NEET Chemistry · NCERT Class 12

Chemical Kinetics notes

Rate of Reaction, Factors Affecting Rate, Order & Molecularity, Integrated Rate Equations and Arrhenius Equation & Collision Theory — explained topic by topic with key points, formulas, memory tricks and common mistakes.

  • 6 topics
  • 50 min read
  • 27 formulas
  • 46 PYQs
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1

Chapter Overview

7 min

Chemical Kinetics studies how fast chemical reactions occur and why their speeds change under different conditions. In NCERT and NEET, the chapter revolves around rate of reaction, rate law, order, molecularity, integrated rate equations, half-life, Arrhenius equation, activation energy and collision theory. Thermodynamics tells whether a reaction is possible, but kinetics tells how quickly it reaches completion. A reaction may be thermodynamically feasible but kinetically slow, like conversion of diamond to graphite. This chapter is highly numerical and graph-based. NEET commonly asks units of rate constant, graphical identification of order, half-life dependence, pseudo first order reactions and activation energy from Arrhenius plots.

Key points6

  • The chapter has two major numerical areas: integrated rate laws and Arrhenius equation.
  • Graph slopes are very important: zero order [R] vs t has slope -k, first order log[R] vs t has slope -k/2.303, and Arrhenius plot log k vs 1/T has slope -Ea/2.303R.
  • Order may be zero, fractional, integral or negative; molecularity is always a positive integer.
  • Rate constant k is independent of concentration but strongly depends on temperature and catalyst.
  • Half-life trends are powerful shortcuts for identifying order.
  • Collision theory explains that only properly oriented collisions with energy greater than threshold energy are effective.

Memory tricks2

Kinetics Keyword: Remember: Kinetics = K for 'Kaise fast?' It answers how fast a reaction occurs, not whether it is possible.

Graph Shortcut: Straight line graph reveals order: [R] vs t means zero order, log[R] vs t means first order, 1/[R] vs t means second order.

Examples2

Rusting of Iron

Rusting is thermodynamically possible but slow. Moisture, oxygen and surface conditions affect its rate.

Food Spoilage

Food spoils slower in a refrigerator because lower temperature decreases rate constants of biochemical reactions.

Common mistakes2

Confusing Thermodynamics and Kinetics

A negative ΔG means a reaction is feasible, but it may still be extremely slow without a suitable pathway.

Assuming Balanced Equation Gives Rate Law

For complex reactions, rate law must be determined experimentally; stoichiometric coefficients are not automatically orders.

Reference tables2

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Formulas4

General Rate of Reaction
Rate =−Δ[Reactant]/Δt=+Δ[Product]/Δt\text{Rate }= -Δ[\text{Reactant}]/Δt = +Δ[\text{Product}]/Δt

Measures concentration change per unit time. Negative sign is used for reactant disappearance, positive sign for product formation.

Δ[Reactant]
Change in concentration of reactant
Δ[Product]
Change in concentration of product
Δt
Time interval
Rate Law
Rate =k[A]x[B]y\text{Rate }= k[A]^x[B]^y

Experimental relationship between reaction rate and molar concentrations of reactants.

k
Rate constant
x, y
Orders with respect to A and B
x + y
Overall order of reaction
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2

Rate of Reaction

8 min

Rate of reaction tells how quickly reactants are consumed or products are formed. For a reaction, it is expressed as change in concentration per unit time and usually has unit mol L⁻¹ s⁻¹. Average rate is calculated over a finite time interval, while instantaneous rate is the slope of tangent to concentration-time curve at a particular moment. Rate expression must consider stoichiometric coefficients so that the same reaction rate is obtained from any reactant or product. Rate law expresses rate in terms of molar concentrations raised to experimentally determined powers. NEET often tests signs, units, stoichiometric division, rate law interpretation and numerical substitution.

Key points6

  • 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.
  • Instantaneous rate is more accurate for changing rates and is obtained graphically from tangent slope.
  • Rate law is not the same as balanced chemical equation for complex reactions.
  • The rate constant k equals rate when all concentration terms are unity.
  • The units of k change with overall order, but unit of rate remains concentration per time.
  • Reaction rate generally decreases with time as reactant concentration decreases.

Memory tricks2

Reactant Negative, Product Positive: Reactants go down, so put minus sign; products go up, so put plus sign.

Rate Law is Experimental: Remember 'LAW is from LAB': rate law comes from experiment, not from just looking at the balanced equation.

Examples2

Solved Numerical: Average Rate

If [A] decreases from 0.50 M to 0.30 M in 20 s, average rate of disappearance of A = -(0.30 - 0.50)/20 = 0.010 mol L⁻¹ s⁻¹.

PYQ Concept: Rate Expression

For 2A + B → 3C, if d[C]/dt = 0.06 M s⁻¹, reaction rate = 1/3 × 0.06 = 0.02 M s⁻¹ and rate of disappearance of A = 2 × 0.02 = 0.04 M s⁻¹.

Common mistakes3

Forgetting Stoichiometric Division

For 2N2O5 → 4NO2 + O2, the reaction rate is -1/2 d[N2O5]/dt, not simply -d[N2O5]/dt.

Wrong Unit of Rate

Rate unit remains mol L⁻¹ s⁻¹ even if the reaction order changes; only the unit of k changes.

Using Product Coefficient as Order

Order is not obtained from product coefficients. It is obtained from the rate law.

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Formulas4

Average Rate
Average rate =−Δ[R]/Δt=+Δ[P]/Δt\text{Average rate }= -Δ[R]/Δt = +Δ[P]/Δt

Used when concentration change is measured over a finite time interval.

Δ[R]
Change in reactant concentration
Δ[P]
Change in product concentration
Δt
Time interval
Instantaneous Rate
Instantaneous rate =−d[R]/dt=+d[P]/dt\text{Instantaneous rate }= -d[R]/dt = +d[P]/dt

Rate at a particular instant; equal to slope of tangent on concentration-time graph.

d[R]
Infinitesimal change in reactant concentration
dt
Infinitesimal time interval
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3

Factors Affecting Rate

7 min

Reaction rate changes because the number and effectiveness of molecular collisions change. Ionic reactions in aqueous solution are often very fast, while covalent bond-breaking reactions may be slower because old bonds must break and new bonds must form. Increasing concentration or pressure usually increases collision frequency. Raising temperature increases kinetic energy and greatly increases the fraction of molecules crossing activation energy. A catalyst provides an alternate pathway with lower activation energy and speeds both forward and reverse reactions. Greater surface area allows more particles to be exposed for reaction. NEET frequently asks qualitative effects, catalyst misconceptions, temperature coefficient and graphical observations.

Key points6

  • Rate increases when effective collision frequency increases.
  • Temperature affects both kinetic energy and Arrhenius factor e^(-Ea/RT).
  • A catalyst does not change ΔG, ΔH or equilibrium constant; it only helps reach equilibrium faster.
  • Surface area matters mainly for heterogeneous reactions involving solids.
  • Pressure affects rate mainly for gaseous reactions.
  • The same factor may affect different reactions differently depending on rate law.

Memory tricks2

FACTS Control Rate: F-A-C-T-S: Frequency of collisions, Activation energy, Concentration, Temperature, Surface area.

Catalyst Shortcut: Catalyst = 'Cut Ea'. It cuts activation energy by providing a new pathway.

Examples3

Experimental Observation: Marble and HCl

Powdered CaCO3 reacts faster with dilute HCl than marble chips because more CaCO3 surface is exposed to acid.

Real Life: Refrigerator

Low temperature slows bacterial and enzymatic reactions, so food lasts longer in a refrigerator.

PYQ Concept: Pressure

Increasing pressure increases rate significantly for gaseous reactants but not directly for reactions in dilute liquid solutions.

Common mistakes3

Thinking Catalyst Changes Equilibrium

Catalyst speeds up attainment of equilibrium but does not change equilibrium constant or final equilibrium composition.

Saying Temperature Lowers Ea

Temperature increases molecular energy; it does not lower activation energy. Catalyst lowers activation energy.

Ignoring Surface Area

A solid lump and same mass of powder can react at very different rates due to exposed area.

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Formulas3

Concentration Dependence
r=k[A]x[B]yr = k[A]^x[B]^y

Changing concentration changes rate according to reaction orders.

x
Order with respect to A
y
Order with respect to B
k
Rate constant at fixed temperature
Temperature Dependence
k=Ae(−Ea/RT)k = A e^(-Ea/RT)

Rate constant increases when temperature increases because the exponential term increases.

Ea
Activation energy
T
Absolute temperature
A
Frequency factor
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4

Order & Molecularity

9 min

Order and molecularity describe reaction dependence and mechanism, but they are not the same. Order is the sum of powers of concentration terms in the experimentally determined rate law. It may be zero, fractional, integral or even negative. Molecularity is the number of reacting species participating in one elementary step and is always a positive integer. Zero order reactions have rate independent of reactant concentration, first order reactions have rate proportional to one concentration term, and second order reactions have total concentration power two. Pseudo first order reactions are actually higher order but behave as first order because one reactant is in large excess, such as hydrolysis of ester in water.

Key points6

  • Overall order = sum of powers in rate law.
  • Molecularity indicates how many species collide in an elementary event.
  • Complex reactions occur through multiple elementary steps, and the slowest step often controls rate.
  • A first order reaction has k unit s⁻¹.
  • For zero order reactions, rate remains constant until reactant is nearly exhausted.
  • Pseudo first order reactions simplify difficult kinetics and are common in solution chemistry.

Memory tricks3

Order is from Rate Law: Order = 'Power in rate law'. Look at exponents, add them, and ignore product terms.

Molecularity Must be Countable: Molecularity counts molecules colliding, so it cannot be zero or fractional.

Pseudo Means Fake Appearance: Pseudo first order is not truly first order; it only appears first order because one reactant is in excess.

Examples3

Numerical Example: Finding Order

If rate law is r = k[A]^1[B]^2, order with respect to A is 1, with respect to B is 2, and overall order is 3.

Pseudo First Order Example

Acid hydrolysis of ethyl acetate: CH3COOC2H5 + H2O → CH3COOH + C2H5OH. Water is solvent and in large excess, so reaction behaves as first order with respect to ester.

PYQ Concept: Doubling Concentration

For r = k[A]^2, if [A] is doubled, new rate = k(2[A])^2 = 4k[A]^2, so rate becomes four times.

Common mistakes3

Taking Stoichiometric Coefficient as Order

For complex reactions, coefficients are not equal to order unless experimentally proven.

Using Molecularity for Overall Complex Reaction

Molecularity is meaningful for elementary steps, not for an overall complex reaction.

Calling Molecularity Zero

Zero molecularity is impossible because no reaction can occur without reacting species.

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Formulas5

Overall Order
Order =x+yforr=k[A]x[B]y\text{Order }= x + y for r = k[A]^x[B]^y

The total order is the sum of powers of concentration terms in the rate law.

x
Order with respect to reactant A
y
Order with respect to reactant B
Zero Order Rate Law
r=k[A]0=kr = k[A]^0 = k

Rate is independent of reactant concentration.

r
Rate of reaction
k
Zero order rate constant
First Order Rate Law
r=k[A]r = k[A]

Rate is directly proportional to concentration of one reactant.

[A]
Molar concentration of reactant A
k
First order rate constant
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5

Integrated Rate Equations

10 min

Integrated rate equations connect reactant concentration with time and allow calculation of rate constant, remaining concentration and half-life. They are essential because experimental concentration-time data is easier to measure than instantaneous rate. For zero order reactions, concentration falls linearly with time and half-life is directly proportional to initial concentration. For first order reactions, logarithmic concentration changes linearly with time and half-life is constant. For second order reactions, reciprocal concentration varies linearly with time and half-life is inversely proportional to initial concentration. NEET repeatedly asks identification of order from straight-line graph, slope formulas, half-life patterns and first order numerical calculations using log values.

Key points6

  • Integrated rate laws are used when concentration changes continuously with time.
  • For first order reactions, equal time intervals reduce concentration by the same fraction.
  • In first order decay, after n half-lives, remaining amount = initial amount/2^n.
  • Graphical representation is a quick way to identify reaction order.
  • Units of k confirm order: zero order M s⁻¹, first order s⁻¹, second order M⁻¹ s⁻¹.
  • Most radioactive decay and many decomposition reactions follow first order kinetics.

Memory tricks3

Graph Order Trick: Zero: [R] straight. First: log[R] straight. Second: 1/[R] straight.

First Order Half-Life: First order half-life is 'free from initial concentration': t1/2 = 0.693/k.

Slope Signs: Reactant concentration graphs go down for zero and first order, so slopes are negative. Reciprocal concentration graph goes up for second order, so slope is positive.

Examples3

Solved Numerical: First Order k

A first order reaction is 75% complete in 20 min. Remaining = 25% = 1/4. k = 2.303/20 log(100/25) = 2.303/20 log4 = 2.303/20 × 0.602 = 0.0693 min⁻¹.

Solved Numerical: Half-Life

If k = 0.0231 s⁻¹ for a first order reaction, t1/2 = 0.693/0.0231 = 30 s.

PYQ Concept: Graph Identification

If a plot of 1/[A] vs time is straight line, the reaction is second order with respect to A.

Common mistakes3

Using Natural Log and Common Log Wrongly

ln([R]0/[R]) = kt, but log([R]0/[R]) needs factor 2.303: k = 2.303/t log([R]0/[R]).

Assuming All Half-Lives are Constant

Only first order half-life is independent of initial concentration. Zero and second order half-lives depend on [R]0.

Wrong Slope for First Order Graph

For log[R] vs t, slope = -k/2.303, so k = -2.303 × slope.

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Formulas6

Zero Order Integrated Rate Law
[R]=[R]0−kt[R] = [R]0 - kt

Reactant concentration decreases linearly with time.

[R]
Concentration at time t
[R]0
Initial concentration
k
Zero order rate constant
t
Time
Zero Order Half-Life
t1/2=[R]0/(2k)t1/2 = [R]0/(2k)

Half-life is directly proportional to initial concentration.

t1/2
Half-life period
[R]0
Initial reactant concentration
k
Zero order rate constant
First Order Integrated Rate Law
k=2.303/tlog([R]0/[R])k = 2.303/t log([R]0/[R])

Most important NEET formula for first order reactions.

k
First order rate constant
[R]0
Initial concentration
[R]
Remaining concentration at time t
t
Time
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6

Arrhenius Equation & Collision Theory

9 min

Arrhenius equation explains the strong dependence of rate constant on temperature and activation energy. Molecules must acquire minimum energy to form an activated complex or transition state before converting into products. This minimum extra energy is activation energy. A small rise in temperature can cause a large increase in rate because many more molecules cross the activation energy barrier. Collision theory adds that not every collision gives product; collisions must have sufficient energy and proper orientation. The Arrhenius plot of log k versus 1/T is a straight line whose slope gives activation energy. NEET commonly asks activation energy calculation, catalyst effect and effective collision concepts.

Key points6

  • Threshold energy is the minimum total energy needed for reaction.
  • Activation energy = threshold energy - average energy of reactants.
  • Only a fraction of collisions are effective.
  • Frequency factor A represents collision frequency and orientation-related factors.
  • Catalyst changes the reaction pathway, not the initial or final energy of reactants and products.
  • For exothermic reactions, Ea for forward and backward reactions differ by enthalpy change.

Memory tricks3

Arrhenius Plot Slope: log k vs 1/T always slopes down because increasing 1/T means decreasing temperature, so k decreases.

Effective Collision: Collision must have 'E plus O': Enough energy plus proper Orientation.

Catalyst Rule: Catalyst changes the road, not the start or finish. It lowers the hill but reactant and product energies remain same.

Examples4

Solved Numerical: Activation Energy from Slope

If slope of log k vs 1/T plot is -5000 K, then slope = -Ea/2.303R. Ea = 5000 × 2.303 × 8.314 = 95720 J mol⁻¹ ≈ 95.7 kJ mol⁻¹.

Solved Numerical: Temperature Effect

If k2/k1 = 2 for a 10 K rise, the reaction rate approximately doubles. This is a common temperature coefficient idea.

Real Life: Enzymes

Enzymes are biological catalysts. They increase reaction rate by lowering activation energy through a suitable active site orientation.

PYQ Concept: Catalyst and Equilibrium

A catalyst increases both forward and reverse rates, so equilibrium is reached faster but equilibrium constant remains unchanged.

Common mistakes4

Using Celsius in Arrhenius Equation

Always use kelvin for T in Arrhenius calculations.

Wrong Sign in Two-Temperature Formula

If T2 > T1, k2 should usually be greater than k1. Check that log(k2/k1) is positive.

Saying Every Collision is Effective

Only collisions with energy greater than or equal to Ea and correct orientation lead to products.

Thinking Catalyst is Used Up

A catalyst participates in the mechanism but is regenerated at the end.

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Formulas5

Arrhenius Equation
k=Ae(−Ea/RT)k = A e^(-Ea/RT)

Relates rate constant to activation energy and temperature.

k
Rate constant
A
Frequency factor
Ea
Activation energy
R
Gas constant, 8.314 J mol⁻¹ K⁻¹
T
Absolute temperature
Logarithmic Arrhenius Equation
logk=logA−Ea/(2.303RT)log k = log A - Ea/(2.303RT)

Used for Arrhenius plot and activation energy from slope.

log k
Common logarithm of rate constant
log A
Intercept of Arrhenius plot
Ea/(2.303R)
Magnitude related to slope
Two-Temperature Arrhenius Equation
log(k2/k1)=Ea/(2.303R)×(T2−T1)/(T1T2)log(k2/k1) = Ea/(2.303R) × (T2 - T1)/(T1T2)

Used to calculate activation energy or rate constant at another temperature.

k1, k2
Rate constants at temperatures T1 and T2
T1, T2
Absolute temperatures in kelvin
Ea
Activation energy
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Videos for Chemical Kinetics

  • 🔥 Chemical Kinetics | 3D Animated | NCERT Class 12 Chemistry | Best for NEET 💯Best for NEET · English
  • CHEMICAL KINETICS | Class 12th ONE SHOTPW NEET English · English
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NEET PYQs from Chemical Kinetics

The latest 3 of 46 previous-year questions from this chapter.

  1. NEET 2026Set 50MediumQ1

    For an elementary chemical reaction, the Arrhenius plot is given below. If the energy of activation is $6.64\ \mathrm{kJ\ mol^{-1}}$ and $R = 8.3\ \mathrm{J\ K^{-1}\ mol^{-1}}$, the temperature at which the rate constant becomes $e^2\ \mathrm{min^{-1}}$, is

    • A$125\ \mathrm{K}$
    • B$150\ \mathrm{K}$
    • C$200\ \mathrm{K}$
    • D$250\ \mathrm{K}$

  2. NEET 2026Set 50MediumQ2

    $2\mathrm{A} \xrightarrow{k} \mathrm{B}$ is a zero-order reaction, where $k = 1.0\ \mathrm{mol\ L^{-1}\ min^{-1}}$. If the initial concentration of A is 2 M, then the time taken to complete $75\%$ of the reaction will be

    • A$1.5\ \mathrm{min}$
    • B$0.75\ \mathrm{min}$
    • C$1.0\ \mathrm{min}$
    • D$2.0\ \mathrm{min}$

  3. NEET 2026Set 11MediumQ3

    Match List I with List II: Choose the correct answer from the options given below :

    • AA–IV, B–II, C–I, D–III
    • BA–IV, B–III, C–I, D–II
    • CA–IV, B–III, C–II, D–I
    • DA–I, B–II, C–III, D–IV
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