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Chemical Kinetics — Class 12 Chemistry

The complete chapter guide: why some reactions finish in a blink and others take centuries — explained rate by rate, for CBSE, AHSEC, NEET, JEE Main and CUET.

CBSE Class 12 AHSEC NEET JEE Main CUET State PSC / Teaching Exams

Why This Chapter Matters

Think about two reactions happening right now, somewhere in the world. Somewhere, iron is rusting — a process so slow you'd need years to see real change. And somewhere else, a firework is exploding — a reaction finished in a fraction of a second. Both are chemical reactions. Both follow the same laws of thermodynamics. So why does one take years and the other take milliseconds?

That question — how fast, and why — is the entire subject of Chemical Kinetics. It's the branch of chemistry that studies the speed (rate) of chemical reactions and the factors that control that speed: concentration, temperature, catalysts, and the surface area of reactants.

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Kinetics isn't abstract — it's why food cooks faster in a pressure cooker (higher temperature and pressure speed up rate), why we refrigerate milk (lower temperature slows spoilage reactions), and why enzymes in your stomach digest food in hours rather than days.

NCERT NOTE This unit is officially Chapter 3 — Chemical Kinetics in the rationalised 2026–27 NCERT Class 12 Chemistry textbook. If you're using an older edition or state board sequence (including some AHSEC textbooks), you may see it numbered as Chapter 4 — the content is identical either way.

Chemical Kinetics carries roughly 6 marks in the CBSE board paper and shows up reliably in NEET and JEE Main — mostly through numerical questions on rate laws, half-life, and the Arrhenius equation. The good news? Once you're comfortable with the formulas, this chapter becomes one of the most scoring and mechanical topics in all of Class 12 Chemistry.

1. Rate of a Chemical Reaction

The rate of reaction is simply how fast reactants turn into products — measured as the change in concentration of a reactant or product per unit time.

FORMULA

Rate = ± Δ[Concentration] / Δt

Negative sign for reactants (their concentration decreases), positive sign for products (their concentration increases). This makes the rate value always positive.

For a reaction R → P:

FORMULA

Rate = −Δ[R]/Δt = +Δ[P]/Δt

Units: mol L⁻¹ s⁻¹ (or mol L⁻¹ min⁻¹, mol L⁻¹ h⁻¹ depending on the time unit used).

Average Rate vs Instantaneous Rate

Average Rate	Instantaneous Rate
Rate measured over a finite, measurable time interval (Δt)	Rate at one particular instant, as Δt → 0
Formula: Δ[R]/Δt	Formula: d[R]/dt (derivative, i.e. slope of the tangent on a concentration-time graph)
Gives an "overall" picture of the reaction	Gives the "true" rate at any given moment — this is what we usually mean by "the rate of reaction"

EXAM TIP If a numerical gives you a graph and asks for rate "at t = 20 s," draw a tangent at that point and calculate its slope — that's instantaneous rate. If it gives you two time points and two concentrations directly, that's average rate — no calculus needed.

Rate of Reaction in Terms of Stoichiometry

For a general reaction aA + bB → cC + dD, the rate expressed in terms of any single species must be divided by its stoichiometric coefficient so that the overall rate is the same regardless of which species you track:

FORMULA

Rate = −(1/a)(d[A]/dt) = −(1/b)(d[B]/dt) = +(1/c)(d[C]/dt) = +(1/d)(d[D]/dt)

2. Factors Affecting Rate of Reaction

• Concentration of reactants — higher concentration means more molecules per unit volume, so more frequent collisions, so faster rate.
• Temperature — raising temperature increases the kinetic energy of molecules, so more molecules cross the activation energy barrier. As a rule of thumb, rate roughly doubles for every 10°C rise (this is approximate, not a strict law).
• Catalyst — provides an alternate reaction pathway with lower activation energy, speeding up the reaction without being consumed itself.
• Surface area — powdered/finely divided solids react faster than large lumps, since more surface is exposed for collisions.
• Nature of reactants — ionic reactions (e.g., in solution) are typically much faster than reactions involving covalent bond-breaking.

🥤

Ever notice an antacid tablet dissolves faster when crushed than when swallowed whole? That's surface area at work — more exposed area means more collisions with stomach acid per second.

3. Rate Law, Order and Molecularity

Rate Law and Rate Constant

The rate law is an experimentally determined equation relating the rate of a reaction to the concentrations of reactants raised to specific powers.

FORMULA

Rate = k [A]^x [B]^y

k is the rate constant (specific reaction rate) — the rate when all reactant concentrations equal 1 mol/L. x and y are experimentally found exponents, NOT taken from the balanced equation.

COMMON MISTAKE Students often assume the rate law exponents equal the stoichiometric coefficients from the balanced equation. This is wrong in general — rate law exponents must be determined experimentally (usually via the method of initial rates). They only coincide with coefficients for elementary reactions.

Order of Reaction vs Molecularity

Order of Reaction	Molecularity of Reaction
Sum of powers of concentration terms in the experimentally determined rate law	Number of reacting species (atoms, ions, molecules) taking part in one elementary step
Can be zero, fractional, negative, or whole number	Always a positive whole number (1, 2, or rarely 3)
Applies to the overall reaction (may involve several steps)	Meaningful only for elementary (single-step) reactions
Experimentally determined	Theoretical concept from the proposed mechanism

MEMORY TRICK "Order is Observed (from experiment); Molecularity is from the Mechanism (theoretical)." Order can break the rules (zero, fraction); molecularity always plays by whole numbers.

Pseudo First-Order Reactions

Some reactions are technically higher order but behave like first order because one reactant is present in large excess (its concentration barely changes during the reaction). Classic NCERT examples:

• Hydrolysis of ethyl acetate: CH₃COOC₂H₅ + H₂O → CH₃COOH + C₂H₅OH (water is in vast excess as solvent)
• Inversion of cane sugar: C₁₂H₂₂O₁₁ + H₂O → C₆H₁₂O₆ + C₆H₁₂O₆ (glucose + fructose)

Zero Order

Rate independent of concentration; straight line, constant slope

First Order

Rate ∝ [A]; exponential decay curve

4. Integrated Rate Equations

NCERT restricts integrated rate laws to zero order and first order only (this is explicitly stated in the syllabus).

Zero Order Reaction

FORMULA

[R] = [R]₀ − kt   or   k = ([R]₀ − [R]) / t

Rate is independent of reactant concentration. Units of k: mol L⁻¹ s⁻¹. Example: decomposition of gaseous ammonia on a hot platinum surface (at high pressure); enzyme-catalysed reactions when the enzyme is saturated.

First Order Reaction

FORMULA

k = (2.303/t) log([R]₀/[R])

Units of k: s⁻¹ (time⁻¹ only — no concentration term). This is the single most frequently tested formula in this chapter across CBSE, NEET, and JEE.

Alternative exponential form:

FORMULA

[R] = [R]₀ e^(−kt)

EXAM TIP Check the units of k given in a question — they instantly tell you the order! mol L⁻¹ s⁻¹ → zero order; s⁻¹ (or min⁻¹) → first order; mol⁻¹ L s⁻¹ → second order. This is a favourite 1-mark MCQ trick in JEE Main.

Order	Differential Rate Law	Integrated Rate Law	Units of k	Half-Life (t½)
Zero	Rate = k	[R]₀ − [R] = kt	mol L⁻¹ s⁻¹	[R]₀ / 2k
First	Rate = k[R]	k = (2.303/t) log([R]₀/[R])	s⁻¹	0.693 / k

5. Half-Life of a Reaction (t½)

The half-life is the time required for the concentration of a reactant to fall to half its initial value.

FORMULA

Zero order: t½ = [R]₀ / 2k

Half-life depends on initial concentration — directly proportional to it.

FORMULA

First order: t½ = 0.693 / k

Half-life is INDEPENDENT of initial concentration — this is the defining fingerprint of a first-order reaction. Every radioactive decay (a first-order process) is tested using exactly this idea.

MEMORY TRICK "First order doesn't care where you started — it always takes the same time to halve." Use this line to instantly identify first-order behaviour in a word problem: if halving time stays constant regardless of starting amount, it's first order.

☢️

Radioactive decay (like Carbon-14 dating) is a textbook first-order process — this is exactly why t½ for carbon-14 is fixed at ~5730 years no matter how much carbon you start with.

6. Collision Theory of Reaction Rates

Collision theory explains reaction rate at the molecular level. NCERT specifies this only needs an elementary, non-mathematical treatment — no derivations expected, but the concepts are frequently asked as short-answer questions.

• Reactions occur when reactant molecules collide with each other.
• Not every collision leads to a reaction — only effective collisions do.
• An effective collision needs two conditions: (1) molecules must collide with energy equal to or greater than a minimum threshold energy, and (2) molecules must collide with proper orientation.

FORMULA

Rate = P × Z_AB × e^(−Ea/RT)

P = steric/probability factor (accounts for orientation); Z_AB = collision frequency; e^(−Ea/RT) = fraction of molecules with energy ≥ activation energy.

Activation Energy (Ea)

The minimum extra energy that reactant molecules must acquire (above their average energy) to reach the "activated complex" or transition state and successfully form products.

NCERT KEY POINT Threshold energy = Activation energy + average kinetic energy of reactants. A catalyst works by lowering the activation energy through an alternate pathway — it does NOT change the enthalpy (ΔH) of the reaction.

7. Arrhenius Equation and Temperature Dependence

Proposed by Svante Arrhenius, this equation quantifies how rate constant k varies with temperature.

FORMULA

k = A e^(−Ea/RT)

A = Arrhenius (frequency) factor; Ea = activation energy; R = universal gas constant (8.314 J K⁻¹ mol⁻¹); T = absolute temperature in Kelvin.

Taking natural log (useful for graphical/linear-plot questions):

FORMULA

ln k = ln A − Ea/RT   or   log k = log A − Ea/(2.303RT)

Plotting log k vs 1/T gives a straight line with slope = −Ea/2.303R and intercept = log A. NEET frequently tests this exact plot.

Two-temperature form (most useful for numericals — given k at two temperatures, find Ea, or vice versa):

FORMULA

log(k₂/k₁) = (Ea/2.303R) × [(T₂−T₁)/(T₁T₂)]

This is the JEE Main favourite — expect it whenever a question gives "rate constant at two different temperatures."

EXAM TIP Always convert temperature to Kelvin before substituting into any Arrhenius calculation. A huge chunk of numerical errors come from forgetting to add 273.

8. Solved Numericals

Numerical 1 — Average Rate

Q. For the reaction R → P, the concentration of R changes from 0.03 M to 0.02 M in 25 minutes. Calculate the average rate in both minutes and seconds.

Solution:

Average rate = −([R]₂ − [R]₁)/(t₂ − t₁) = −(0.02 − 0.03)/25 = 0.01/25 = 4 × 10⁻⁴ mol L⁻¹ min⁻¹

Converting to seconds: 4 × 10⁻⁴ / 60 = 6.66 × 10⁻⁶ mol L⁻¹ s⁻¹

Numerical 2 — First-Order Rate Constant

Q. A first-order reaction has a rate constant of 1.15 × 10⁻³ s⁻¹. How long will 5 g of this reactant take to reduce to 3 g?

Solution:

Using k = (2.303/t) log([R]₀/[R]):

t = (2.303/k) log([R]₀/[R]) = (2.303 / 1.15×10⁻³) × log(5/3)

t = 2002.6 × log(1.667) = 2002.6 × 0.2219

t ≈ 444.4 seconds

Numerical 3 — Half-Life (First Order)

Q. A first-order reaction takes 40 minutes for 30% decomposition. Calculate its t½.

Solution:

If [R]₀ = 100, then [R] = 70 after 30% decomposition.

k = (2.303/40) log(100/70) = (2.303/40) × 0.1549 = 0.00892 min⁻¹

t½ = 0.693/k = 0.693/0.00892 ≈ 77.7 minutes

Numerical 4 — Arrhenius Equation (Two Temperatures)

Q. The rate constant of a reaction is 2 × 10⁻² s⁻¹ at 300 K and 8 × 10⁻² s⁻¹ at 340 K. Calculate Ea. (R = 8.314 J K⁻¹mol⁻¹)

Solution:

log(k₂/k₁) = (Ea/2.303R) × (T₂−T₁)/(T₁T₂)

log(8×10⁻²/2×10⁻²) = log(4) = 0.602

0.602 = (Ea / (2.303 × 8.314)) × (40/(300×340))

0.602 = (Ea / 19.15) × 3.92×10⁻⁴

Ea = (0.602 × 19.15) / (3.92×10⁻⁴) ≈ 29,415 J/mol ≈ 29.4 kJ/mol

9. Practice MCQs (With Answers)

Q1. The unit of rate constant for a zero-order reaction is:

1. s⁻¹
2. mol L⁻¹ s⁻¹
3. mol⁻¹ L s⁻¹
4. mol⁻² L² s⁻¹

Show Answer

Answer: B — mol L⁻¹ s⁻¹. Zero-order rate constant carries the same units as rate itself, since rate = k (concentration-independent).

Q2. For a first-order reaction, the half-life period is:

1. Proportional to initial concentration
2. Independent of initial concentration
3. Inversely proportional to initial concentration
4. Proportional to the square of initial concentration

Show Answer

Answer: B. t½ = 0.693/k for first order — no concentration term appears, so it's constant regardless of starting amount.

Q3. Which of the following best describes molecularity?

1. Can be zero or fractional
2. Determined experimentally for the overall reaction
3. Always a whole number, meaningful only for elementary steps
4. Same as order for all reactions

Show Answer

Answer: C. Molecularity is a theoretical, whole-number concept applying only to single elementary steps — unlike order, which is experimental and can be fractional or zero.

Q4. The inversion of cane sugar is an example of a:

1. Zero-order reaction
2. Second-order reaction
3. Pseudo first-order reaction
4. Third-order reaction

Show Answer

Answer: C. Although two reactants are involved, water is present in vast excess, so the reaction behaves kinetically as first order.

Q5. A plot of log k versus 1/T gives a straight line whose slope equals:

1. −Ea/R
2. Ea/2.303R
3. −Ea/2.303R
4. 2.303R/Ea

Show Answer

Answer: C. From log k = log A − Ea/2.303RT, the slope of log k vs 1/T is −Ea/2.303R.

Q6. Which factor does a catalyst primarily alter in a reaction?

1. Enthalpy of reaction (ΔH)
2. Activation energy
3. Equilibrium constant
4. Order of reaction

Show Answer

Answer: B. A catalyst lowers activation energy by providing an alternate pathway; it does not change ΔH or the equilibrium constant.

10. Previous Year Questions (CBSE / NEET / JEE Pattern)

CBSE BOARD — 3 MARKS (RECURRING PATTERN) Define rate of reaction. Write the rate expression for the reaction 2A + B → 3C in terms of each reactant and product.

CBSE BOARD — NUMERICAL (RECURRING PATTERN) A first-order reaction has a rate constant 1 × 10⁻³ s⁻¹. Calculate the time required to reduce the initial concentration of the reactant to its 1/4th value.

(Hint: use [R]₀/[R] = 4, then apply k = (2.303/t) log([R]₀/[R]).)

NEET UG PATTERN For a chemical reaction, the rate is doubled when temperature is increased from 300 K to 310 K. Calculate the activation energy using the Arrhenius relation.

JEE MAIN PATTERN The rate constants of a reaction at two temperatures 293 K and 313 K are given. Compute the activation energy Ea using the two-temperature Arrhenius equation.

CBSE BOARD — CONCEPTUAL (SHORT ANSWER) Distinguish between order and molecularity of a reaction with a suitable example of each.

11. Frequently Asked Questions

Is Chemical Kinetics Chapter 3 or Chapter 4 in Class 12 Chemistry?

In the current rationalised NCERT syllabus (2026–27), Chemical Kinetics is Chapter 3. Some older editions and a few state boards number it as Chapter 4 — the content covered is the same either way.

Can order of a reaction be negative or fractional?

Yes. Unlike molecularity, order is purely experimental and can be zero, a fraction (e.g., 1.5), or even negative if increasing a species' concentration decreases the rate.

Why is molecularity meaningless for complex reactions?

Because molecularity applies only to a single elementary step. A complex, multi-step reaction has different molecularity for each step, so quoting one overall "molecularity" for the whole reaction has no physical meaning — only the overall order is meaningful there.

What is the difference between rate constant and rate of reaction?

Rate of reaction changes continuously as concentrations change during the reaction. Rate constant (k) is a fixed value at a given temperature — it doesn't depend on concentration, only on temperature (via the Arrhenius equation).

Is this chapter important for NEET and JEE Main?

Yes — Chemical Kinetics regularly contributes numerical questions in both exams, especially on integrated rate laws, half-life, and the Arrhenius equation. It's considered a high-yield, formula-driven chapter once practiced.

Chapter Summary — Quick Recall

• Rate of reaction = change in concentration ÷ time; can be average or instantaneous.
• Rate law is experimental: Rate = k[A]^x[B]^y, where x, y need not equal stoichiometric coefficients.
• Order (experimental, can be zero/fractional) ≠ Molecularity (theoretical, always a whole number, elementary steps only).
• Zero order: [R]₀ − [R] = kt, units of k = mol L⁻¹ s⁻¹, t½ = [R]₀/2k.
• First order: k = (2.303/t) log([R]₀/[R]), units of k = s⁻¹, t½ = 0.693/k (independent of initial concentration).
• Pseudo first-order: higher-order reaction that behaves as first order because one reactant is in large excess.
• Collision theory: only collisions with sufficient energy (≥ activation energy) and correct orientation are effective.
• Arrhenius equation: k = Ae^(−Ea/RT); log k vs 1/T plot gives slope −Ea/2.303R.
• Catalysts lower activation energy via an alternate pathway; they don't alter ΔH or equilibrium constant.

Jnaanangkur — The Learning Hub · Class 12 Chemistry Series

Content verified against the NCERT 2026–27 rationalised syllabus. For CBSE, AHSEC, NEET, JEE Main, and CUET preparation.