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📌 Illustrated Diagram: Activation Energy

Chemical Kinetics – Complete CBSE Class 12 Chemistry Notes

Unit 3 | Chemical Kinetics | CBSE 2026–27

1. Rate of Reaction

Rate measures the change in concentration with time. For a reactant, concentration decreases; for a product, it increases.

2. Factors Affecting Rate

• Concentration/pressure of reactants
• Temperature
• Nature of reactants
• Surface area for suitable heterogeneous reactions
• Catalyst

3. Rate Law and Order

For rate = k[A]^m[B]ⁿ, order = m+n. The order must be determined experimentally; it should not automatically be copied from stoichiometric coefficients.

4. Zero and First Order

Worked Example

For a first-order reaction with k = 0.231 min⁻¹, t½ = 0.693/0.231 = 3.0 min.

5. Arrhenius Equation

k = Ae^−Ea/RT. The logarithmic form allows activation energy to be determined from temperature-dependent rate constants.

6. Collision Theory

Not every collision produces reaction. Effective collisions require sufficient energy and suitable orientation. A catalyst offers an alternative pathway with lower activation energy and therefore can increase reaction rate.

🧠 Energy Profile

📝 Case-Based Question

A reaction is studied at two temperatures and its rate constant increases significantly at the higher temperature.

1. Which Arrhenius parameter explains this temperature sensitivity?
2. What happens to the fraction of molecules having energy ≥ activation energy?
3. Would a catalyst change the activation energy pathway?
4. Would a catalyst change the equilibrium constant of the reaction?

🎯 Board Practice

1. Determine reaction order from experimental rate data.
2. Calculate k and half-life for a first-order reaction.
3. Use two temperature-rate data points to determine activation energy.
4. Distinguish rate law from stoichiometric equation using an example.

MCQs

1. Order is determined experimentally. True.
2. For a first-order reaction, half-life is independent of initial concentration. True.
3. A catalyst changes the equilibrium constant. False.
4. The Arrhenius factor explains temperature dependence of k. True.

Deep-Dive: Rate Laws

The rate law connects reaction rate with reactant concentrations. For a reaction with rate=k[A]^m[B]^n, the overall order is m+n.

Integrated First-Order Equation

For a first-order reaction, k=(2.303/t)log(a/(a−x)). Its half-life is t₁/₂=0.693/k and is independent of initial concentration.

Worked Numerical

If k=0.231 min⁻¹, t₁/₂=0.693/0.231=3.00 min.

Collision Theory

Effective collisions require sufficient energy and suitable orientation. Increasing temperature raises the fraction of molecules able to overcome activation energy.

Arrhenius Equation

k=Ae^(−E_a/RT). A plot of ln k against 1/T has slope −E_a/R.

Common Mistakes

• Confusing order with stoichiometric coefficients.
• Using the first-order half-life formula for another order.
• Forgetting units of k depend on reaction order.

Practice

1. Determine order from a rate-law table.
2. Calculate k and half-life for a first-order reaction.
3. Use Arrhenius data at two temperatures.
4. Interpret an energy profile diagram.

Chemical Kinetics — Reaction Rate Map

Rate Law

For rate=k[A]^m[B]ⁿ, the overall order is m+n. The reaction order must be obtained from experimental data unless it follows directly from an elementary-step mechanism.

First-Order Reaction

For a first-order reaction, k=(2.303/t)log(a/(a−x)). Its half-life is t_1/2=0.693/k and is independent of the initial concentration.

Worked Numerical

If k=0.231 min⁻¹ for a first-order reaction, t_1/2=0.693/0.231=3.00 min.

Arrhenius Equation

k=Ae^−E_a/RT. Increasing temperature generally increases the fraction of molecules having energy at or above the activation energy.

Exam Traps

• Reaction order and stoichiometric coefficients are not automatically identical.
• Check units of k for the order involved.
• Do not use the first-order half-life formula for every reaction.

CBSE Application Practice

1. Determine order from experimental rate data.
2. Calculate k and half-life.
3. Use Arrhenius data to determine activation energy.
4. Interpret a concentration-time graph.