Monitoring chemical reactions — rates, energy changes, reversible reactions and equilibrium
Rates of reaction
Rate of reaction = change in quantity of reactant or product ÷ time. Units: g/s, cm³/s, mol/dm³/s.
Measuring rate in OCR J258:
- Gas volume collected over time (using a gas syringe or upturned measuring cylinder over water).
- Mass loss on a balance (if a gas escapes).
- Colour change / absorption spectroscopy (for coloured products).
- Turbidity (cloudiness) — e.g. sodium thiosulfate + HCl → sulfur precipitate; time until a cross drawn under the flask disappears (PAG C4).
Factors affecting rate
All these factors increase rate by increasing the frequency of successful collisions:
| Factor | Effect on rate | Explanation (collision theory) |
|---|---|---|
| Concentration ↑ | Rate ↑ | More particles per unit volume → more frequent collisions |
| Temperature ↑ | Rate ↑ | Particles have more kinetic energy → more frequent collisions AND more particles exceed the activation energy |
| Particle size ↓ (surface area ↑) | Rate ↑ | More surface area exposed → more collisions possible per second |
| Catalyst | Rate ↑ | Provides an alternative pathway with lower activation energy; not consumed |
| Pressure ↑ (gases only) | Rate ↑ | Same as concentration for gases — more frequent collisions |
Activation energy: the minimum kinetic energy that colliding particles must have for a reaction to occur. A catalyst provides a route with lower activation energy (shown on an energy profile diagram as a lower "hill").
OCR PAG C4 — investigating rate of reaction (thiosulfate + HCl, or magnesium + acid): students control one variable at a time and plot results.
Energy changes in reactions
Chemical reactions involve breaking bonds (requiring energy — endothermic) and making bonds (releasing energy — exothermic). The overall energy change determines whether the reaction is exothermic or endothermic:
- Exothermic: energy released to surroundings (products have lower energy than reactants). Temperature of surroundings increases. ΔH is negative.
- Examples: combustion, neutralisation, oxidation of metals, respiration.
- Endothermic: energy absorbed from surroundings (products have higher energy than reactants). Temperature of surroundings decreases. ΔH is positive.
- Examples: thermal decomposition, photosynthesis, dissolving ammonium chloride.
Energy profile diagrams: reactants and products at different energy levels; the "hill" is the activation energy. Catalyst lowers the hill but does not change the overall energy change (ΔH).
Bond energies (Higher): ΔH = sum of energies required to break bonds − sum of energies released in making bonds. If bond-breaking energy > bond-making energy → endothermic (positive ΔH). If bond-making energy > bond-breaking energy → exothermic (negative ΔH).
Reversible reactions and equilibrium
A reversible reaction can proceed in both the forward and reverse directions: A + B ⇌ C + D
At dynamic equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. The concentrations of reactants and products remain constant (not necessarily equal). The equilibrium is maintained in a closed system only.
Le Chatelier's Principle (Higher)
If a system at equilibrium is disturbed, the equilibrium shifts to oppose the change:
| Change | Shift direction | Effect on yield of products |
|---|---|---|
| Increase temperature | Towards endothermic direction | Depends on which direction is endothermic |
| Decrease temperature | Towards exothermic direction | |
| Increase pressure | Towards fewer moles of gas | |
| Decrease pressure | Towards more moles of gas | |
| Increase concentration of reactant | Towards products | More products formed |
| Increase concentration of product | Towards reactants | |
| Add catalyst | No change to equilibrium position | Equilibrium reached faster; same yield |
Haber process application: N₂ + 3H₂ ⇌ 2NH₃ (ΔH = −92 kJ/mol, exothermic forward).
- High pressure (200 atm): favours fewer moles of gas (right side) → more NH₃.
- Low temperature: favours exothermic direction (more NH₃) BUT too slow.
- Compromise temperature ~450 °C and iron catalyst: reasonable rate AND yield.
OCR PAG C6 covers energy changes in reactions — students measure temperature changes to classify reactions as exo/endothermic.
AI-generated · claude-opus-4-7 · v3-ocr-chemistry