Chemical Kinetics Notes for NEET and JEE Main 2027: Key Points, PDF

Chemical Kinetics is Unit 3 of NCERT Class 12 Chemistry Part I and Unit 8 of both the NEET UG and JEE Main chemistry syllabus. These notes cover reaction rate, rate law, order and molecularity, integrated rate equations for zero and first order reactions, half-life, pseudo first order reactions, the Arrhenius equation, catalysts and collision theory, with a free PDF.

Chemical Kinetics Notes for NEET and JEE Main

Section-wise key points from NCERT Unit 3, formula tables, high-yield facts and common mistakes for 2027 revision.

Chemical Kinetics at a Glance

The topic lists follow the NEET UG 2026 syllabus released by the National Medical Commission (NMC) on 22 December 2025 and the JEE Main 2026 syllabus released by NTA on 31 October 2025, which stay the reference until the 2027 syllabi are out. Full lists: NEET syllabus and JEE Main syllabus. Nuclear chemistry, which the older version of this page covered, is not in the chemistry syllabus of either exam, so it is no longer included.

Rate of a Chemical Reaction

The rate of a reaction is the change in concentration of a reactant or product per unit time; chemical kinetics studies these rates and the mechanisms behind them.

• Thermodynamics tells whether a reaction is feasible; kinetics tells how fast it goes. Diamond should turn into graphite, but the change is too slow to notice.
• Very fast: precipitation of AgCl on mixing AgNO3 and NaCl. Very slow: rusting of iron. Moderate: inversion of cane sugar, hydrolysis of starch.
• Average rate = −Δ[R]/Δt = +Δ[P]/Δt. The minus sign keeps the rate positive because [R] decreases.
• Instantaneous rate = −d[R]/dt = d[P]/dt, found from the slope of the tangent to the concentration-time curve.
• Units: mol L⁻¹ s⁻¹; for gases, with partial pressures, atm s⁻¹.
• Divide by stoichiometric coefficients so all species give the same rate. For 2HI → H2 + I2: rate = −½ d[HI]/dt = d[H2]/dt = d[I2]/dt.

Rate Law, Rate Constant and Order

The rate law expresses rate in terms of reactant concentrations raised to powers that must be found by experiment, not read from the balanced equation.

• For aA + bB → products: rate = k[A]ˣ[B]ʸ, where x and y may or may not equal a and b. k is the rate constant.
• Order with respect to A is x, with respect to B is y, and the overall order is x + y. Order can be 0, 1, 2, 3 or a fraction.
• 2NO + O2 → 2NO2: doubling [NO] makes the rate four times, doubling [O2] doubles it, so rate = k[NO]²[O2].
• CHCl3 + Cl2 → CCl4 + HCl has rate = k[CHCl3][Cl2]^½; for ethyl acetate hydrolysis rate = k[CH3COOC2H5][H2O]⁰. Here exponents differ from coefficients.

Order and Molecularity: Elementary and Complex Reactions

An elementary reaction happens in one step; a complex reaction is a sequence of elementary steps (its mechanism), and the slowest step controls the overall rate.

• Unimolecular: NH4NO2 → N2 + 2H2O. Bimolecular: 2HI → H2 + I2. Termolecular: 2NO + O2 → 2NO2. Molecularity three is rare because simultaneous collision of three species is unlikely.
• KClO3 + 6FeSO4 + 3H2SO4 → KCl + 3Fe2(SO4)3 + 3H2O looks tenth order but is second order, so it must occur in several steps.
• Iodide-catalysed decomposition of H2O2 in alkaline medium: rate = k[H2O2][I⁻]. Step 1 (slow, rate-determining): H2O2 + I⁻ → H2O + IO⁻. Step 2: H2O2 + IO⁻ → H2O + I⁻ + O2. IO⁻ is an intermediate.
• For a complex reaction, the order is given by the slowest step, and the molecularity of the slowest step equals the overall order.

Integrated Rate Equations: Zero and First Order

Integrating the rate law links concentrations measured at different times directly to the rate constant.

• Zero order on a metal surface: at high pressure the surface is saturated with gas, so the rate stops depending on concentration. Some enzyme-catalysed reactions are also zero order.
• First order gas-phase reaction A(g) → B(g) + C(g): k = (2.303/t) log[pi/(2pi − pt)], where pi is the initial pressure and pt the total pressure at time t.
• Between two times: k = [2.303/(t2 − t1)] log([R]1/[R]2).
• For a first order reaction, 99.9% completion takes 10 half-lives (t = 6.909/k).

Pseudo First Order Reactions

A higher order reaction behaves as first order when one reactant is in such large excess that its concentration hardly changes.

• Acid hydrolysis of ethyl acetate: 0.01 mol ester with 10 mol water leaves 9.99 mol water, so rate = k[CH3COOC2H5].
• Inversion of cane sugar: C12H22O11 + H2O → C6H12O6 (glucose) + C6H12O6 (fructose) in acid; rate = k[C12H22O11].

Effect of Temperature: Arrhenius Equation

For most reactions the rate constant nearly doubles for a 10° rise in temperature; the Arrhenius equation describes this exactly.

• k = A e^(−Ea/RT), where A is the Arrhenius (frequency or pre-exponential) factor and Ea the activation energy in J mol⁻¹. van’t Hoff first proposed it; Arrhenius gave its physical justification.
• ln k = −Ea/RT + ln A, so a plot of ln k vs 1/T is a straight line with slope −Ea/R and intercept ln A.
• Two temperatures: log(k2/k1) = [Ea/(2.303R)] × [(T2 − T1)/(T1T2)].
• Reactants pass through an unstable activated complex; the energy needed to form it is the activation energy.
• The factor e^(−Ea/RT) is the fraction of molecules with energy equal to or greater than Ea. Raising temperature shifts the Maxwell-Boltzmann curve maximum to higher energy and broadens it; the area under the curve stays constant.
• Half-life of N2O5 decomposition: 12 min at 50 °C, 5 h at 25 °C and 10 days at 0 °C.
• Higher temperature or lower activation energy gives a faster reaction.

Effect of a Catalyst

A catalyst increases the rate without itself undergoing any permanent chemical change, by giving an alternative pathway with lower activation energy.

• Example: MnO2 catalyses 2KClO3 → 2KCl + 3O2.
• A substance that reduces the rate is an inhibitor, not a catalyst.
• Intermediate complex theory: the catalyst forms temporary bonds with reactants, giving a short-lived intermediate that breaks down to products and the catalyst.
• A catalyst does not change ΔG or the equilibrium constant. It speeds up forward and backward reactions equally, so equilibrium is reached sooner.
• It catalyses spontaneous reactions only, and a small amount catalyses a large amount of reactants.

Collision Theory of Chemical Reactions

Collision theory (Max Trautz and William Lewis, 1916-18), based on the kinetic theory of gases, treats molecules as hard spheres that react when they collide effectively.

• Collision frequency (Z) is the number of collisions per second per unit volume of the reaction mixture.
• For A + B → products: rate = Z_AB e^(−Ea/RT); comparing with Arrhenius, A is related to collision frequency.
• Effective collisions need enough energy (threshold energy) and proper orientation. Threshold energy = activation energy + energy possessed by the reacting species.
• Steric (probability) factor P accounts for orientation: rate = P Z_AB e^(−Ea/RT).
• Drawback: molecules are treated as hard spheres and their structure is ignored, so predictions deviate for complex molecules.

Most-asked Facts for NEET and JEE Main: Chemical Kinetics

Revise these one-line facts from the NCERT unit just before the exam.

• Rate law must be found experimentally, not from the balanced equation.
• Order can be zero or fractional; molecularity is only 1, 2 or 3.
• Molecularity applies only to elementary reactions.
• Units of k: zero order mol L⁻¹ s⁻¹; first order s⁻¹; second order mol⁻¹ L s⁻¹.
• Zero order t½ = [R]0/2k; first order t½ = 0.693/k.
• First order: 99.9% completion = 10 half-lives.
• NH3 decomposition on hot Pt at high pressure is zero order.
• All radioactive decay follows first order kinetics.
• Ester hydrolysis and inversion of cane sugar are pseudo first order.
• k = A e^(−Ea/RT); slope of ln k vs 1/T = −Ea/R.
• Rate constant nearly doubles for a 10° rise in temperature.
• A catalyst lowers Ea and does not change ΔG or K.
• Collision theory: rate = P Z_AB e^(−Ea/RT).

Common Mistakes to Avoid

Watch for these mix-ups when you answer questions on this unit.

• Writing the rate law from stoichiometric coefficients. Exponents come from experiment.
• Giving molecularity to a complex reaction. It has meaning only for elementary steps.
• Forgetting to divide by stoichiometric coefficients when relating rates of different species.
• Saying a first order half-life depends on initial concentration. It does not; a zero order half-life does.
• Using log instead of ln without the 2.303 factor, or mixing units of time between k and t.
• Thinking a catalyst shifts the equilibrium. It only helps equilibrium be reached faster.
• Treating every collision as effective. Collisions need enough energy and proper orientation.
• Plotting [R] vs t for a first order reaction. The straight line is ln[R] vs t.
• Mixing up the slope of ln k vs 1/T (−Ea/R) with the slope of log k vs 1/T (−Ea/2.303R).

Chemical Kinetics Notes PDF and Related Pages

Download the AglaSem notes for Chemical Kinetics (key notes and important questions with answers) free: View PDF

• All chapters: NEET and JEE Main Chemistry notes
• Full topic list: NEET syllabus and JEE Main syllabus
• Practice: NEET question papers and JEE Main question papers
• Source text: NCERT Unit 3 (NCERT, Class 12 Chemistry Part I)

Chemistry Notes for NEET and JEE Main: All Chapters

• Some Basic Concepts of Chemistry
• Structure of Atom
• Chemical Kinetics
• Equilibrium
• Organic Chemistry: Some Basic Principles and Techniques
• Electrochemistry
• Amines
• Solutions
• Thermodynamics
• Coordination Compounds
• Haloalkanes and Haloarenes
• p-Block Elements
• Alcohols, Phenols and Ethers
• Hydrocarbons
• Aldehydes, Ketones and Carboxylic Acids

All chapters with key points and notes PDFs are on the Chemistry notes for NEET and JEE Main page.

Chemical Kinetics Notes for NEET and JEE Main FAQs