CLASS 12 CHEMISTRY
CHAPTER 3 — CHEMICAL KINETICSChemical KineticsChemical kinetics is the study of the speed (rate) at which chemical reactions occur and the factors that influence reaction rates. Thermodynamics tells IF a reaction can occur; kinetics tells HOW FAST.
Conceptual • NCERT-Focused • NEET-Oriented
What is this chapter really about?
Thermodynamics tells you IF a reaction can happen; kinetics tells you HOW FAST it actually happens. This chapter is built around one experimentally-determined equation — the rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. — and everything else (order, rate constantrate constantThe proportionality constant in a rate law for a reaction at a fixed temperature., integrated rate equations, half-lifehalf-lifeThe time required for the concentration of a reactant to fall to half its initial value., Arrhenius equationArrhenius equationThe equation relating the rate constant to activation energy and temperature., collision theorycollision theoryA model stating that reactions occur through collisions with sufficient energy and suitable orientation.) exists to help you find, use, or explain that rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations.. Never derive a rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. from a balanced equation by inspection: it must come from experimental data.
Reactants collide → Effective collisions (right energy + right orientation) → Products form → Rate measured experimentally
1. RATE OF A REACTION
Rate = change in concentration of reactant/product per unit time. For R → P:
Average rateAverage rateAverage rate is the change in concentration over a finite time interval. Instantaneous rate is the rate at a specific moment (limit as delta t -> 0). = −Δ[R]/Δt = +Δ[P]/Δt
Instantaneous rate is the average rateAverage rateAverage rate is the change in concentration over a finite time interval. Instantaneous rate is the rate at a specific moment (limit as delta t -> 0). as Δt → 0, i.e. the slope of the tangent on a concentration–time graph at a chosen instant.
| 📖 NCERT FOCUS For a reaction with unequal stoichiometric coefficients, e.g. 2HI(g) → H₂(g) + I₂(g), the rate must be divided by each species' coefficient to give one single, unambiguous rate: Rate = −(1/2)Δ[HI]/Δt = Δ[H₂]/Δt = Δ[I₂]/Δt. |
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| ⚠ NEET TRAP Units of rate are always concentration × time⁻¹ (e.g. mol L⁻¹ s⁻¹) — regardless of the reaction's order. Do not confuse rate units with rate CONSTANTrate constantThe proportionality constant in a rate law for a reaction at a fixed temperature. units, which change with order. |
2. RATE LAWrate lawAn experimentally determined equation relating reaction rate to reactant concentrations., ORDER & RATE CONSTANTrate constantThe proportionality constant in a rate law for a reaction at a fixed temperature.
For aA + bB → products, the experimentally-determined rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. is Rate = k[A]ˣ[B]ʸ, where x and y are NOT necessarily equal to a and b.
| 🔥 HIGH-YIELD Rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. can ONLY be found from experimental data (initial-rate method) — never by simply reading stoichiometric coefficients off the balanced equation. This is the single most repeated conceptual NEET line in this chapter. |
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Order of reactionorder of reactionThe sum of the powers of reactant concentrations in the experimentally determined rate law. = x + y (sum of powers in the rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations., found experimentally)
| Order | Rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. shape | Units of k |
|---|---|---|
| Zero order | Rate = k | mol L⁻¹ s⁻¹ |
| First order | Rate = k[A] | s⁻¹ |
| Second order | Rate = k[A]² or k[A][B] | mol⁻¹ L s⁻¹ |
| nth order | Rate = k[A]ⁿ | mol⁽¹⁻ⁿ⁾ Lⁿ⁻¹ s⁻¹ |
| 🧠 MEMORY HOOK To find units of k quickly: units of k = (conc.)^(1−order) × time⁻¹. | ||
| ⚠ NEET TRAP Order can be zero, a positive/negative integer, OR a fraction (e.g. 3/2) — but molecularitymolecularityThe number of reacting species colliding in a single elementary step. can NEVER be zero or a fraction. |
3. MOLECULARITYmolecularityThe number of reacting species colliding in a single elementary step. VS ORDER
| Order | MolecularitymolecularityThe number of reacting species colliding in a single elementary step. | |
|---|---|---|
| Definition | Sum of powers of concentration terms in the experimental rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. | Number of reacting species colliding simultaneously in an ELEMENTARY step |
| Determined by | Experiment | The mechanism (theoretical, from the balanced elementary equation) |
| Can be | 0, fraction, negative, or integer | Only 1, 2, or 3 (never 0 or fraction) |
| Applies to | Elementary AND complex (multi-step) reactions | Only elementary reactions/steps |
| 📖 NCERT FOCUS For a complex, multi-step reaction, the ORDER equals the molecularitymolecularityThe number of reacting species colliding in a single elementary step. of the SLOWEST step — the rate-determining step (RDS) — just like a relay race is only as fast as its slowest runner. | ||
| ⚠ NEET TRAP A reaction that 'looks' high order from its overall balanced equation (e.g. KClO₃ + 6FeSO₄ + 3H₂SO₄ → ...) can actually be experimentally second order — proof that it proceeds through several elementary steps, not in one shot. | ||
| 🔥 HIGH-YIELD Classic example: H₂O₂ + I⁻ → H₂O + IO⁻ (slow, rate-determining, bimolecular) followed by H₂O₂ + IO⁻ → H₂O + I⁻ + O₂ (fast). Both steps are bimolecular elementary reactions; overall order = 2 (first in H₂O₂, first in I⁻), matching the slow step exactly. |
4. INTEGRATED RATE EQUATIONS
Zero Order Reactions
Rate is independent of reactant concentration — common on saturated catalyst surfaces (e.g. Pt) or in some enzyme-catalysed reactions, since the surface is already 'full' and further concentration increase cannot speed things up.
[R] = −kt + [R]0 → k = ([R]0 − [R]) / t
Straight-line plot: [R] vs t → slope = −k, intercept = [R]0. Example: 2NH₃(g) → N₂(g) + 3H₂(g) on hot Pt surface at high pressure.
First Order Reactions
ln[R] = −kt + ln[R]0 → k = (2.303/t)·log([R]0/[R])
Straight-line plot: ln[R] vs t → slope = −k. Or log([R]0/[R]) vs t → slope = k/2.303.
| 🔥 HIGH-YIELD All radioactive decay follows first-order kinetics. Other classics: N₂O₅ and N₂O decomposition, hydrogenation of ethene. |
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First order gas-phase reactions using pressure
For A(g) → B(g) + C(g), using total pressure pt at time t and initial pressure pi:
k = (2.303/t)·log[pi / (2pi − pt)]
5. HALF-LIFEhalf-lifeThe time required for the concentration of a reactant to fall to half its initial value. (T½)
The time in which the reactant concentration falls to half its initial value.
| Order | Half-lifehalf-lifeThe time required for the concentration of a reactant to fall to half its initial value. formula | Depends on initial concentration? |
|---|---|---|
| Zero order | t½ = [R]0 / 2k | YES — directly proportional |
| First order | t½ = 0.693 / k | NO — constant, independent of [R]0 |
| 🧠 MEMORY HOOK First-order half-lifehalf-lifeThe time required for the concentration of a reactant to fall to half its initial value. is a FIXED number for a given reaction — this is exactly why radioactive decay always quotes a single half-lifehalf-lifeThe time required for the concentration of a reactant to fall to half its initial value. value regardless of how much sample you start with. | ||
| ⚠ NEET TRAP Time for 99.9% completion of a first-order reaction = 10 × t½. Time for 99% completion = ~6.64 × t½ ≈ 2 × (time for 90% completion) — a frequently tested numeric relationship. |
6. PSEUDO FIRST ORDER REACTIONS
A higher-order reaction that behaves experimentally as first order because one reactant is present in such large excess that its concentration barely changes during the reaction.
| Reaction | True order | Reason it behaves as 1st order |
|---|---|---|
| Hydrolysis of ethyl acetate: CH3COOC2H5 + H2O → CH3COOH + C2H5OH | 2 (1st in ester, 1st in water) | Water taken in large excess — its concentration ≈ constant |
| Inversion of cane sugar: C12H22O11 + H2O → glucose + fructose | 2 | Water in large excess (also acid-catalysed) |
7. TEMPERATURE DEPENDENCE — ARRHENIUS EQUATIONArrhenius equationThe equation relating the rate constant to activation energy and temperature.
Rule of thumb: rate constantrate constantThe proportionality constant in a rate law for a reaction at a fixed temperature. roughly DOUBLES for every 10°C rise in temperature.
k = A·e^(−Ea/RT)
A = Arrhenius (pre-exponential/frequency) factor, related to collision frequency; Ea = activation energyactivation energyThe minimum energy barrier reactant molecules must overcome to form products. (J/mol); e^(−Ea/RT) = fraction of molecules with energy ≥ Ea.
ln k = −Ea/RT + ln A → Plot: ln k (y-axis) vs 1/T (x-axis) → straight line, slope = −Ea/R, intercept = ln A
Two-temperature form
log(k2/k1) = (Ea/2.303R)·[(T2−T1)/(T1T2)]
| 📖 NCERT FOCUS Activation energyactivation energyThe minimum energy barrier reactant molecules must overcome to form products. Ea is the minimum extra energy that colliding reactant molecules need (over and above their average energy) to cross the energy barrier and form the unstable activated complex/transition state. |
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| 🔥 HIGH-YIELD Maxwell–Boltzmann distribution curve: raising temperature shifts the peak (most probable kinetic energy) to a higher value AND broadens the curve, increasing the fraction of molecules with energy ≥ Ea — this is WHY rate increases with temperature, in molecular terms. |
| ⚠ NEET TRAP A catalyst LOWERS activation energyactivation energyThe minimum energy barrier reactant molecules must overcome to form products. by providing an alternate reaction pathway — it does NOT change ΔG, ΔH, or the equilibrium constant K of the reaction. It speeds up the forward AND backward reactions equally, so equilibrium is reached faster but the equilibrium position itself is unchanged. |
8. COLLISION THEORYcollision theoryA model stating that reactions occur through collisions with sufficient energy and suitable orientation.
Builds on Arrhenius's ideas using the kinetic theory of gases: a reaction occurs only when reactant molecules collide.
Rate = ZAB·e^(−Ea/RT) (ZAB = collision frequency)
But not every collision leads to product — only EFFECTIVE collisions do, which need BOTH: (1) sufficient kinetic energy (≥ threshold energy) and (2) proper orientation of colliding molecules.
Rate = P·ZAB·e^(−Ea/RT) (P = steric/probability factor)
| 🧠 MEMORY HOOK Threshold energy = Activation energyactivation energyThe minimum energy barrier reactant molecules must overcome to form products. + energy already possessed by the reacting species. |
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| ⚠ NEET TRAP Collision theorycollision theoryA model stating that reactions occur through collisions with sufficient energy and suitable orientation. works well for simple/atomic species but shows significant deviation for complex molecules, precisely because it ignores structural/orientation requirements unless the steric factor P is included. |
MUST-KNOW DIAGRAMS
⭐ Concentration vs time — zero and first order
Structure → Label → Function: Zero order gives a straight declining line ([R] vs t); first order gives a curved exponential decay in [R] vs t but a straight declining line in ln[R] vs t.
⭐ Potential energy vs reaction coordinate
Structure → Label → Function: Reactants at one energy level, an energy hump (activated complex, at the top) representing Ea, then products at their own (possibly lower) energy level; ΔH = energy of products − energy of reactants. A catalyst redraws this hump lower without changing the start/end energy levels.
Understand the structure — Maxwell–Boltzmann distribution
A skewed curve of fraction of molecules vs kinetic energy; raising T shifts the peak right and flattens/broadens the curve, increasing the shaded area beyond Ea (i.e. more molecules qualify as 'energetic enough').
Understand the structure — Arrhenius plot
ln k (y-axis) vs 1/T (x-axis) is a straight line with a NEGATIVE slope of magnitude Ea/R and a y-intercept of ln A.
LAST-MINUTE CONFUSION BUSTERS
Rate lawrate lawAn experimentally determined equation relating reaction rate to reactant concentrations. is experimental — never assume it equals stoichiometric coefficients.
Order can be zero/fraction/negative; molecularitymolecularityThe number of reacting species colliding in a single elementary step. is always a positive integer (1, 2, or 3).
For multi-step reactions, order = molecularitymolecularityThe number of reacting species colliding in a single elementary step. of the rate-determining (slowest) step.
First-order half-lifehalf-lifeThe time required for the concentration of a reactant to fall to half its initial value. is constant; zero-order half-lifehalf-lifeThe time required for the concentration of a reactant to fall to half its initial value. depends on initial concentration.
A catalyst changes the RATE (via lower Ea) but never changes ΔG°, K, or the equilibrium position.
Pseudo first order = true higher-order reaction where one reactant is in large excess.
Rate constantrate constantThe proportionality constant in a rate law for a reaction at a fixed temperature. units change with order; rate units (mol L⁻¹ s⁻¹) never change.