Chemical Kinetics – Chemistry Quick Revision Notes 2025

1. Introduction to Chemical Kinetics

This is a branch of physical chemistry that studies the speed or rate at which chemical reactions take place and the factors that can change this speed. It also explains how a reaction happens; the step-by-step path is called the mechanism.

Importance of Chemical Kinetics (PGT interview points)

• It helps industries design faster and safer chemical processes.
• It predicts reaction time, remaining concentration of reactant, product concentration
• It explains: How does food spoil? And how do preservatives work?
• Helps scientists control explosions & combustion reactions.
• Supports environmental chemistry (e.g., rate of ozone depletion in the stratosphere, formation of photochemical smog)

2. Basic Terms

3. Types of Reaction Rates

Average Rate

$$\text{Average Rate} = -\frac{\Delta [\text{Reactant}]}{\Delta t} = \frac{\Delta [\text{Product}]}{\Delta t}$$

Used when data is collected after a time interval.

Instantaneous Rate

$$\text{Instantaneous Rate} = \frac{d[\text{Product}]}{dt}$$

Initial Rate

Rate measured just when the reaction starts. Useful for determining rate law experimentally.

4. Rate Law and Rate Constant

The rate law expresses rate as a function of concentration:$$\text{Rate} = k[A]^m [B]^n$$

• m, n are experimentally determined, not from balanced equation.
• The overall order = m + n.

Key Points (PGT level):

• Rate constant has units that depend on order.
• Rate law tells how concentration affects rate, but doesn’t describe mechanism directly.

5. Order of Reaction

5.1 Zero Order

$$\text{Rate} = k$$

• Rate is independent of concentration.
• Example: Decomposition of NH₃ on Pt catalyst.

5.2 First Order

$$\text{Rate} = k[A]$$

Example: Radioactive decay, H₂O₂ decomposition, hydrolysis of esters.

5.3 Second Order

$$\text{Rate} = k[A]^2 \quad \text{or} \quad k[A][B]$$

Example: Reaction between NO and O₂.

5.4 Pseudo First Order

When one reactant is taken in large excess → effective first order.
Example: Hydrolysis of ester in presence of excess water.

6. Half-Life (t_½) and Integrated Rate Equations

Interview Tip: First-order reactions have constant half-life, a very common question.

7. Molecularity vs Order: Comparison Table

8. Collision Theory

To react, molecules must:

1. Collide
2. Collide with enough energy (≥ Ea)
3. Collide with proper orientation

The rate of reaction increases when collisions become:

• More frequent
• More energetic
• Better oriented

Limitations

• Works well for gas-phase, simple reactions only.
• Fails for complex reactions because orientation factor is difficult to calculate.
• Does not predict reaction pathways.

9. Transition State Theory

Also called Activated Complex Theory.

Main Concepts

• During reaction, reactants form a temporary, unstable structure called activated complex or transition state.
• This structure has maximum energy.
• From here, the reaction can move towards products or revert back to reactants.

Energy Profile Diagram

Reactants → Activated Complex (peak) → Products
Height of peak = Activation Energy

Why TST is considered better than Collision Theory?

• Includes entropy and molecular arrangement effects.
• Explains orientation in more realistic terms.
• Efficient for complex mechanisms.

10. Arrhenius Equation — Temperature Dependence of Rate

$$k = Ae^{-\frac{E_a}{RT}}$$

Where:

• A = frequency factor
• Ea = activation energy
• R = gas constant
• T = temperature in Kelvin

Arrhenius Plot

A plot of $\ln k$ vs1/T gives a straight line.
Slope = $-E_a/R$

Important Observations

• Rate roughly doubles for every 10°C rise (common interview point). [WBSLST IX-X 2016]
• Higher Ea → slower reaction.
• Catalysts lower Ea → faster reaction.

11. Catalysis

Types of Catalysis

1. Homogeneous Catalysis:
   Catalyst and reactants both are in same phase.
   Example: Hydrolysis of ester by acide/base
2. Heterogeneous Catalysis:
   Catalyst in different phase (usually solid surface).
   Example: Haber process using Fe(s).
3. Enzyme Catalysis:
   Highly specific biological catalysts.
   Example: Amylase converting starch → sugar.
4. Auto-catalysis:
   Product acts as a catalyst.
   Example: Mn²⁺ in KMnO₄ decomposition.

PG Interview-Level Concepts

• Adsorption theory for heterogeneous catalysis.
• Lock-and-key theory & induced-fit theory for enzyme catalysis.
• Catalyst poisons (e.g., lead poisoning Pt catalyst).
• Promoters (e.g., Mo in Haber process).

12. Reaction Mechanism

A mechanism explains how a reaction proceeds step-by-step.

Features of a Good Mechanism

• Sum of elementary steps must equal overall reaction.
• Mechanism must match experimentally determined rate law.
• Must identify the rate-determining step (slowest step).

Example

Reaction:$$2NO + O_2 \rightarrow 2NO_2$$

Mechanism:

1. $NO + O_2 \rightarrow NO_3$(slow)
2. $NO_3 + NO \rightarrow 2NO_2$

Since rate depends on slow step:$$\text{Rate} = k[NO]^2[O_2]$$

13. Factors Affecting Reaction Rate

1. Nature of Reactants

• Ionic reactions are fast (no bond breaking).
• Covalent reactions slow (bond breaking required).

2. Temperature

• Higher temperature = higher kinetic energy.
• More molecules cross activation barrier.

3. Concentration / Pressure

• Higher concentration → more collisions.
• Pressure increases rate in gases.

4. Catalyst

• Lowers activation energy.
• Provides alternative pathway.

5. Surface Area

More area → more collisions; important in heterogeneous catalysis.

6. Light (Photochemical Reactions)

Example: Photosynthesis, halogenation.

14. Steady-State Approximation

Used in complex mechanisms when intermediate concentration remains almost constant.

Assumption:$$\frac{d[\text{Intermediate}]}{dt} \approx 0$$

This simplifies rate expressions and is commonly used in PG questions.

15. Chain Reactions

Chain reactions involve:

1. Initiation
2. Propagation
3. Termination

Example:
Chlorination of methane.

Why chain reactions are fast?
Because each radical can create many more radicals → explosive growth.

16. Photochemical Kinetics

Light energy can start reactions by supplying activation energy.

Example:$$Cl_2 \xrightarrow{hv} 2Cl^\bullet$$

Quantum Yield:$$\phi = \frac{\text{No. of molecules reacted}}{\text{No. of photons absorbed}}$$

High quantum yield is characteristic of chain reactions.

17. Free Energy and Reaction Rate

Even if a reaction is thermodynamically feasible (ΔG < 0), it may still be slow due to high activation energy.

Example:
Diamond → Graphite
(ΔG negative but rate extremely slow.)

This is a favourite interviewer question:
“Why doesn’t diamond turn into graphite quickly although it is thermodynamically unstable?”
Answer: Because the activation energy barrier is extremely high.

18. Kinetics in Real-Life Systems

• 1. Food Spoilage: Reactions responsible for spoilage speed up at higher temperatures → refrigerators slow them down.
• 2. Medicine Breakdown: Drugs degrade by first-order kinetics → determines shelf life.
• 3. Environmental Chemistry: Ozone depletion reactions follow complex chain mechanisms.
• 4. Combustion: Combustion reactions are fast because activation energies are crossed instantly at high temperature.

19. Typical PGT / PG Interview Questions with Short Answers

Q1: Why is molecularity never zero or fractional?
Answer: Because molecularity refers to the number of molecules participating in a single elementary step; you cannot have a fraction of a molecule participating.

Q2: Can order be zero? Why?
Answer: Yes. If rate does not depend on concentration, it becomes zero order — often due to surface saturation or enzyme saturation.

Q3: Why does a catalyst not change equilibrium constant?
Answer: Because it speeds up both forward and backward reactions equally.

Q4: What is the significance of activation energy?
Answer: It determines how easily reactants convert to products. Lower Ea → higher rate.

Q5: What is the difference between collision theory and transition state theory?
Answer: Collision theory depends only on collision frequency and energy. Transition state theory considers formation of activated complex and includes entropy factors; more accurate for complex reactions.

Q6. Give a Humanized Explanation of Activation Energy.
Answer:

• Imagine you and your friends trying to jump over a low wall. If the wall is low, most of you will easily cross it. But if the wall is high, only a few strong friends can jump across. This wall is like activation energy.
• Adding a catalyst is like placing a small box beside the wall so everyone can climb easily. This is exactly how catalysts work: they lower the height of the barrier.
• When temperature increases, it’s like giving everyone extra energy so more people can cross the wall. This explains why reactions become faster at higher temperatures.