Rates of reaction
Collision theory
A chemical reaction occurs when reactant particles collide with sufficient energy (at least the activation energy) and with the correct orientation. The rate of reaction is the amount of product formed (or reactant used) per unit time.
rate = change in amount ÷ time (in mol/s, g/s, or cm³/s)
Factors that increase the rate of reaction all work by increasing the frequency of successful collisions and/or lowering the activation energy:
| Factor | Effect on rate | Explanation |
|---|---|---|
| Increase temperature | Increases rate | Particles have more kinetic energy → more frequent collisions AND higher proportion of collisions exceed Eₐ |
| Increase concentration | Increases rate | More particles per unit volume → more frequent collisions |
| Increase pressure (gases) | Increases rate | Same as concentration — particles closer together |
| Increase surface area | Increases rate | More particles exposed at surface → more collision opportunities |
| Add catalyst | Increases rate | Provides alternative reaction pathway with lower activation energy |
Edexcel Core Practical CP4 — Rate of reaction (HCl + Mg or marble chips)
HCl + Mg ribbon: Mg + 2HCl → MgCl₂ + H₂↑ Measure: volume of H₂ gas collected in a gas syringe vs time. Plot a graph: rate = gradient of tangent to the curve.
HCl + marble chips (CaCO₃): CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂↑ Measure: mass loss (CO₂ escapes) vs time on a balance, or volume of CO₂ in gas syringe.
Variables investigated: temperature (use water baths), concentration of HCl, surface area of marble (powdered vs lumps), catalysts.
Interpreting graphs: a steeper initial gradient = faster rate; curve levels off when a reactant is used up.
Catalysts
A catalyst is a substance that increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy. It is not consumed — it remains chemically unchanged at the end of the reaction (though it may be physically changed).
Homogeneous catalyst: in the same phase as the reactants (e.g. enzymes in biological reactions, H₂SO₄ catalyst in ester formation). Heterogeneous catalyst: in a different phase (e.g. solid Fe catalyst in the Haber process; Pt/Pd in catalytic converters).
Energy profile diagram: for a catalysed reaction, the activation energy peak is lower. The overall ΔH (enthalpy change) is the same — a catalyst does NOT change the thermodynamics.
Exothermic and endothermic reactions
Exothermic: energy is released to the surroundings. ΔH is negative. Temperature of surroundings increases. Examples: combustion, respiration, neutralisation, oxidation reactions.
Endothermic: energy is absorbed from the surroundings. ΔH is positive. Temperature of surroundings decreases. Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate in water.
Bond energy calculations (Higher)
Bond breaking requires energy (endothermic); bond forming releases energy (exothermic). ΔH = energy in (bonds broken) − energy out (bonds formed)
If ΔH < 0 (negative): exothermic (more energy released than absorbed). If ΔH > 0 (positive): endothermic (more energy absorbed than released).
⚠Common mistakes
- Catalyst and equilibrium: a catalyst does not change the position of equilibrium — it only speeds up reaching it.
- Rate vs yield: rate and yield are different. High temperature increases rate but may decrease yield (for exothermic reactions at equilibrium).
- Surface area misconception: smaller particle size = larger surface area = faster rate. Powdered marble reacts faster than lumps.
- Energy profile: the catalyst lowers Eₐ, but ΔH (difference between reactant and product energy) is unchanged.
- Bond energy sign: always positive for bond-breaking; total ΔH can be negative or positive.
AI-generated · claude-opus-4-7 · v3-edexcel-chemistry