Class 12 Chemistry Chapter 3: Chemical Kinetics

Chapter 3: Chemical Kinetics (रासायनिक बलगतिकी) is a high-weightage chapter in the MP Board Class 12 Chemistry syllabus, typically contributing 10–14 marks in the board exam. This chapter explores the rates of chemical reactions — how fast reactants turn into products and what factors influence that speed. From the concept of reaction rate to integrated rate laws, activation energy, and collision theory, these comprehensive notes cover every key topic with clear formulas, MP Board–focused exam tips, and previous year pattern analysis for the 2027 board exam.

⏱️ 1. Rate of a Chemical Reaction

Chemical kinetics is the branch of physical chemistry that deals with the study of rates of chemical reactions and the factors affecting them. The rate of a reaction is defined as the change in concentration of a reactant or product per unit time.

Average Rate vs Instantaneous Rate

Average rate (r_avg): Change in concentration over a finite time interval.

r_avg = −Δ[R] / Δt = +Δ[P] / Δt

Instantaneous rate (r_inst): Rate at a particular instant, obtained as the slope of the concentration vs time graph at that point.

r_inst = −d[R]/dt = +d[P]/dt

Units of Rate of Reaction

For a gaseous reaction: atm s⁻¹ or bar s⁻¹
For reactions in solution: mol L⁻¹ s⁻¹ (also written as M s⁻¹ or mol dm⁻³ s⁻¹)

🔬 2. Factors Affecting Rate of Reaction

📊 3. Rate Law and Order of Reaction

Rate Law Expression

For a general reaction: aA + bB → Products, the rate law is:
Rate = k [A]^x [B]^y

Where k = rate constant (specific rate), x = order w.r.t. A, y = order w.r.t. B, and (x + y) = overall order of the reaction.

Order of Reaction — Important Characteristics

• Order is experimentally determined — cannot be predicted from the balanced chemical equation
• Order can be zero, integer, fractional, or even negative
• Order is defined only for elementary reactions — for complex reactions, it applies to the rate-determining step
• The sum of powers (x + y) gives the overall order

Units of Rate Constant (k)

General formula: Unit of k = (mol L⁻¹)¹⁻ⁿ s⁻¹, where n = overall order

🧬 4. Molecularity of a Reaction

Molecularity is defined as the number of molecules (or atoms) that collide simultaneously to bring about a chemical reaction in an elementary step.

Types of Molecularity

• Unimolecular: Single molecule decomposes or rearranges (e.g., radioactive decay, NH₄NO₂ → N₂ + 2H₂O)
• Bimolecular: Two molecules collide (e.g., 2HI → H₂ + I₂, CH₃COOC₂H₅ + NaOH → CH₃COONa + C₂H₅OH)
• Termolecular: Three molecules collide simultaneously — very rare (e.g., 2NO + O₂ → 2NO₂)

📐 5. Integrated Rate Laws

Integrated rate laws express the concentration of reactants as a function of time. These are essential for solving numerical problems in the MP Board exam.

Zero Order Reaction

Rate is independent of reactant concentration: Rate = k[R]⁰ = k

Integrated equation: [R] = −kt + [R]₀

Graph: [R] vs t → straight line with negative slope = −k

Example: Photochemical reactions, decomposition of HI on gold surface, enzyme-catalysed reactions at high substrate concentration

First Order Reaction

Rate ∝ [R]¹: Rate = k[R]

Integrated equation: k = (2.303/t) × log₁₀([R]₀/[R])

Or: log₁₀([R]₀/[R]) = kt / 2.303

Graph: log[R] vs t → straight line with slope = −k/2.303

Examples: Radioactive decay, Hydrogenation of ethene, Inversion of cane sugar, Decomposition of N₂O₅

Second Order Reaction

Type 1 — Same reactants (2A → Products): Rate = k[A]²

Integrated equation: kt = 1/[A] − 1/[A]₀

Graph: 1/[A] vs t → straight line with slope = k

Type 2 — Different reactants (A + B → Products): If [A] = [B], same as Type 1. If [A] ≠ [B], integrated form is more complex.

⏲️ 6. Half-Life of a Reaction

Half-life (t₁/₂) is the time required for the concentration of a reactant to reduce to half of its initial value. It is a key parameter in chemical kinetics.

Half-Life Formulas

🧪 7. Pseudo First Order Reaction

A reaction that is biomolecular (second order) but behaves as first order because one reactant is present in large excess is called a pseudo first order reaction.

Classic Example: Hydrolysis of ethyl acetate (ester hydrolysis)
CH₃COOC₂H₅ + H₂O → CH₃COOH + C₂H₅OH

Here, water is in large excess, so its concentration remains practically constant. The rate law becomes:
Rate = k'[CH₃COOC₂H₅] (where k’ = k[H₂O])

Examples of Pseudo First Order Reactions

🌡️ 8. Temperature Dependence — Arrhenius Equation

The Arrhenius equation describes how the rate constant (k) varies with temperature:

k = Ae^−E_a/RT

Where:
k = rate constant
A = Arrhenius constant (frequency factor / pre-exponential factor)
E_a = Activation energy (J/mol or kJ/mol)
R = Universal gas constant (8.314 J mol⁻¹ K⁻¹)
T = Temperature in Kelvin

Logarithmic Form

Taking natural log: ln k = ln A − E_a/RT

Taking log₁₀: log₁₀ k = log₁₀ A − E_a/(2.303RT)

Graph: log k vs 1/T → straight line with slope = −E_a/(2.303R) and intercept = log A

Calculation of Activation Energy

If rate constants k₁ and k₂ are known at temperatures T₁ and T₂:

log₁₀(k₂/k₁) = E_a/(2.303R) × (1/T₁ − 1/T₂)

Effect of Catalyst on Activation Energy

A catalyst provides an alternative reaction pathway with lower activation energy. This means more molecules have sufficient energy to overcome the barrier, increasing the reaction rate. The catalyst itself is not consumed in the reaction.

💥 9. Collision Theory

According to the collision theory, a chemical reaction occurs when molecules collide with sufficient energy and proper orientation. For a reaction to occur, three conditions must be satisfied:

1. Collision: Reactant molecules must collide with each other
2. Energy: The collision must have energy ≥ activation energy (E_a)
3. Orientation: Molecules must collide with proper orientation for bond rearrangement

Important Concepts from Collision Theory

• Effective collisions — only collisions with sufficient energy and proper orientation lead to product formation
• Threshold energy — the minimum energy that colliding molecules must possess for reaction to occur
• Activation energy — E_a = Threshold energy − Average energy of reactants
• Steric factor (P) — accounts for orientation requirements; Rate = P × Z × e^−E_a/RT, where Z = collision frequency
• The Arrhenius constant A = P × Z (frequency factor corrected for orientation)

📋 10. Important Questions for MP Board 2027

✏️ Very Short Answer (1 Mark)

1. Define rate of a chemical reaction.
2. What is the unit of rate constant for a first-order reaction?
3. Define half-life of a reaction.
4. What is the molecularity of a reaction?
5. State the Arrhenius equation.

📝 Short Answer (2–3 Marks)

1. Distinguish between order of a reaction and molecularity.
2. Write the integrated rate equation for a first-order reaction. Derive the expression for its half-life period.
3. Explain with an example what is meant by a pseudo first order reaction.
4. What is the effect of temperature on the rate of a reaction? Explain with Arrhenius equation.
5. The half-life of a first-order reaction is 100 seconds. Calculate the rate constant.

📚 Long Answer (5–6 Marks)

1. Derive the integrated rate equation for a first-order reaction. Show that the half-life is independent of initial concentration. (Graph required)
2. Explain collision theory of chemical reactions. Discuss how temperature and catalyst affect reaction rates based on this theory.
3. What is activation energy? Derive the Arrhenius equation and explain how activation energy can be calculated graphically from log k vs 1/T plot.
4. A first order reaction is 50% complete in 40 minutes. Calculate: (a) the rate constant, (b) the time required for 90% completion, (c) the time required for 99% completion.

📖 More Resources: Download MP Board Class 12 Chemistry previous year papers and practice sets from mpboard.ai. For video explanations and chapter-wise PYQ discussions, visit our Class 12 Chemistry Course. Stay connected with mpboard.ai for the latest MP Board 2027 exam updates, study materials, and expert guidance.