Factors Affecting Reaction Rate – Complete Study Notes
The rate of a chemical reaction is defined as the speed at which reactants are converted into products. It is measured as the change in concentration of a reactant or product per unit time. Understanding what controls this rate is essential for industrial chemistry, pharmacology, and food science.
Unit of Rate: mol L⁻¹ s⁻¹ (moles per litre per second). The rate can be expressed as: Rate = –Δ[Reactant]/Δt = +Δ[Product]/Δt
1. Concentration of Reactants
Increasing the concentration of reactants increases the number of reacting particles per unit volume. This leads to more frequent collisions between molecules and therefore a higher rate of reaction.
- For the reaction: A + B → Products, Rate = k[A]ᵐ[B]ⁿ (Rate Law expression)
- The order of reaction (m + n) is determined experimentally, not from stoichiometry.
- Example: Doubling the concentration of HCl doubles the rate of reaction with marble chips (first-order in HCl).
2. Temperature
Temperature is one of the most important factors. As temperature rises, molecules gain more kinetic energy, leading to more frequent and more energetic collisions. The fraction of molecules with energy greater than or equal to the activation energy (Ea) increases dramatically.
- Temperature Coefficient (Q₁₀): For every 10°C rise in temperature, the rate of reaction approximately doubles (Q₁₀ ≈ 2).
- Arrhenius Equation: k = Ae^(–Ea/RT). This equation relates the rate constant (k) to temperature (T) and activation energy (Ea). A plot of ln(k) vs 1/T gives a straight line with slope = –Ea/R.
- This is why food is refrigerated (to slow microbial reactions) and why reactions in a pressure cooker are faster (higher temperature).
3. Surface Area
For reactions involving solids, the surface area exposed to the reactant is crucial. Smaller particles (powders) have a much greater surface area per unit mass than larger lumps.
- Powdered calcium carbonate reacts much faster with hydrochloric acid than marble chips of the same mass.
- This principle is used in industrial processes: catalysts are made as fine powders or porous pellets to maximize surface contact.
- Dust explosions in flour mills and coal mines occur because fine particles have enormous surface area, allowing explosive combustion rates.
4. Catalyst
A catalyst increases the rate of reaction without being permanently consumed. It provides an alternative reaction pathway with a lower activation energy. This means more molecules have sufficient energy to react, and the rate increases.
- Positive Catalyst: Increases rate (e.g., MnO₂ in decomposition of H₂O₂).
- Negative Catalyst (Inhibitor): Decreases rate (e.g., glycerin added to H₂O₂ for stability).
- A catalyst does not change the enthalpy (ΔH) or the equilibrium constant (K) of the reaction — only the rate at which equilibrium is reached.
5. Pressure (for Gaseous Reactions)
For reactions involving gases, increasing pressure effectively increases the concentration of gas molecules per unit volume, leading to a higher reaction rate. This is important in industrial processes like the Haber Process for ammonia synthesis.
6. Nature of Reactants
The inherent reactivity of the substances involved also matters. Ionic reactions (between ions in solution) are almost instantaneous because no bond-breaking is needed — just ion recombination. Covalent reactions are generally slower because covalent bonds must be broken first.
7. Collision Theory and Activation Energy
According to Collision Theory, for a reaction to occur, molecules must:
- Collide with each other.
- Have energy ≥ the Activation Energy (Ea) — the minimum energy required for the reaction to proceed.
- Have the correct orientation at the moment of collision.
An activated complex (transition state) is formed momentarily at the peak of the energy profile diagram. It is highly unstable and quickly breaks down into either products or reactants.
Key Exam Tips
- Catalyst does NOT change K (equilibrium constant) or ΔH — only rate and activation energy.
- Zero order reactions: rate is independent of concentration.
- First order reactions: rate is directly proportional to concentration (e.g., radioactive decay).
- Pseudo-first-order reactions occur when one reactant is in large excess (e.g., hydrolysis of esters in water).