Chemical Kinetics — chemistry Class 12 Notes (CBSE & HBSE)
Free NCERT chemistry notes for Chemical Kinetics (Class 12) on Siksha Sarovar, aligned to CBSE and Haryana Board (HBSE). This chapter is broken into 3 topics with clear explanations, formulas, solved examples and board-pattern practice — free to read, no sign-up required.
Board exam focus — Chemical Kinetics (CBSE & HBSE)
Rate of reaction, rate law, order and molecularity, integrated rate equations, half-life, temperature dependence, Arrhenius equation, activation energy and collision theory.
Rate of Reaction and Rate Law
Chemical Kinetics
Rate of Reaction
Rate of reaction = change in concentration / time For A → B: Rate = -d[A]/dt = +d[B]/dt
For aA + bB → cC + dD: Rate = -(1/a) d[A]/dt = -(1/b) d[B]/dt = +(1/c) d[C]/dt = +(1/d) d[D]/dt
Units: mol L^-1 s^-1 (or mol L^-1 min^-1 etc.)
Instantaneous rate: rate at a specific instant = -d[A]/dt (tangent to [A] vs t curve) Average rate: rate over a time interval = -Delta[A]/Delta t
Rate Law (Rate Expression)
Determined experimentally; not from stoichiometry. For aA + bB → products: Rate = k [A]^m [B]^n where k = rate constant, m = order w.r.t. A, n = order w.r.t. B.
Overall order = m + n
Order of Reaction
- Zero order: Rate = k (rate independent of concentration)
- First order: Rate = k[A]
- Second order: Rate = k[A]^2 or k[A][B]
- Pseudo-first order: Actually second order but [B] is very large/constant (e.g., hydrolysis of ethyl acetate in excess water)
Molecularity
Number of molecules colliding in elementary step.
- Unimolecular: 1 molecule (A → products)
- Bimolecular: 2 molecules (A + B → products)
- Termolecular: 3 molecules (rare)
Order vs Molecularity:
| Property | Order | Molecularity |
|---|---|---|
| Determined by | Experiment | Mechanism |
| Applies to | Overall reaction | Elementary step |
| Can be fractional | Yes | No (always integer) |
| Can be zero | Yes | No |
Units of Rate Constant (k)
| Order | Units |
|---|---|
| Zero | mol L^-1 s^-1 |
| First | s^-1 |
| Second | L mol^-1 s^-1 |
| nth order | (mol L^-1)^(1-n) s^-1 |
Factors Affecting Rate
- Concentration: Higher [A] = faster rate
- Temperature: Higher T = faster rate (usually doubles per 10 C)
- Catalyst: Lowers activation energy
- Surface area: More surface = more collisions
- Nature of reactants: Bond polarity, bond energy
Integrated Rate Equations and Half-Life
Integrated Rate Equations
Zero Order Reaction
Rate = k; d[A]/dt = -k Integrated: [A]t = [A]0 - kt Linear plot: [A] vs t (slope = -k) Half-life: t1/2 = [A]0 / 2k (depends on initial concentration) Example: Enzyme-catalyzed reactions at saturating substrate, photochemical reactions.
First Order Reaction
Rate = k[A]; -d[A]/dt = k[A] Integrated: ln[A]t = ln[A]0 - kt Or: [A]t = [A]0 e^(-kt) Or: log([A]0/[A]t) = kt/2.303
Linear plot: ln[A] vs t (slope = -k) or log[A] vs t (slope = -k/2.303)
Half-life: t1/2 = 0.693/k = ln2/k (independent of initial concentration)
Examples of first order reactions:
- Radioactive decay
- Decomposition of N2O5: 2N2O5 → 4NO2 + O2
- Decomposition of NH3 on Pt surface
- Inversion of sucrose (pseudo first order)
- Hydrolysis of esters in dilute acid (excess water)
Second Order Reaction
Rate = k[A]^2; -d[A]/dt = k[A]^2 Integrated: 1/[A]t = 1/[A]0 + kt Linear plot: 1/[A] vs t (slope = k) Half-life: t1/2 = 1/(k[A]0) (depends on initial concentration)
Identifying Reaction Order
| Method | Zero order | First order | Second order |
|---|---|---|---|
| Plot | [A] vs t (linear) | ln[A] vs t (linear) | 1/[A] vs t (linear) |
| t1/2 | depends on [A]0 | independent of [A]0 | depends on [A]0 |
Radioactive Decay (First Order)
N = N0 e^(-lambda t) lambda = decay constant = 0.693/t1/2 Activity A = lambda N
Carbon dating: Uses C-14 decay (t1/2 = 5730 years) to date organic matter.
Temperature Dependence and Arrhenius Equation
Temperature Dependence of Rate
Temperature Coefficient
Rate of reaction typically doubles for every 10 C rise in temperature. Temperature coefficient (mu) = k(T+10)/k(T) ≈ 2
Arrhenius Equation
k = A x e^(-Ea/RT)
where:
- k = rate constant
- A = pre-exponential factor (frequency factor, Arrhenius factor)
- Ea = activation energy (minimum energy needed for reaction)
- R = gas constant (8.314 J/mol/K)
- T = temperature (Kelvin)
Logarithmic form: ln k = ln A - Ea/RT log k = log A - Ea/(2.303 RT)
Linear form (plot of log k vs 1/T): log k = -(Ea/2.303 R) x (1/T) + log A Slope = -Ea/2.303R (negative slope) Intercept = log A
Calculating Ea from Two Temperature Data Points
log(k2/k1) = (Ea/2.303 R) x (1/T1 - 1/T2) = Ea(T2-T1) / (2.303 R T1 T2)
Activation Energy
The minimum energy above the average energy of reactants that molecules must possess to react.
Energy profile diagram: Reactants → transition state (highest energy) → products Activation energy = E(transition state) - E(reactants) DeltaH = E(products) - E(reactants)
Effect of catalyst: Catalyst provides an alternative pathway with lower activation energy. Same DeltaH but lower Ea → faster reaction.
Collision Theory
For a reaction to occur:
- Molecules must collide
- Collision energy >= activation energy (threshold energy)
- Collision geometry must be correct (orientation factor, probability factor p)
Rate = p x Z x e^(-Ea/RT) Z = collision frequency p = steric factor (usually << 1)
Transition State Theory: Reactants form an activated complex (transition state) at energy maximum. Activated complex then converts to products.
Frequently asked questions
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Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Chemical Kinetics.