What you'll learn
This revision guide covers rates of reaction and energy changes, two fundamental topics in Edexcel GCSE Chemistry. You'll understand how to measure and control reaction speeds using collision theory, and how to analyse energy transfers in chemical reactions. These concepts apply to industrial processes, environmental chemistry, and everyday reactions.
Key terms and definitions
Rate of reaction — the speed at which reactants are converted into products, measured as the change in quantity of reactant or product per unit time (e.g. g/s or cm³/s)
Activation energy — the minimum energy that colliding particles must have for a reaction to occur
Catalyst — a substance that increases the rate of reaction without being used up, by providing an alternative pathway with lower activation energy
Exothermic reaction — a reaction that transfers energy to the surroundings, causing the temperature to increase (ΔH is negative)
Endothermic reaction — a reaction that takes in energy from the surroundings, causing the temperature to decrease (ΔH is positive)
Collision theory — the concept that chemical reactions can only occur when particles collide with sufficient energy (activation energy) and the correct orientation
Reaction profile — a diagram showing the energy changes during a chemical reaction, with energy on the y-axis and progress of reaction on the x-axis
Enthalpy change (ΔH) — the overall energy change in a reaction, measured in kJ/mol
Core concepts
Measuring rates of reaction
The rate of reaction can be measured by monitoring either the disappearance of reactants or the appearance of products over time.
Methods for measuring rate:
- Gas collection — measure the volume of gas produced using a gas syringe or upturned measuring cylinder filled with water (suitable for reactions producing CO₂, H₂, O₂)
- Mass loss — place the reaction vessel on a balance and record the decrease in mass as gas escapes (works for reactions producing gas)
- Colour change — use a colorimeter or visual observation to track colour intensity changes
- Precipitation — time how long it takes for a solution to become cloudy enough that a cross underneath cannot be seen (e.g. sodium thiosulfate and hydrochloric acid)
Calculating rate:
Rate = amount of reactant used or product formed ÷ time
For example, if 48 cm³ of gas is produced in 20 seconds: Rate = 48 ÷ 20 = 2.4 cm³/s
The steeper the gradient on a graph of product formed against time, the faster the rate of reaction. The rate is fastest at the start of the reaction when reactant concentration is highest.
Factors affecting rates of reaction
Five main factors affect reaction rates, all explainable using collision theory.
Concentration (solutions) and pressure (gases):
- Increasing concentration means more particles in the same volume
- More frequent collisions between reactant particles
- Higher rate of successful collisions
- Example: 2 mol/dm³ hydrochloric acid reacts faster with calcium carbonate than 1 mol/dm³ acid
Temperature:
- Increasing temperature gives particles more kinetic energy
- Particles move faster and collide more frequently
- More particles have energy equal to or greater than the activation energy
- Even a small temperature increase (e.g. 10°C) can significantly increase rate
- Example: magnesium ribbon reacts vigorously with hot acid but slowly with cold acid
Surface area (solid reactants):
- Smaller pieces or powdered solids have larger surface area to volume ratio
- More particles are exposed to react
- More collisions can occur
- Example: powdered calcium carbonate reacts faster with acid than large marble chips
Catalysts:
- Provide an alternative reaction pathway with lower activation energy
- More particles have sufficient energy to react
- Increase rate without being consumed in the reaction
- Specific to particular reactions (e.g. iron in the Haber process, manganese(IV) oxide for hydrogen peroxide decomposition)
Presence of light:
- Some reactions are affected by light (photochemical reactions)
- Not commonly tested at GCSE but worth noting (e.g. photosynthesis, photographic film)
Collision theory and activation energy
For a reaction to occur, particles must:
- Collide with each other
- Have sufficient energy (at least the activation energy)
- Have the correct orientation
Not all collisions result in a reaction. Only successful collisions — those with energy ≥ activation energy and correct orientation — lead to products.
Factors that increase the rate work by either:
- Increasing collision frequency (concentration, pressure, surface area, temperature)
- Increasing the proportion of successful collisions (temperature, catalysts)
Temperature is unique because it affects both collision frequency AND the energy of collisions.
Energy changes in reactions
All chemical reactions involve energy changes as bonds are broken and formed.
Bond breaking and bond making:
- Breaking bonds requires energy (endothermic process)
- Making bonds releases energy (exothermic process)
- The overall energy change depends on the difference between energy in and energy out
Exothermic reactions:
- Energy released from forming new bonds > energy needed to break original bonds
- Products have less energy than reactants
- Temperature of surroundings increases
- ΔH is negative (e.g. ΔH = -92 kJ/mol)
- Examples: combustion, neutralisation, oxidation reactions, many displacement reactions
Endothermic reactions:
- Energy needed to break original bonds > energy released from forming new bonds
- Products have more energy than reactants
- Temperature of surroundings decreases
- ΔH is positive (e.g. ΔH = +178 kJ/mol)
- Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate in water
Reaction profiles
Reaction profiles show the energy pathway from reactants to products.
Exothermic reaction profile:
- Reactants start at higher energy level
- Products finish at lower energy level
- Energy difference between reactants and products = enthalpy change (negative)
- Activation energy shown as peak from reactant level
Endothermic reaction profile:
- Reactants start at lower energy level
- Products finish at higher energy level
- Energy difference = enthalpy change (positive)
- Activation energy shown as peak from reactant level
Effect of catalysts on profiles:
- Catalysts lower the activation energy
- Shown as a lower peak on the profile
- Same start and end points (same ΔH)
- Alternative pathway with reduced energy barrier
Calculating energy changes
You can calculate the overall energy change using bond energies.
Method:
- List all bonds broken in reactants (energy in)
- List all bonds formed in products (energy out)
- Calculate: Energy change = Energy in - Energy out
If the result is:
- Negative → exothermic reaction
- Positive → endothermic reaction
Example calculation structure:
For the reaction: CH₄ + 2O₂ → CO₂ + 2H₂O
Bonds broken: 4 × C-H, 2 × O=O Bonds formed: 2 × C=O, 4 × O-H
Energy in = (4 × 412) + (2 × 496) = 2640 kJ Energy out = (2 × 743) + (4 × 463) = 3338 kJ Energy change = 2640 - 3338 = -698 kJ/mol (exothermic)
Worked examples
Example 1: Calculating rate from a graph
Question: A student investigates the reaction between magnesium ribbon and hydrochloric acid by measuring the volume of hydrogen gas produced. The graph shows their results.
At 30 seconds, 60 cm³ of gas has been produced. At 50 seconds, 80 cm³ has been produced.
Calculate the mean rate of reaction between 30 and 50 seconds. [3 marks]
Answer:
Volume of gas produced = 80 - 60 = 20 cm³ [1 mark]
Time taken = 50 - 30 = 20 s [1 mark]
Rate = 20 ÷ 20 = 1.0 cm³/s [1 mark]
Examiner tip: Always show your working. Include units in your final answer. Subtract initial values from final values to find the change.
Example 2: Explaining the effect of temperature
Question: Explain, in terms of particles, why increasing the temperature increases the rate of reaction between sodium thiosulfate solution and hydrochloric acid. [4 marks]
Answer:
At higher temperature, particles have more kinetic energy [1 mark]
Particles move faster and collide more frequently [1 mark]
More particles have energy greater than or equal to the activation energy [1 mark]
Therefore there are more successful collisions per second / higher proportion of collisions are successful [1 mark]
Examiner tip: For 'explain' questions worth 4 marks, aim for 4 distinct points. Link particle behaviour to collision theory and then to successful collisions.
Example 3: Energy change calculation
Question: Use the bond energies in the table to calculate the energy change for this reaction:
N₂ + 3H₂ → 2NH₃
| Bond | Bond energy (kJ/mol) |
|---|---|
| N≡N | 945 |
| H-H | 436 |
| N-H | 391 |
[4 marks]
Answer:
Bonds broken: 1 × N≡N and 3 × H-H [1 mark] Energy in = 945 + (3 × 436) = 945 + 1308 = 2253 kJ [1 mark]
Bonds formed: 6 × N-H (2 NH₃ molecules, each with 3 N-H bonds) [1 mark] Energy out = 6 × 391 = 2346 kJ [1 mark]
Energy change = 2253 - 2346 = -93 kJ/mol [1 mark]
(The reaction is exothermic) [1 mark if stated]
Examiner tip: Count bonds carefully — each NH₃ has 3 N-H bonds. Show all calculations. Include the sign (+ or -) in your answer.
Common mistakes and how to avoid them
Confusing rate with time — Rate is how fast something happens (amount per second). A faster reaction has a higher rate but takes less time. Don't mix these up when describing results.
Thinking catalysts are used up — Catalysts are not consumed in reactions. They participate but are regenerated. Write "catalyst" not "reactant" when describing their role.
Wrongly explaining surface area — Don't say "more surface area means more particles" — the number of particles stays the same. Say "more particles are exposed at the surface" or "larger surface area available for collisions."
Reversing bond energy calculations — Breaking bonds always requires energy (positive values). Making bonds always releases energy. The confusion comes in the final calculation: Energy change = Energy to break bonds - Energy from making bonds.
Misreading reaction profile diagrams — The activation energy is measured from the reactant level to the peak, not from the x-axis. The enthalpy change (ΔH) is the difference between reactants and products.
Using inappropriate methods — You cannot measure the rate of a precipitation reaction by collecting gas. Match the measurement method to the products formed: gas collection for gases, mass loss for gas escaping, colorimetry for colour changes.
Exam technique for "Rates of Reaction and Energy Changes"
Command word awareness — "Explain" questions require reasoning (because/therefore). "Describe" needs observations or patterns only. "Calculate" always requires working and units. For 3-4 mark explanations, use collision theory terms: frequency of collisions, energy of collisions, activation energy, successful collisions.
Graph questions — The steepest part of a curve shows the fastest rate. When the line becomes horizontal, the reaction has finished. Calculate rate from a graph using gradient = change in y ÷ change in x. Draw tangent lines carefully using a ruler for instantaneous rate.
Practical questions — Know control variables for rate experiments (volume of acid, mass of solid, temperature, surface area). Identify independent variable (what you change) and dependent variable (what you measure). Suggest improvements like repeating for reliability or using a data logger for accuracy.
Energy calculations — Set out bond energy calculations clearly in a table or list. Count all bonds including those in coefficients (e.g., 2H₂O means 4 O-H bonds). Remember: positive ΔH = endothermic, negative ΔH = exothermic.
Quick revision summary
Rates of reaction depend on collision frequency and energy. Increasing concentration, temperature, surface area, or adding a catalyst increases rate. Collision theory explains this: particles must collide with sufficient energy (activation energy) for reaction. Exothermic reactions release energy (negative ΔH), endothermic reactions absorb energy (positive ΔH). Reaction profiles show activation energy and enthalpy change. Calculate energy changes using bond energies: energy in (breaking bonds) minus energy out (making bonds). Catalysts lower activation energy without being consumed.