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Chemical Kinetics – Class 12 Chemistry Chapter 3 Notes & NCERT PDF

Class 12ChemistryChapter 3NCERT book: Chemistry Part I - Class 12

Notes and a simple summary of Chapter 3 of the NCERT Class 12 Chemistry book, with the official chapter PDF, flashcards, an MCQ quiz and an AI tutor for your doubts.

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Chapter 3: Chemical Kinetics

Introduction

Thermodynamics tells us whether a reaction is feasible; chemical kinetics tells us how fast it goes and by what path (mechanism). Some reactions such as ionic reactions are very fast, rusting of iron is slow, and inversion of cane sugar is moderate. This unit covers average and instantaneous rates, factors affecting rate, rate laws, order and molecularity, integrated rate equations, half-life, the effect of temperature (Arrhenius equation), catalysts and collision theory.

Key Concepts

3.1 Rate of a Chemical Reaction

  • Rate = change in concentration of a reactant or product per unit time; units mol L⁻¹ s⁻¹ (for gases, atm s⁻¹).
  • Average rate = −Δ[R]/Δt = +Δ[P]/Δt over a time interval; instantaneous rate = −d[R]/dt, the slope of the tangent to the concentration–time curve.
  • With unequal stoichiometry, divide by coefficients: for 2HI → H₂ + I₂, rate = −½ Δ[HI]/Δt = Δ[H₂]/Δt = Δ[I₂]/Δt.

    • 3.2 Factors Influencing Rate

    Concentration of reactants (pressure for gases), temperature and catalyst.

  • Rate law: experimentally determined, e.g. for 2NO + O₂ → 2NO₂, Rate = k[NO]²[O₂]. It cannot be predicted from the balanced equation.
  • Order = sum of powers of concentration terms in the rate law; can be 0, 1, 2, 3 or even a fraction.
  • Units of k: zero order mol L⁻¹ s⁻¹; first order s⁻¹; second order mol⁻¹ L s⁻¹.
  • Elementary reactions occur in one step; complex reactions occur as a sequence of steps. The slowest step is rate determining.
  • Molecularity = number of reacting species taking part in an elementary reaction that must collide simultaneously: unimolecular (NH₄NO₂ → N₂ + 2H₂O), bimolecular (2HI → H₂ + I₂), termolecular (rare). Molecularity is always a whole number, is never zero, and has no meaning for complex reactions; order applies to both.

    • 3.3 Integrated Rate Equations

    Zero order: [R] = [R]₀ − kt; k = ([R]₀ − [R])/t; t½ = [R]₀/2k (proportional to initial concentration).

    Examples: decomposition of gaseous ammonia on a hot platinum surface at 1130 K (at high pressure the metal surface is saturated with gas molecules), and thermal decomposition of HI on a gold surface.

    First order: ln[R] = ln[R]₀ − kt, or

    k = (2.303/t) log([R]₀/[R]); t½ = 0.693/k (independent of initial concentration)

    Examples: hydrogenation of ethene, and all natural and artificial radioactive decay of unstable nuclei. For a gas-phase first-order reaction A(g) → B(g) + C(g), k = (2.303/t) log[pᵢ/(2pᵢ − p_t)].

    Pseudo first-order reactions: reactions of higher order that behave as first order because one reactant (usually water) is in large excess — e.g. acid hydrolysis of ethyl acetate (0.01 mol ester with 10 mol water), and inversion of cane sugar (C₁₂H₂₂O₁₁ + H₂O → glucose + fructose), Rate = k[C₁₂H₂₂O₁₁].

    3.4 Temperature Dependence — Arrhenius Equation

    For a chemical reaction, a 10° rise in temperature nearly doubles the rate constant.

    k = A e^(−Ea/RT); ln k = −Ea/RT + ln A; log(k₂/k₁) = (Ea/2.303R)(1/T₁ − 1/T₂)
  • A = Arrhenius (frequency / pre-exponential) factor; Ea = activation energy — the energy needed to form the unstable activated complex, e.g. in H₂ + I₂ → 2HI.
  • A plot of ln k vs 1/T is a straight line with slope −Ea/R and intercept ln A.
  • The Maxwell–Boltzmann distribution (Ludwig Boltzmann and James Clerk Maxwell) shows that raising temperature by 10° roughly doubles the fraction of molecules with energy ≥ Ea, so the rate roughly doubles.

    • 3.5 Effect of Catalyst

    A catalyst increases the rate without undergoing any permanent chemical change (e.g. MnO₂ in 2KClO₃ → 2KCl + 3O₂). It provides an alternate pathway with lower activation energy by forming an intermediate complex. It does not alter ΔG, the equilibrium constant or the position of equilibrium — it helps reach equilibrium faster by catalysing forward and backward reactions equally. A small amount of catalyst can catalyse a large amount of reactants. Inhibitors decrease the rate.

    3.6 Collision Theory

    Developed by Max Trautz and William Lewis (1916–18). Reactant molecules are hard spheres; reaction occurs when they collide. The number of collisions per second per unit volume is the collision frequency (Z). Only effective collisions — with energy ≥ threshold energy and proper orientation — form products:

    Rate = P Z_AB e^(−Ea/RT)

    where P is the probability or steric factor accounting for orientation (e.g. formation of methanol from bromoethane depends on how the reacting molecules are oriented).

    Worked Examples

  • First order k (Ex. 3.5): N₂O₅ in CCl₄: 1.24 → 0.20 mol L⁻¹ in 60 min. k = (2.303/60) log(1.24/0.20) = 0.0304 min⁻¹.
  • Gas phase first order (Ex. 3.6): N₂O₅ at 318 K, pressure change gives k ≈ 4.92 × 10⁻⁴ s⁻¹ (the textbook prints 4.98 × 10⁻⁴ s⁻¹; direct calculation of (2.303/100) log(0.5/0.476) gives 4.92 × 10⁻⁴).
  • Half-life (Ex. 3.7): k = 5.5 × 10⁻¹⁴ s⁻¹ → t½ = 0.693/k = 1.26 × 10¹³ s.
  • 99.9% completion (Ex. 3.8): t = (2.303/k) log 1000 = 6.909/k, while t½ = 0.693/k, so t(99.9%) ≈ 10 t½.
  • Activation energy (Ex. 3.9): k = 0.02 s⁻¹ at 500 K and 0.07 s⁻¹ at 700 K → Ea ≈ 18.23 kJ mol⁻¹, A ≈ 1.61 s⁻¹.
  • Arrhenius prediction (Ex. 3.10): k = 1.60 × 10⁻⁵ s⁻¹ at 600 K, Ea = 209 kJ mol⁻¹ → k at 700 K ≈ 6.36 × 10⁻³ s⁻¹.

    • Formula Sheet

  • Zero order: k = ([R]₀ − [R])/t; t½ = [R]₀/2k
  • First order: k = (2.303/t) log([R]₀/[R]); t½ = 0.693/k
  • Gas phase first order: k = (2.303/t) log[pᵢ/(2pᵢ − p_t)]
  • Arrhenius: k = A e^(−Ea/RT); log(k₂/k₁) = (Ea/2.303R)[(T₂ − T₁)/(T₁T₂)]
  • Collision theory: Rate = P Z_AB e^(−Ea/RT)
  • Units of k for order n: (mol L⁻¹)^(1−n) s⁻¹

    • Board Exam Focus

  • Numericals: first-order k and t½; time for x% completion; Ea from two rate constants; k at a new temperature.
  • Distinguish order vs molecularity (at least four points); average vs instantaneous rate.
  • Graphs: [R] vs t (zero order — straight line, slope −k), ln[R] vs t (first order — slope −k), ln k vs 1/T (slope −Ea/R).
  • Reasoning: Why does a catalyst not change equilibrium? Why is inversion of cane sugar pseudo first order? Why do all collisions not lead to products?
  • Answer outline — effect of temperature: 10° rule, Arrhenius equation, Maxwell–Boltzmann curve showing more molecules above Ea at higher T.

    • Key Terms

  • Rate of reaction (अभिक्रिया वेग)
  • Rate constant (वेग स्थिरांक)
  • Order of reaction (अभिक्रिया की कोटि)
  • Molecularity (आण्विकता)
  • Half-life (अर्धायु)
  • Activation energy (सक्रियण ऊर्जा)
  • Activated complex (सक्रियित संकुल)
  • Catalyst (उत्प्रेरक)
  • Collision frequency (संघट्ट आवृत्ति)
  • Pseudo first order reaction (छद्म प्रथम कोटि अभिक्रिया)

    • Common Mistakes

  • Writing the rate law from the balanced equation — it must be found experimentally.
  • Confusing order with molecularity; molecularity can never be zero or fractional.
  • Using t½ = 0.693/k for zero-order reactions (it is [R]₀/2k).
  • Mixing ln and log: remember the factor 2.303.
  • Believing a catalyst shifts equilibrium or changes ΔG.
  • Getting the sign wrong in log(k₂/k₁) — keep T₂ > T₁ with (T₂ − T₁) positive.

    • 💡 Key Learning Points

    • ✓Define average and instantaneous rates and express rate using stoichiometric coefficients.
    • ✓Distinguish order and molecularity and deduce units of rate constants.
    • ✓Use integrated rate equations and half-life for zero and first order reactions.
    • ✓Apply the Arrhenius equation to find activation energy and rate constants at other temperatures.
    • ✓Explain the role of catalysts and the collision theory of reaction rates.

    👨‍🏫 Teaching Tips

    • →Compare fast (precipitation), moderate (sugar inversion) and slow (rusting) reactions to introduce rates.
    • →Draw a concentration–time curve and let students find average and instantaneous rates using tangents.
    • →Use a table of initial-rate data to let students deduce the rate law themselves.
    • →Model first-order decay with dice or coins removed each round to visualise constant half-life.
    • →Sketch the Maxwell–Boltzmann curve at two temperatures to show why rate doubles for 10°.
    • →Draw energy profiles with and without catalyst to show lowered activation energy.
    • →Drill Arrhenius numericals, insisting on T₂ > T₁ and consistent units of R.

    📋 Assessment Questions

    1. Determine order and rate law from initial-rate data.
    2. Calculate k and t½ for first-order reactions, including time for x% completion.
    3. Derive or use the integrated rate equation for zero-order reactions.
    4. Calculate Ea from rate constants at two temperatures.
    5. Explain effective collisions, steric factor and the role of a catalyst.

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    All chapters in Chemistry Part I - Class 12

    1. 1. Solutions
    2. 2. Electrochemistry
    3. 3. Chemical Kinetics
    4. 4. The d- and f-Block Elements
    5. 5. Coordination Compounds