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Class 12 Chemistry Notes Chapter 4 (Chemical kinetics) – Chemistry-I Book

Philoid

12 Jul 2025 — 8 min read

Chemistry-I
Alright, let's get straight into Chapter 4: Chemical Kinetics. This is a crucial chapter, focusing on the speed or rate at which chemical reactions occur and the factors that influence these rates. Understanding this is vital not just for your board exams but also for various competitive government exams where chemistry is a component.

Chapter 4: Chemical Kinetics - Detailed Notes

1. Introduction

Chemical Kinetics: The branch of chemistry dealing with the study of reaction rates, the factors affecting these rates (like concentration, temperature, pressure, catalyst), and the mechanism by which reactions proceed.
Importance: Helps predict how quickly a reaction reaches equilibrium, optimize reaction conditions in industry, and understand biological processes.

2. Rate of a Chemical Reaction

Definition: The change in concentration of any one of the reactants or products per unit time.
Average Rate: The rate measured over a finite, measurable interval of time.
- For a reaction: R → P
- Average Rate = - Δ[R] / Δt = + Δ[P] / Δt
  - Δ[R] = Change in concentration of reactant R
  - Δ[P] = Change in concentration of product P
  - Δt = Time interval
  - Negative sign for reactants indicates decrease in concentration.
  - Positive sign for products indicates increase in concentration.
- For a general reaction: aA + bB → cC + dD
- Rate = - (1/a) Δ[A]/Δt = - (1/b) Δ[B]/Δt = + (1/c) Δ[C]/Δt = + (1/d) Δ[D]/Δt
Instantaneous Rate: The rate of reaction at a specific instant in time. It is the slope of the tangent drawn to the concentration vs. time curve at that instant.
- Mathematically: Rate_inst = - (1/a) d[A]/dt = - (1/b) d[B]/dt = + (1/c) d[C]/dt = + (1/d) d[D]/dt
- Where d[ ]/dt represents the derivative (instantaneous change) of concentration with respect to time.
Units of Rate: Concentration / Time (e.g., mol L⁻¹ s⁻¹, mol L⁻¹ min⁻¹, atm s⁻¹ for gaseous reactions).

3. Factors Influencing Rate of Reaction

Concentration of Reactants: Generally, the rate increases with increasing concentration of reactants (more collisions).
Temperature: Reaction rates almost always increase significantly with increasing temperature (molecules have higher kinetic energy, more effective collisions). A common rule of thumb is the rate doubles or triples for every 10°C rise.
Nature of Reactants and Products: Physical state (gas > liquid > solid), bond strengths, and complexity of molecules affect the rate. Reactions involving ionic species are usually very fast.
Presence of a Catalyst: A catalyst increases the rate of reaction without being consumed itself, by providing an alternative reaction pathway with lower activation energy.
Surface Area of Reactants: For heterogeneous reactions (reactants in different phases), increasing the surface area of the solid reactant increases the rate (more contact points).
Exposure to Radiation: Some reactions (photochemical reactions) are initiated or accelerated by absorbing light of a suitable wavelength (e.g., reaction of H₂ and Cl₂ in sunlight).

4. Rate Law and Rate Constant

Rate Law (or Rate Equation): An experimentally determined expression that relates the rate of a reaction to the molar concentrations of the reactants, each raised to some power.
- For a general reaction: aA + bB → Products
- Rate Law: Rate = k [A]^x [B]^y
  - k = Rate Constant (or specific reaction rate)
  - x = Order of reaction with respect to A
  - y = Order of reaction with respect to B
  - x and y are determined experimentally and may or may not be equal to the stoichiometric coefficients a and b.
Order of Reaction: The sum of the powers to which the concentration terms are raised in the experimentally determined rate law.
- Overall Order = x + y
- Order can be zero, fractional, or integer.
Rate Constant (k):
- It is the rate of reaction when the concentration of each reactant is unity (1 mol L⁻¹).
- It is characteristic of a specific reaction at a given temperature.
- It depends on temperature but is independent of reactant concentrations.
- Units of k: Depend on the overall order (n) of the reaction.
  - General Unit: (Concentration)^(1-n) (Time)⁻¹ or (mol L⁻¹)⁽¹⁻ⁿ⁾ s⁻¹
  - Zero Order (n=0): mol L⁻¹ s⁻¹ (Same as rate)
  - First Order (n=1): s⁻¹
  - Second Order (n=2): L mol⁻¹ s⁻¹
Molecularity: (Applies only to elementary reactions - reactions occurring in a single step)
- The number of reacting species (atoms, ions, or molecules) that must collide simultaneously to bring about the chemical reaction in a single step.
- It is a theoretical concept.
- It must be a whole number (1, 2, or 3). It cannot be zero or fractional.
- For complex reactions (multi-step), molecularity has no meaning for the overall reaction. The slowest step (rate-determining step) determines the overall rate.

Difference between Order and Molecularity:

Feature	Order of Reaction	Molecularity
Definition	Sum of powers in rate law	No. of species in elementary step
Determination	Experimental	Theoretical (based on mechanism)
Values	0, fraction, integer	Integer (1, 2, 3 usually)
Applicability	Overall reaction / Elementary steps	Only Elementary steps
Meaning	Relates rate to concentration	Describes collision mechanism

5. Integrated Rate Equations

These equations relate concentration to time directly.
Zero-Order Reaction: Rate is independent of reactant concentration. Rate = k[R]⁰ = k
- Integrated Rate Law: [R] = -kt + [R]₀ (where [R]₀ is initial concentration, [R] is concentration at time t)
- Plot of [R] vs t is a straight line with slope -k and intercept [R]₀.
- Half-life (t½): Time taken for concentration to reduce to half its initial value.
- t½ = [R]₀ / 2k (Half-life is directly proportional to initial concentration)
- Examples: Decomposition of gases on metal surfaces at high pressure (e.g., NH₃ on hot Pt).
First-Order Reaction: Rate is directly proportional to the concentration of one reactant. Rate = k[R]¹
- Integrated Rate Law:
  - ln[R] = -kt + ln[R]₀
  - or ln([R]/[R]₀) = -kt
  - or [R] = [R]₀ * e^(-kt)
  - or k = (2.303 / t) * log([R]₀ / [R]) (Most commonly used form for calculations)
- Plot of ln[R] vs t is a straight line with slope -k and intercept ln[R]₀.
- Plot of log[R] vs t is a straight line with slope -k / 2.303 and intercept log[R]₀.
- Half-life (t½):
- t½ = ln(2) / k = 0.693 / k (Half-life is independent of initial concentration - very important!)
- Examples: Radioactive decay, decomposition of N₂O₅, hydrolysis of sucrose (in acidic medium).
Pseudo-First Order Reaction: Reactions which are bimolecular but behave as first order. This happens when one reactant is present in large excess, and its concentration remains virtually constant during the reaction.
- Example: Acid hydrolysis of ethyl acetate: CH₃COOC₂H₅ + H₂O --(H⁺)--> CH₃COOH + C₂H₅OH
- Rate = k' [CH₃COOC₂H₅] [H₂O]. If [H₂O] is very large (solvent), it's constant.
- Rate = k [CH₃COOC₂H₅], where k = k' [H₂O]. The reaction behaves as first order.

6. Temperature Dependence of Reaction Rate & Arrhenius Equation

Activation Energy (Ea): The minimum extra energy that reacting molecules must possess (over their average energy) to overcome the energy barrier and form products upon collision.
Activated Complex (Transition State): An unstable, high-energy intermediate formed during the conversion of reactants to products.
Arrhenius Equation: Quantitatively relates the rate constant (k) to temperature (T) and activation energy (Ea).
- k = A * e^(-Ea / RT)
  - k = Rate constant
  - A = Arrhenius factor or Pre-exponential factor or Frequency factor (related to collision frequency and orientation)
  - Ea = Activation Energy (usually in J mol⁻¹ or kJ mol⁻¹)
  - R = Gas Constant (8.314 J K⁻¹ mol⁻¹)
  - T = Absolute Temperature (in Kelvin)
- Taking natural logarithm: ln k = ln A - Ea / RT
- Converting to base 10 logarithm: log k = log A - Ea / (2.303 RT)
- Plot of ln k vs 1/T gives a straight line with slope -Ea / R and intercept ln A.
- Plot of log k vs 1/T gives a straight line with slope -Ea / (2.303 R) and intercept log A.
Calculating Ea from rate constants at two temperatures (T₁ and T₂):
- log(k₂ / k₁) = (Ea / 2.303 R) * [(T₂ - T₁) / (T₁ T₂)]

7. Collision Theory of Chemical Reactions

Provides a qualitative explanation for reaction rates.
Basic Postulate: Reactions occur when reactant molecules collide.
Effective Collisions: Not all collisions lead to product formation. For a collision to be effective, molecules must possess:
- Sufficient Energy: Kinetic energy equal to or greater than the activation energy (Ea) – called the energy barrier.
- Proper Orientation: Molecules must collide in a specific orientation that allows bonds to break and form – called the orientation barrier.
Rate Expression based on Collision Theory:
- Rate = P * Z_AB * e^(-Ea / RT)
  - Z_AB = Collision frequency (number of collisions per second per unit volume) between A and B.
  - e^(-Ea / RT) = Fraction of molecules having energy ≥ Ea.
  - P = Probability or Steric factor (accounts for proper orientation, typically < 1).
- Comparing with Arrhenius equation, A = P * Z_AB.
Limitations: Mainly applicable to simple gaseous reactions; doesn't fully account for complex molecules or solvent effects.

8. Effect of Catalyst

Definition: A substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change.
Mechanism: Provides an alternative reaction pathway or mechanism with a lower activation energy (Ea). It does not alter the Gibbs energy change (ΔG) or equilibrium constant (K) of the reaction.
Characteristics:
- Required in small quantities.
- Highly specific in action.
- Does not initiate a reaction, only accelerates it.
- Does not change the position of equilibrium (accelerates both forward and backward reactions equally).
- Activity can be affected by temperature, pH, promoters (enhance activity), and poisons (decrease activity).
Energy Profile Diagram: A catalyst lowers the peak of the energy barrier on the reaction coordinate diagram.

Multiple Choice Questions (MCQs)

The unit of rate constant for a zero-order reaction is:
a) s⁻¹
b) L mol⁻¹ s⁻¹
c) mol L⁻¹ s⁻¹
d) L² mol⁻² s⁻¹
For a first-order reaction A → Products, the half-life (t½) is 10 minutes. What percentage of A will be left after 40 minutes?
a) 50%
b) 25%
c) 12.5%
d) 6.25%
The rate law for a reaction is found to be Rate = k[A]²[B]. If the concentration of A is doubled and B is halved, the rate of reaction will:
a) Increase by 4 times
b) Increase by 2 times
c) Remain the same
d) Decrease by 2 times
Which of the following statements is incorrect about the order of a reaction?
a) It is determined experimentally.
b) It can be zero or fractional.
c) It is always equal to the sum of stoichiometric coefficients in the balanced equation.
d) It represents the sum of powers of concentration terms in the rate law.
Activation energy (Ea) of a chemical reaction can be determined by:
a) Evaluating rate constant at standard temperature.
b) Evaluating velocities of molecules.
c) Evaluating rate constants at two different temperatures.
d) Evaluating the number of collisions per unit time.
The role of a catalyst is to change the:
a) Gibbs energy of reaction
b) Enthalpy of reaction
c) Activation energy of reaction
d) Equilibrium constant
Molecularity of a reaction:
a) Is always determined experimentally.
b) Can be zero or fractional.
c) Is the number of molecules involved in the rate-determining step of a complex reaction.
d) Is the number of reacting species colliding simultaneously in an elementary reaction step.
For a reaction 2A + B → C, the rate law is given by Rate = k[A][B]. If the concentration of B is kept constant and the concentration of A is doubled, the rate will:
a) Double
b) Halve
c) Increase by four times
d) Remain unchanged
The hydrolysis of ethyl acetate in acidic medium (CH₃COOC₂H₅ + H₂O --(H⁺)--> CH₃COOH + C₂H₅OH) is an example of:
a) Zero-order reaction
b) Second-order reaction
c) Pseudo-first order reaction
d) Third-order reaction
According to the collision theory, the rate of reaction increases with temperature primarily because of:
a) Increase in the number of collisions.
b) Increase in the activation energy.
c) Increase in the fraction of molecules possessing energy equal to or greater than the activation energy.
d) Decrease in the steric factor.

Answers to MCQs:

c) mol L⁻¹ s⁻¹
d) 6.25% (After 1 t½=10min, 50% left; after 2 t½=20min, 25% left; after 3 t½=30min, 12.5% left; after 4 t½=40min, 6.25% left)
b) Increase by 2 times (Rate' = k[2A]²[B/2] = k * 4[A]² * (1/2)[B] = 2 * k[A]²[B] = 2 * Rate)
c) It is always equal to the sum of stoichiometric coefficients in the balanced equation. (This is only true if the reaction is elementary).
c) Evaluating rate constants at two different temperatures (using the Arrhenius equation).
c) Activation energy of reaction
d) Is the number of reacting species colliding simultaneously in an elementary reaction step.
a) Double (Rate' = k[2A][B] = 2 * k[A][B] = 2 * Rate)
c) Pseudo-first order reaction (Water is in large excess).
c) Increase in the fraction of molecules possessing energy equal to or greater than the activation energy (represented by the e^(-Ea/RT) term).

Study these notes thoroughly. Pay close attention to the definitions, formulas (especially for integrated rate laws, half-lives, and Arrhenius equation), units, and the differences between concepts like order and molecularity. Good luck with your preparation!

Class 12 Chemistry Notes Chapter 4 (Chemical kinetics) – Chemistry-I Book

Philoid

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