Chemical kinetics, collision theory, rate curves, and four key factors — understand why some reactions are fast as lightning while others take centuries.
Section 1 — What is the Rate of Reaction?
Units: mol dm⁻³ s⁻¹ | cm³ s⁻¹ | g s⁻¹
How Do We Measure Rate? (Measurable Properties)
| Method | What's Measured | Example Reaction |
|---|---|---|
| Gas syringe / graduated cylinder | Volume of gas produced over time | Mg + HCl → H₂ gas collected |
| Mass balance | Decrease in mass as gas escapes | CaCO₃ + HCl → CO₂ escapes, mass falls |
| Cloudiness / turbidity | Time for solution to turn cloudy (cross experiment) | Na₂S₂O₃ + HCl → S precipitate forms |
| Colour intensity | Change in colour of a coloured reactant/product | Bleaching of dye by KMnO₄ |
Section 2 — Collision Theory
For a reaction to occur, bonds must break in reactants and new bonds must form in products. But not every collision causes a reaction!
Three Conditions for an Effective Collision
- 1Collision — Reactant particles MUST collide with each other. No collision = no reaction.
- 2Activation Energy — Particles must collide with enough energy (≥ activation energy, Ea) to break existing bonds.
- 3Correct Orientation — Particles must be aligned correctly so energy can be transferred to the bonds that need to break.
Section 3 — Rate Curves
A rate curve plots a measurable property (e.g. volume of gas, mass of flask) against time. All rate curves share the same characteristic shape.
| Rate Curve Feature | What It Means |
|---|---|
| Steep gradient at start | Reaction is FASTEST — reactant concentration is highest, most collisions per second |
| Gradient becomes shallower | Reaction SLOWING DOWN — reactants being used up, fewer collisions |
| Curve becomes flat (horizontal) | Reaction has STOPPED — limiting reactant fully used up |
| Same final level for faster/slower variants | Same total product made — limiting reactant amount is unchanged |
Worked Example — Calculating Average Rate
Section 4 — Factors Affecting Rate of Reaction
Effect: ↑ concentration → ↑ rate
Collision theory: More particles per unit volume → collisions happen more frequently → more effective collisions per second.
🌍 Bleach works faster when concentrated — same principle!
Effect: ↑ temperature → ↑ rate. Every +10°C roughly doubles the rate!
Collision theory: Particles gain kinetic energy → move faster (more frequent collisions) AND hit harder (more collisions exceed activation energy). Both effects increase effective collisions.
🌍 Refrigerators slow food spoilage by lowering reaction rates. Pressure cookers cook faster by raising temperature above 100°C.
Effect: Smaller particles (greater surface area) → ↑ rate
Collision theory: Reactions occur AT THE SURFACE of solids. Grinding solid into smaller pieces exposes MORE surface to the other reactant → more collisions per second.
⚠️ DANGER: Finely divided flour or coal dust in air can cause devastating explosions — enormous surface area means incredibly fast reaction with O₂!
Effect: Catalyst → ↑ rate without being permanently changed
Collision theory: Provides an alternative reaction pathway with lower activation energy. More collisions have enough energy to react — without needing higher temperature.
Example: MnO₂ catalyses decomposition of H₂O₂:
2H₂O₂(aq) →[MnO₂] 2H₂O(l) + O₂(g)
🌍 Catalytic converters use Pt/Pd to convert toxic exhaust gases. Enzymes are biological catalysts — without them, your body's reactions would be too slow to sustain life!
Section 5 — Effect on Rate Curves
When Rate INCREASES
- Steeper gradient at the start
- Horizontal sooner — completes faster
- Same final level — same total product made
When Rate DECREASES
- Shallower gradient at the start
- Horizontal later — takes longer to complete
- Same final level — if limiting reactant unchanged
Section 6 — Resources & Simulations
Section 7 — CSEC Practice Questions
a Rate (0–30 s) = ΔV ÷ Δt = (54 − 0) ÷ 30 = 1.8 cm³ s⁻¹
b Rate (60–90 s) = (106 − 88) ÷ 30 = 18 ÷ 30 = 0.60 cm³ s⁻¹
c At 30 s, the concentration of HCl is much higher (reactant barely used up). More HCl particles per unit volume → collisions are more frequent → more effective collisions per second → higher rate. By 90 s, HCl has been largely consumed → lower concentration → fewer collisions → lower rate.
1 Effect 1 — Increased collision frequency: At higher temperatures, particles have more kinetic energy and therefore move faster. Faster-moving particles collide with each other more frequently per second.
2 Effect 2 — More collisions exceed activation energy: Because particles move faster, they also collide with greater energy. A larger proportion of these collisions now have energy equal to or greater than the activation energy (Ea). More effective collisions per second → faster rate.
a MnO₂ acts as a catalyst — it increases the rate of decomposition of H₂O₂ without itself undergoing any permanent chemical change.
b 2H₂O₂(aq) → 2H₂O(l) + O₂(g) [gas produced = oxygen, test: glowing splint relights]
c Filter off the MnO₂ at the end of the reaction. Wash, dry and weigh it — the mass will be the same as at the start. You could also re-add it to fresh H₂O₂ solution to confirm it still catalyses the reaction.
d MnO₂ provides an alternative reaction pathway with a lower activation energy. More collisions now have enough energy to result in a reaction, so the rate increases without raising the temperature.
1 The new curve would have a steeper gradient at the start — the reaction proceeds faster.
2 The curve would become horizontal sooner — the reaction would be complete in less time.
3 The final level would be the same — the same mass of CaCO₃ is used, so the same total volume of CO₂ is produced. Only the rate has changed, not the amount of limiting reactant.
Explanation (collision theory): Powder has a much larger surface area than crystals. More CaCO₃ surface is exposed to the HCl acid. More acid particles can simultaneously collide with CaCO₃ particles → frequency of collisions increases → more effective collisions per second → higher rate.
(A) Increasing the temperature (B) Using marble powder instead of chips (C) Using a larger volume of the same concentration of acid (D) Increasing the concentration of the acid
A Would INCREASE rate — higher temperature gives particles more KE, more frequent and more energetic collisions.
B Would INCREASE rate — powder has greater surface area than chips, more frequent collisions.
C Using MORE volume of the SAME concentration means the same number of particles per cm³ but a larger total amount. The concentration (particles per cm³) is unchanged, so collision frequency per unit volume is unchanged. Rate does NOT increase (though more product will eventually be made).
D Would INCREASE rate — more acid particles per cm³ → more frequent collisions.