/requires more energy.
Missing units in rate calculations: Always include units with your answer (e.g., cm³/s, g/s). Check what the question measures and use those units with "/s" or "/min".
Exam technique for "C3: Chemical Reactions"
"Explain" questions require reasoning: Don't just state what happens — explain why using collision theory or Le Chatelier's principle. Link cause and effect explicitly for full marks.
Use data from graphs: When asked to determine rate from a graph, draw a tangent or calculate gradient using two clear points. Show your working: rate = change in y / change in x.
Write balanced symbol equations when calculating masses: Even if not explicitly asked, examiners expect to see the equation to award method marks. Show moles calculations clearly.
Command words matter: "State" needs a simple answer; "Explain" requires reasoning; "Calculate" must show working; "Suggest" allows reasonable answers based on knowledge applied to unfamiliar contexts.
Quick revision summary
Chemical reactions involve energy changes — exothermic reactions release energy while endothermic reactions absorb it. Reaction rates depend on concentration, temperature, surface area, and catalysts, explained by collision theory. Reversible reactions in closed systems reach dynamic equilibrium. Le Chatelier's principle predicts how equilibrium position responds to changes in concentration, temperature, or pressure. Balanced equations enable quantitative calculations of reacting masses using moles. Bond energy calculations determine overall energy changes in reactions.