Chemical Kinetics Class 12 Chapter Notes

Key Terms in Chemical Kinetics

Rate of Reaction

The change in concentration of a reactant or product per unit time. Units: mol L⁻¹ s⁻¹.

Rate Law (or Rate Expression)

An expression that relates the rate of a reaction to the concentration of the reactants raised to some power.

Order of Reaction

The sum of powers to which the concentration terms are raised in the rate law expression. It is an experimentally determined quantity.

Molecularity of Reaction

The number of reacting species (atoms, ions or molecules) that must collide simultaneously in an elementary reaction to bring about a chemical change.

Rate Constant (k)

The proportionality constant in the rate law. It is equal to the rate of reaction when the concentration of each reactant is unity.

Half-Life (t₁/₂)

The time taken for the concentration of a reactant to be reduced to half of its initial concentration.

Activation Energy (Ea)

The minimum extra amount of energy that must be supplied to the reactants to enable them to cross the energy barrier and form products.

Collision Frequency (Z)

The number of collisions per second per unit volume of the reaction mixture.

Integrated Rate Equations & Key Formulas

Order of Reaction vs. Molecularity

Collision Theory of Reaction Rates

The Collision Theory provides a qualitative explanation for how chemical reactions occur and why rates differ. It is based on the kinetic theory of gases. For a reaction to happen, the reacting molecules must collide with each other. However, not all collisions lead to product formation. A collision is only effective if it satisfies two main conditions:

1. Energy Barrier (Activation Energy): The colliding molecules must possess a certain minimum amount of energy, called the threshold energy. The extra energy that reactants need to acquire to reach this threshold energy is known as the activation energy (Ea). Only molecules that collide with kinetic energy greater than or equal to the activation energy can form an activated complex and subsequently, products.

1. Orientation Barrier: The colliding molecules must also have a proper orientation relative to each other at the moment of collision. If the molecules are not oriented correctly, they will simply bounce off each other, even if they have sufficient energy. For example, in the formation of CO₂ from CO and O₂, the oxygen atom of O₂ must collide with the carbon atom of CO for the reaction to proceed.

Therefore, the rate of reaction depends on the number of effective collisions, which is only a small fraction of the total collisions. The Arrhenius equation mathematically incorporates these factors: k = A * e^(-Ea/RT). Here, the term e^(-Ea/RT) represents the fraction of molecules with energy greater than or equal to Ea, and the pre-exponential factor 'A' (Arrhenius factor) is related to the collision frequency and the orientation factor.

Must Remember for Chemical Kinetics

• The rate constant 'k' depends only on temperature and the presence of a catalyst; it is independent of reactant concentration.
• Units of rate constant for an nth order reaction are (mol L⁻¹)^(1-n) s⁻¹.
• The half-life of a first-order reaction is constant and independent of the initial concentration.
• For a zero-order reaction, the rate is independent of the concentration of reactants.
• Molecularity is defined only for elementary reactions, whereas order is defined for both elementary and complex reactions.
• A catalyst increases the rate of both forward and backward reactions to the same extent by providing a new path with lower activation energy. It does not affect the enthalpy (ΔH) or Gibbs Free Energy (ΔG) of the reaction.
• Pseudo-first-order reactions are reactions which are not truly first order but are made to behave like one under certain conditions (e.g., hydrolysis of ester with excess water).
• The rate of reaction generally doubles for every 10°C rise in temperature. This is known as the temperature coefficient.

Worked Example: First-Order Kinetics

• {"problem":"A first-order reaction has a rate constant of 1.15 × 10⁻³ s⁻¹. How long will it take for 5 g of this reactant to reduce to 3 g?","solution":"For a first-order reaction, t = (2.303/k) log([R]₀/[R]).\nGiven: k = 1.15 × 10⁻³ s⁻¹, Initial amount [R]₀ = 5 g, Final amount [R] = 3 g.\n(Note: For first-order, ratio of concentrations is same as ratio of mass/moles).\nt = (2.303 / 1.15 × 10⁻³) log(5/3)\nt = (2.0026 × 10³) (log 5 - log 3)\nt = (2.0026 × 10³) (0.6990 - 0.4771)\nt = (2.0026 × 10³) * 0.2219\nt = 444.3 s (approx.)"}

Board Exam Traps & Tips

1. Units are Key: Always write the units for the rate of reaction (mol L⁻¹ s⁻¹) and the rate constant (k). The unit of 'k' can help you identify the order of the reaction if it's not given. Forgetting units can cost you marks.
2. Graphs: Be prepared for graphical questions. For a zero-order reaction, a plot of [R] vs t is a straight line with slope -k. For a first-order reaction, a plot of ln[R] vs t is a straight line with slope -k. The slope of log k vs 1/T gives -Ea / 2.303R.
3. Logarithms: Many numericals involve log. Remember log(a/b) = log a - log b. Practice with common log values (log 2, 3, 5, 7, 10) to be quick in calculations.
4. Pseudo-First Order: Questions on hydrolysis of sucrose or esters are classic examples. Identify that one reactant (like water) is in such large excess that its concentration is considered constant, making the reaction behave as first-order.

Quick Revision Check

• What is the unit of the rate constant 'k' for a second-order reaction? L mol⁻¹ s⁻¹
• For a reaction, A + B → Product, the rate law is given by, r = k[A]¹/² [B]². What is the order of the reaction? The overall order of the reaction is the sum of the powers: 1/2 + 2 = 2.5
• How does a catalyst increase the rate of a reaction? A catalyst provides an alternative reaction pathway with a lower activation energy, increasing the number of effective collisions.
• What is the half-life of a first-order reaction if its rate constant is 2.0 x 10⁻³ s⁻¹? t₁/₂ = 0.693 / k = 0.693 / (2.0 x 10⁻³ s⁻¹) = 346.5 s

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