Chemical Kinetics
Chemical kinetics studies the rate (speed) of chemical reactions and the factors that influence it — essential for understanding how to control reactions in industry and biology.
A. Rate of Reaction
Defining Reaction Rate
Rate of Reaction
Rate = Change in concentration / Time
Unit: mol/L/s or mol dm⁻³ s⁻¹ · Can be measured by decrease in reactant OR increase in product concentration over time
- Rate can be measured by: change in mass, volume of gas produced, colour change, pH change, or conductivity
- Rate is fastest at the start (highest reactant concentration) and slows as reactants are used up
- A rate-time graph shows a curve that flattens when the reaction stops (reactant exhausted)
B. Factors Affecting Rate of Reaction
Six Factors
Temperature
↑ Temp → ↑ Rate
Higher temperature → particles have more kinetic energy → more frequent AND more energetic collisions → more particles exceed activation energy
Concentration
↑ Conc → ↑ Rate
Higher concentration → more particles per unit volume → more frequent collisions between reactant particles
Surface Area
↑ Surface Area → ↑ Rate
Powder reacts faster than lumps — more surface exposed for collisions. E.g. flour dust in mills can explode; iron filings react faster than iron block
Catalyst
↑ Rate (not consumed)
Provides an alternative reaction pathway with LOWER activation energy → more particles have enough energy to react. Not used up — can be recycled
Pressure (gases)
↑ Pressure → ↑ Rate
Same as increasing concentration — gas particles pushed closer together → more frequent collisions. Only applies to gas-phase reactions
Light
↑ Light → ↑ Rate (some reactions)
Photochemical reactions are initiated or accelerated by light — light provides energy for activation. E.g. photosynthesis, photography, bleaching
C. Collision Theory
What Makes a Successful Collision?
For a reaction to occur, colliding particles must satisfy both conditions:
- Sufficient energy: the collision energy must be ≥ the activation energy (Eₐ) — the minimum energy needed to break bonds and start the reaction
- Correct orientation: the particles must collide in the right geometric arrangement for bond breaking/forming to occur
Activation Energy (Eₐ)
Minimum energy required for a reaction to occur
A catalyst lowers Eₐ → more particles have sufficient energy → faster rate at the same temperature
⚡ MCQ Tip Not ALL collisions lead to reaction — only those with enough energy (≥ Eₐ) AND correct orientation. Increasing temperature doesn't just increase collision frequency — it increases the proportion of particles with energy ≥ Eₐ (Maxwell-Boltzmann distribution shifts right).
D. Energy Profile Diagrams
Reading Energy Profile Diagrams
| Feature | Exothermic | Endothermic |
|---|---|---|
| Products vs Reactants energy | Products LOWER than reactants | Products HIGHER than reactants |
| ΔH | Negative (−) | Positive (+) |
| Peak of diagram | Transition state (highest energy point) | Transition state (highest energy point) |
| Effect of catalyst | Lowers peak (Eₐ) — same ΔH | Lowers peak (Eₐ) — same ΔH |
⚡ MCQ Tip A catalyst lowers Eₐ but does NOT change ΔH (enthalpy change) or the energy of reactants/products. The peak of the energy profile = transition state (activated complex) — highest energy point. Catalyst = provides an alternative lower-energy pathway.
Quick MCQ Revision
| Fact | Answer |
|---|---|
| Rate of reaction formula | Rate = Change in concentration ÷ Time |
| Higher temperature → rate | Increases — more energetic, frequent collisions; more particles ≥ Eₐ |
| Catalyst does what | Lowers activation energy (Eₐ) — faster rate, NOT consumed, does NOT change ΔH |
| Collision theory requirements | Sufficient energy (≥ Eₐ) AND correct orientation |
| Activation energy (Eₐ) | Minimum energy needed for reaction to occur |
| Surface area ↑ → rate | Increases — powder reacts faster than lumps |
| Transition state on energy profile | The peak — highest energy point of the reaction pathway |
| Catalyst on energy profile | Lowers the peak (Eₐ) without changing reactant/product energy levels |