Schooling · Class 12 · Chemistry · Chemical Kinetics · Chapter on a page
CBSE Class 12 · Chemistry · NCERT Chemistry-I
Chemical Kinetics: chapter on a page
A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).
About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.
This note's current text has no applied fact-check on record.
Mind map
Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.
The mind map as a list
- Key ideas
- Rate of reaction
- Rate law
- Rate constant (k)
- Order of reaction
- Integrated rate equations
- Half-life (t₁/₂)
- Arrhenius equation
- Formulas and facts
- Average rate
- Instantaneous rate
- First-order integrated rate law
- Half-life for first order
- Zero-order integrated rate law
- Arrhenius equation (logarithmic form)
- Two-temperature form
- Worked examples
- Calculating average rate
- First-order rate constant and…
- Applying the Arrhenius equation
- Watch out for
- Confusing order with stoichiometric…
- Using natural log (ln) formula but…
- Forgetting to include the negative…
- Assuming half-life is constant for…
- Mixing up activation energy (Ea)…
What this chapter is about
Chemical kinetics is the branch of chemistry that studies the rates of chemical reactions and the factors that influence them. While thermodynamics tells us whether a reaction is feasible and how much product can form, kinetics tells us how fast the reaction proceeds. Many thermodynamically favourable reactions are practically useless because they occur too slowly, and understanding kinetics helps us control reaction speeds in industrial processes, biological systems and everyday life.
Key ideas
- Rate of reaction is defined as the change in concentration of a reactant or product per unit time. For a reaction A → B, rate = –d[A]/dt = +d[B]/dt, where the negative sign accounts for decreasing reactant concentration.
- Rate law is an experimentally determined expression that relates the rate to the concentrations of reactants raised to certain powers: rate = k[A]^m[B]^n. The exponents m and n are the orders with respect to A and B, and their sum is the overall order.
- Rate constant (k) is a proportionality constant specific to a reaction at a given temperature; its units depend on the overall order of the reaction.
- Order of reaction cannot be predicted from the stoichiometric equation; it must be found by experiment. Reactions can be zero order, first order, second order, or even fractional order.
- Integrated rate equations connect concentration with time. For first-order reactions: ln[A]₀/[A] = kt, or [A] = [A]₀ × e^(–kt).
- Half-life (t₁/₂) is the time for half the reactant to be consumed. For first-order reactions, t₁/₂ = 0.693/k, which is independent of initial concentration.
- Arrhenius equation describes how the rate constant changes with temperature: k = A × e^(–Ea/RT), where Ea is the activation energy and A is the pre-exponential factor.
- Catalysts increase the reaction rate by providing an alternative pathway with lower activation energy; they are regenerated and do not appear in the overall stoichiometry.
Formulas and facts to remember
- Average rate = Δ[product]/Δt or –Δ[reactant]/Δt (units: mol L⁻¹ s⁻¹).
- Instantaneous rate = d[product]/dt, found as the slope of the concentration–time curve at a point.
- First-order integrated rate law: ln[A] = ln[A]₀ – kt, or equivalently k = (2.303/t) × log([A]₀/[A]).
- Half-life for first order: t₁/₂ = 0.693/k (constant, independent of [A]₀).
- Zero-order integrated rate law: [A] = [A]₀ – kt; half-life t₁/₂ = [A]₀/(2k).
- Arrhenius equation (logarithmic form): log k = log A – Ea/(2.303 RT).
- Two-temperature form: log(k₂/k₁) = (Ea/2.303 R) × (1/T₁ – 1/T₂).
- Units of k: For zero order, mol L⁻¹ s⁻¹; for first order, s⁻¹; for second order, L mol⁻¹ s⁻¹.
Common mistakes
- Confusing order with stoichiometric coefficient → Order is found only from experiment; never assume it equals the coefficient in the balanced equation.
- Using natural log (ln) formula but substituting log₁₀ values → Remember ln x = 2.303 × log₁₀ x; keep logarithm bases consistent.
- Forgetting to include the negative sign when writing rate in terms of a reactant → Rate is always positive; the negative sign compensates for decreasing reactant concentration.
- Assuming half-life is constant for all reaction orders → Half-life is concentration-independent only for first-order reactions; for zero and second order, it depends on initial concentration.
- Mixing up activation energy (Ea) with enthalpy change (ΔH) → Ea is the energy barrier to reach the transition state; ΔH is the overall energy difference between reactants and products.
Quick revision
- Rate = change in concentration per unit time; units are mol L⁻¹ s⁻¹.
- For first-order reactions, plot ln[A] vs t gives a straight line with slope = –k, and t₁/₂ = 0.693/k.
- Order must be determined experimentally; it can be zero, fractional or integer.
- Arrhenius equation: k = A × e^(–Ea/RT); higher temperature or lower Ea means faster reaction.
- Catalysts lower Ea without changing the equilibrium position or being consumed.