Unit 3
General and Physical Chemistry
Class 12 Chemistry
Chemical KineticsClass 12 Chemistry – Chemical Kinetics Notes PDF
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Introduction
This unit covers average and instantaneous rates, rate law, rate constant, order and molecularity, integrated equations for zero- and first-order reactions, half-life, collision theory, activation energy, Arrhenius equation, catalysts and related numerical problems.
1. Rate of Reaction
For a reaction aA + bB → cC + dD:
Average Rate
Instantaneous Rate
The instantaneous rate is the slope of the concentration–time curve at a particular instant.
Fig. 1 – Average vs Instantaneous Rate
2. Rate Law and Rate Constant
For a reaction involving reactants A and B, an experimentally determined rate law may be:
Here k is the rate constant, while m and n are experimentally determined orders with respect to A and B.
The unit of k depends on overall reaction order.
| Overall Order | Typical Unit of k |
|---|---|
| Zero | mol L⁻¹ s⁻¹ |
| First | s⁻¹ |
| Second | L mol⁻¹ s⁻¹ |
Fig. 2 – Components of a Rate Law
3. Order and Molecularity
| Feature | Order | Molecularity |
|---|---|---|
| Meaning | Sum of powers of concentration terms in experimentally determined rate law. | Number of reacting species involved in a single elementary step. |
| How obtained | Experimentally. | From elementary mechanism step. |
| Possible values | Can be zero, integer or sometimes fractional. | Positive integer for an elementary step. |
| Applies to | Overall reaction rate expression. | Elementary reaction step. |
4. Zero-Order Reaction
A plot of [A] versus t is linear with slope −k.
Fig. 3 – Zero-Order Integrated Plot
5. First-Order Reaction
For a first-order reaction, half-life is independent of initial concentration.
Fig. 4 – First-Order Plots
6. Collision Theory, Activation Energy and Activated Complex
Collision theory states that reacting particles must collide with sufficient energy and suitable orientation for reaction to occur.
- Effective collision: collision that leads to product formation.
- Activation energy, Ea: minimum energy barrier that reacting particles must overcome.
- Activated complex: unstable, high-energy arrangement near the top of the energy barrier.
Fig. 5 – Energy Profile and Activation Energy
7. Temperature Effect and Arrhenius Equation
As temperature increases, a larger fraction of molecules has energy equal to or greater than Ea, so k generally increases.
Fig. 6 – Maxwell–Boltzmann Concept and Activation Energy
8. Catalysis
| Type | Description | Example Concept |
|---|---|---|
| Homogeneous | Catalyst and reactants are in the same phase. | Acid-catalyzed reactions in solution. |
| Heterogeneous | Catalyst and reactants are in different phases. | Gas reaction on a solid metal surface. |
| Enzyme catalysis | Biological catalysts accelerate biochemical reactions with high specificity. | Enzyme + substrate → products. |
Fig. 7 – Catalyst Lowers Activation Energy
9. Worked Numerical Patterns
10. Quick Revision & Exam Points
- Define chemical kinetics, average rate and instantaneous rate.
- Explain rate law and rate constant.
- Differentiate order and molecularity.
- Derive integrated rate equation for zero-order reaction.
- Derive integrated rate equation for first-order reaction.
- Derive half-life expressions for zero and first order.
- Explain collision theory, activation energy and activated complex.
- State and use Arrhenius equation.
- Explain effects of concentration, temperature and catalyst on rate.
- Differentiate homogeneous, heterogeneous and enzyme catalysis.
- Rate law is determined experimentally.
- Order = sum of rate-law exponents.
- Zero-order [A] vs t is linear.
- First-order ln[A] vs t is linear.
- First-order t₁/₂ = 0.693/k.
- Effective collision needs sufficient energy and orientation.
- Eₐ is the reaction energy barrier.
- Higher T generally raises k.
- Catalyst lowers activation energy.
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