Rates of Reaction
Summary: Reaction rate depends on concentration, temperature, surface area, and catalysts. Explained by collision theory — particles must collide with sufficient energy (activation energy) and correct orientation. Tags: igcse chemistry reactions rates collision-theory Created: 2026-07-14 Last Updated: 2026-07-16
What is Rate of Reaction?
The rate of reaction measures how quickly a reaction happens — specifically, the change in amount of reactant used up or product formed per unit time.
- A fast reaction has a high rate (e.g., explosion, neutralisation)
- A slow reaction has a low rate (e.g., rusting of iron)
Reactions slow down over time as reactants are used up, so the rate is not constant throughout a reaction.
Collision Theory
Collision theory is the fundamental explanation for all reaction rates. For a reaction to occur between particles:
- Particles must collide with each other
- They must collide with sufficient energy — at least the activation energy (Ea)
- They must collide with the correct orientation (the right way around)
Activation energy = the minimum energy that colliding particles must have for a successful reaction to occur.
If any of these three conditions is not met, the particles simply bounce off each other and no reaction occurs.
Factors Affecting Rate of Reaction
There are five key factors, all explained by collision theory:
| Factor | Effect on Rate | Collision Theory Explanation |
|---|---|---|
| Concentration (solutions) | Increase concentration → rate increases | More particles per unit volume → particles are closer together → more frequent collisions per unit time → greater chance of successful collisions |
| Pressure (gases only) | Increase pressure → rate increases | Same number of particles in a smaller volume → particles are closer together → more frequent collisions per unit time |
| Surface area (solids) | Increase surface area (smaller pieces) → rate increases | More solid particles are exposed at the surface available to react → more frequent collisions per unit time between reactant particles |
| Temperature | Increase temperature → rate increases | Particles gain more kinetic energy → they move faster, causing more frequent collisions AND a greater proportion of particles have energy ≥ activation energy → more successful collisions per unit time |
| Catalyst | Adding a catalyst → rate increases | A catalyst provides an alternative reaction pathway with a lower activation energy → a greater proportion of particles have sufficient energy to react → more successful collisions per unit time |
Temperature is unique: it is the only factor that increases rate for TWO reasons — more frequent collisions AND more collisions exceeding activation energy.
Measuring Rates of Reaction
There are several experimental methods to measure rate.
| Method | What is measured | How rate is found | Suitable for |
|---|---|---|---|
| Gas syringe | Volume of gas collected | Plot volume vs. time; rate = gradient | Reactions that produce a gas (e.g., metal + acid, decomposition of H₂O₂) |
| Upside-down measuring cylinder / burette (over water) | Volume of gas collected | Same as gas syringe | Reactions producing gases that are insoluble in water |
| Mass loss (balance) | Decrease in mass as gas escapes | Plot mass vs. time; rate = gradient (often done with cotton wool plug to prevent liquid spray loss) | Reactions producing CO₂ or other dense gases (e.g., CaCO₃ + HCl) |
| Disappearing cross | Time for a cross/drawn mark to become obscured | Rate ∝ 1/time taken (arbitrary units: s⁻¹) | Reactions producing a precipitate (e.g., Na₂S₂O₃ + HCl → sulfur precipitate) |
| Colour change (colorimetry) | Time for a colour change or intensity measurement | Rate ∝ 1/time; or measure absorbance over time | Reactions with coloured reactants/products |
| pH change | Change in pH over time using a pH meter/data logger | Plot pH vs. time | Reactions involving acids or alkalis |
The disappearing cross method:
- Sodium thiosulfate (Na₂S₂O₃) reacts with hydrochloric acid (HCl) to produce a fine yellow sulfur precipitate that makes the solution cloudy
- A cross drawn on paper is placed under the conical flask
- Measure the time taken for the cross to “disappear” (become no longer visible from above)
- Rate can be expressed as: rate = 1 / time (unit: s⁻¹)
Interpreting Graphs
Rate graphs typically plot amount of product (or reactant) on the y-axis against time on the x-axis.
Key features of a rate graph:
- Steeper gradient = faster rate of reaction
- Gradient = rate at any point on the curve
- Horizontal/flat line = reaction is complete (all limiting reactant used up)
- The graph is steepest at the start (rate is fastest) because reactant concentration is highest
- The graph gradually flattens as reactants are used up
Comparing two graphs on the same axes:
| Scenario | Graph feature | What it means |
|---|---|---|
| Higher concentration / pressure | Steeper initial gradient; reaches same final amount faster | More particles per volume → faster rate |
| Higher temperature | Steeper initial gradient; reaches same final amount faster | More kinetic energy → faster rate |
| Greater surface area | Steeper initial gradient; reaches same final amount faster | More exposed particles → faster rate |
| Catalyst present | Steeper initial gradient; reaches same final amount faster | Lower activation energy → faster rate |
| More reactant used | Same initial gradient; higher final amount (higher plateau) | More substance to react, not a rate change |
Example 1: Two experiments are carried out using the same mass of marble chips (CaCO₃) and the same volume and concentration of HCl. In Experiment A, large chips are used; in Experiment B, powdered CaCO₃ is used.
- Both graphs reach the same final volume of CO₂ (same mass of CaCO₃)
- Graph B (powder) has a steeper initial gradient (faster rate due to larger surface area)
- Graph B reaches the plateau sooner
Example 2: Two experiments: Experiment X uses 50 cm³ of 2.0 mol/dm³ HCl; Experiment Y uses 50 cm³ of 1.0 mol/dm³ HCl. Both use excess marble chips.
- Graph X has a steeper initial gradient (higher concentration → faster rate)
- Graph X reaches a higher final plateau (more HCl → more CO₂ produced)
Explaining Each Factor in Terms of Collision Theory
Effect of concentration: Increasing the concentration increases the rate of reaction. A higher concentration means there are more particles per unit volume. The particles are closer together. This means there are more frequent collisions per unit time. This gives a greater chance of successful collisions per unit time. Therefore, the rate of reaction increases.
Effect of temperature: Increasing the temperature increases the rate of reaction. Particles gain more kinetic energy at a higher temperature. They move faster, so there are more frequent collisions per unit time. More importantly, a greater proportion of particles have energy greater than or equal to the activation energy. This means there is a greater chance of successful collisions per unit time. Temperature has a particularly large effect because it increases BOTH collision frequency AND the proportion of particles exceeding activation energy.
Effect of surface area: Using smaller pieces / powder increases the rate of reaction. A solid with smaller pieces has a larger total surface area. This means more particles are exposed and available to react. There are more frequent collisions per unit time between reactant particles. This gives a greater chance of successful collisions, so the rate increases.
Effect of a catalyst: A catalyst increases the rate of reaction without being used up. It provides an alternative reaction pathway with a lower activation energy. A greater proportion of particles now have energy greater than or equal to this lower activation energy. This means there is a greater chance of successful collisions per unit time, so the rate increases.
Catalysts
Definition: A catalyst is a substance that increases the rate of a chemical reaction but is chemically unchanged at the end of the reaction (it is not used up).
Key facts about catalysts:
- They are NOT included in the chemical equation (they are written above the arrow)
- They lower the activation energy by providing an alternative reaction pathway
- They do NOT change the energy of reactants or products — only the pathway between them
- They do NOT affect the position of equilibrium (only how fast equilibrium is reached)
- Only a small amount is needed (they are not consumed)
Enzymes are biological catalysts — proteins that catalyse reactions in living organisms. Examples:
- Amylase (breaks down starch)
- Catalase (breaks down hydrogen peroxide in cells)
- Protease and lipase (digestive enzymes)
Industrial catalysts:
| Catalyst | Process | Reaction |
|---|---|---|
| Iron (Fe) | Haber process | N₂ + 3H₂ ⇌ 2NH₃ (manufacture of ammonia) |
| Vanadium(V) oxide (V₂O₅) | Contact process | 2SO₂ + O₂ ⇌ 2SO₃ (manufacture of sulfuric acid — Stage 2) |
| Nickel (Ni) | Hydrogenation of alkenes | C₂H₄ + H₂ → C₂H₆ (hardening of vegetable oils to make margarine) |
| Platinum / rhodium | Catalytic converters | 2CO + 2NO → 2CO₂ + N₂ (in car exhausts) |
| Manganese(IV) oxide (MnO₂) | Decomposition of hydrogen peroxide | 2H₂O₂ → 2H₂O + O₂ |
Key Practical Methods
| Practical | Reaction | Factor Investigated | How to measure rate | Key variables to control |
|---|---|---|---|---|
| HCl + Na₂S₂O₃ (disappearing cross) | Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + S(s) + H₂O(l) | Concentration (vary conc. of Na₂S₂O₃) | Time for cross to disappear; rate = 1/t | Temperature, total volume, concentration of HCl |
| Marble chips + HCl | CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l) | Surface area (large chips vs. small chips vs. powder) | Volume of CO₂ collected (gas syringe) or mass loss | Mass of CaCO₃, concentration and volume of HCl, temperature |
| Same reaction, vary temperature | Same as above | Temperature (water bath at different temps) | Volume of CO₂ or mass loss | Mass of CaCO₃, concentration and volume of HCl, surface area of chips |
| Catalytic decomposition of H₂O₂ | 2H₂O₂(aq) → 2H₂O(l) + O₂(g) | Catalyst (MnO₂ used as catalyst) | Volume of O₂ collected (gas syringe) or rate of effervescence | Concentration and volume of H₂O₂, temperature, mass of MnO₂ |
Key Points
- Rate = change in amount / time
- Collision theory: particles must collide, with sufficient energy (≥ Ea), with correct orientation
- Increasing concentration/pressure: more particles per volume → more frequent collisions → faster rate
- Increasing surface area: more exposed particles → more frequent collisions → faster rate
- Increasing temperature: more kinetic energy → more frequent collisions AND more particles ≥ Ea → faster rate
- Catalyst: alternative pathway with lower Ea → more particles ≥ Ea → faster rate
- Catalyst is NOT used up and is chemically unchanged at the end
- Rate graphs: steeper gradient = faster rate; horizontal line = reaction complete
- Disappearing cross rate ∝ 1/time (units: s⁻¹)
- Industrial catalysts: Fe (Haber), V₂O₅ (Contact), Ni (hydrogenation), MnO₂ (H₂O₂ decomposition)
- Enzymes are biological catalysts
Key Concepts from Past Papers
- More frequent collisions between particles
- More particles have energy greater than or equal to the activation energy
- A catalyst provides an alternative reaction pathway with a lower activation energy
- Steeper gradient = faster rate; the graph levels off when one reactant is used up
- For the disappearing cross: rate = 1/time and the cross disappears because a precipitate of sulfur forms
Keywords from Past Papers
rate, reaction, increases, speeds, substance, lower, temperature, increasing, decreases, slower, faster, concentration, higher, idea, linear, collision, activation, energy, catalyst, surface, area, gradient, graph
Related Notes
- Reversible Reactions
- Redox Reactions
- Energetics of Reactions
- IGCSE-Chem-Index — Full IGCSE Chemistry index
Sources
- OpenStax Chemistry 2e — [Chapter 12: Kinetics], Rice University (free, CC BY 4.0)
- BBC Bitesize GCSE Chemistry — [Rates of Reaction], BBC (free educational resource)
- Cambridge IGCSE Chemistry 0620 — Syllabus Section 6: Chemical Reactions (Rates of Reaction), Cambridge Assessment International Education
- CK-12 Chemistry for High School — [Chapter 18: Chemical Kinetics], CK-12 Foundation (free, CC BY-NC 3.0)
Past Paper Sources
- 0620/32 Feb/March 2015: Q33(d)(i) (2m)
- 0620/32 Feb/March 2018: Q66(a)(ii) (1m)
- 0620/32 May/June 2018: Q33(a)(iii) (1m)
- 0620/33 May/June 2017: Q66(b)(iv) (1m)
- 0620/33 May/June 2019: Q55(d)(ii) (1m)
- 0620/33 May/June 2021: Q55(c)(iv) (2m), Q44(a)(ii) (0m)
- 0620/33 October/November 2019: Q66(d)(i) (1m)
- 0620/33 October/November 2020: Q22(c)(i) (1m)
- 0971/31 May/June 2018: Q33(b)(i) (1m)
- 0971/31 May/June 2021: Q55(b)(ii) (2m)
- 0971/32 May/June 2023: Q88(a)(ii) (1m)
Common Misconceptions
| Misconception | Reality |
|---|---|
| ”Catalysts are used up during the reaction” | Catalysts are NOT used up — they are chemically unchanged at the end. They can be reused. |
| ”Increasing temperature only makes particles collide more often” | Increasing temperature does TWO things: more frequent collisions AND a greater proportion of particles exceed activation energy. The second reason is often the more important one. |
| ”A catalyst increases the yield of a reaction” | A catalyst only affects RATE, not yield. It does not change the position of equilibrium or the amount of product formed. |
| ”If I use more reactant, the rate increases” | Using more reactant may increase the TOTAL amount of product but does NOT necessarily increase rate. Rate depends on concentration, not total mass. |
| ”Powdered solids are a different chemical with different properties” | Powdering only changes the surface area (a physical change), not the chemical nature of the substance. The same mass of powder and large lumps will produce the same amount of product. |
| ”Rate is constant throughout the reaction” | Rate is fastest at the start (highest concentration of reactants) and decreases over time as reactants are used up. |