Class 12 Chemistry Notes Chapter 2 (Chemical Kinetics) – Lab Manual (English) Book
Detailed Notes with MCQs of Chapter 2: Chemical Kinetics. This is a crucial chapter, not just for your board exams but also for various government competitive exams where chemistry is a component. We'll cover the essential concepts systematically.
Chemical Kinetics: Detailed Notes
1. Introduction
- Definition: Chemical Kinetics is the branch of chemistry that deals with the study of the rates (or speeds) of chemical reactions, the factors affecting these rates, and the mechanism by which reactions occur.
- Importance: Helps optimize reaction conditions in industries, understand biological processes (enzyme kinetics), and predict reaction pathways.
2. Rate of a Chemical Reaction
- Definition: The change in concentration of any one of the reactants or products per unit time.
- Average Rate: The rate measured over a considerable interval of time.
- For a reaction: R → P
- Average Rate = - Δ[R] / Δt = + Δ[P] / Δt
- Where Δ[R] = Change in concentration of reactant R, Δ[P] = Change in concentration of product P, Δt = Time interval.
- Negative sign for reactants indicates decrease in concentration; Positive sign for products indicates increase in concentration.
- Instantaneous Rate: The rate of reaction at a particular instant of time. It is the limit of the average rate as the time interval approaches zero (Δt → 0).
- Instantaneous Rate = - d[R] / dt = + d[P] / dt
- Graphically, it is the slope of the tangent drawn to the concentration vs. time curve at that specific instant.
- Units of Rate: Concentration / Time (e.g., mol L⁻¹ s⁻¹, mol L⁻¹ min⁻¹, atm s⁻¹ for gaseous reactions).
- Stoichiometry and Rate: For a general reaction: 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
- The rate of disappearance/appearance of each species is divided by its stoichiometric coefficient to get a unique value for the reaction rate.
3. Factors Affecting Reaction Rate
- Concentration of Reactants: Generally, the rate increases with an increase in the concentration of reactants (more collisions). This is quantified by the Rate Law.
- Temperature: Reaction rates generally increase significantly with an increase in temperature (usually doubles for every 10°C rise). This is because the kinetic energy of molecules increases, leading to more frequent and more energetic collisions (explained by Arrhenius Equation).
- Nature of Reactants and Products: The physical state, chemical reactivity, bond strengths, etc., influence the rate. Reactions involving ionic species are usually very fast, while those involving breaking covalent bonds are slower.
- Presence of a Catalyst: A catalyst increases the rate of reaction without being consumed itself by providing an alternative reaction pathway with lower activation energy.
- Surface Area of Reactants: For reactions involving solids, the rate increases with an increase in surface area (e.g., powdered solid reacts faster than a lump). More surface area means more contact points for reaction.
- Presence of Light (Radiation): Some reactions, called photochemical reactions, are initiated or accelerated by light (e.g., H₂(g) + Cl₂(g) → 2HCl(g)).
4. Rate Law and Rate Constant
- Rate Law (or Rate Equation): An experimentally determined expression that relates the rate of a reaction to the molar concentrations of the reactants, each raised to some power.
- For a reaction: aA + bB → Products
- Rate = k [A]ˣ [B]ʸ
- Where:
- k = Rate Constant (or specific reaction rate)
- x = Order of reaction with respect to A
- y = Order of reaction with respect to B
- Important: x and y are determined experimentally and may or may not be equal to the stoichiometric coefficients (a and b).
- Order of Reaction: The sum of the powers to which the concentration terms are raised in the experimentally determined rate law.
- Overall Order = x + y
- Order can be zero, positive, negative, or fractional.
- Rate Constant (k):
- Defined as the rate of reaction when the concentration of each reactant is unity (1 mol L⁻¹).
- It is characteristic of a reaction at a given temperature. Its value changes with temperature but is independent of reactant concentrations.
- Units of k: Depend on the overall order (n) of the reaction.
- Units of k = (Concentration)¹⁻ⁿ (Time)⁻¹ = (mol L⁻¹)¹⁻ⁿ s⁻¹
- Zero order (n=0): mol L⁻¹ s⁻¹
- First order (n=1): s⁻¹
- Second order (n=2): L mol⁻¹ s⁻¹
5. Molecularity of a Reaction
- Definition: The number of reacting species (atoms, ions, or molecules) that must collide simultaneously in an elementary reaction (a reaction occurring in a single step) to bring about a chemical change.
- Characteristics:
- It is a theoretical concept, applicable only to elementary reactions.
- It must be a whole number (1, 2, or 3). Molecularity greater than 3 is rare due to the low probability of simultaneous collision of more than three molecules.
- It can be determined from the stoichiometry of an elementary reaction.
- Difference between Order and Molecularity:
Feature Order of Reaction Molecularity of Reaction Definition Sum of powers in the rate law No. of species colliding in elementary step Nature Experimental concept Theoretical concept Values Can be 0, fraction, integer Always a positive integer (1, 2, 3) Applicability Overall reaction & elementary steps Only elementary steps Determination From Rate Law (experiment) From mechanism (elementary step stoichiometry)
6. Integrated Rate Equations
- These relate concentration directly to time.
- Zero-Order Reaction: Rate is independent of reactant concentration. Rate = k[R]⁰ = k.
- Integrated Rate Law: k = ([R]₀ - [R]) / t or [R] = [R]₀ - kt
- [R]₀ = Initial concentration, [R] = Concentration at time t.
- Plot of [R] vs. t is a straight line with slope = -k and intercept = [R]₀.
- Half-life (t½): Time taken for the concentration to reduce to half its initial value.
- t½ = [R]₀ / 2k (Half-life is directly proportional to initial concentration).
- First-Order Reaction: Rate is directly proportional to the first power of reactant concentration. Rate = k[R]¹.
- Integrated Rate Law: k = (2.303 / t) log ([R]₀ / [R]) or ln([R]₀ / [R]) = kt or [R] = [R]₀ e⁻ᵏᵗ
- Plot of log[R] vs. t is a straight line with slope = -k / 2.303 and intercept = log[R]₀.
- Plot of ln[R] vs. t is a straight line with slope = -k and intercept = ln[R]₀.
- Half-life (t½): t½ = 0.693 / k (Half-life is independent of initial concentration).
- Examples: Radioactive decay, decomposition of N₂O₅, decomposition of H₂O₂.
7. Pseudo First-Order Reactions
- Reactions that are bimolecular but whose order is experimentally found to be one.
- This happens when one reactant is present in large excess, and its concentration remains practically constant during the reaction.
- Example: Acid hydrolysis of ethyl acetate (ester).
- CH₃COOC₂H₅ + H₂O (excess) --[H⁺]--> CH₃COOH + C₂H₅OH
- Rate = k' [CH₃COOC₂H₅] [H₂O]
- Since [H₂O] is very large and constant, Rate = k [CH₃COOC₂H₅], where k = k'[H₂O].
- The reaction behaves as first order.
- Example: Inversion of cane sugar.
- C₁₂H₂₂O₁₁ + H₂O (excess) --[H⁺]--> C₆H₁₂O₆ (glucose) + C₆H₁₂O₆ (fructose)
- Rate = k [C₁₂H₂₂O₁₁]
8. Temperature Dependence of Reaction Rate - Arrhenius Equation
- Concept: Increasing temperature increases the kinetic energy of molecules, leading to a larger fraction of molecules possessing energy equal to or greater than the activation energy (Ea).
- Activation Energy (Ea): The minimum extra energy that reacting molecules must possess (in addition to their average kinetic energy) to overcome the energy barrier and form products.
- Arrhenius Equation: Quantifies the relationship between rate constant (k) and temperature (T).
- k = A e⁻ᴱᵃ/ᴿᵀ
- Where:
- k = Rate constant
- A = Arrhenius factor or pre-exponential factor or frequency factor (related to collision frequency and orientation)
- Ea = Activation energy (usually in J mol⁻¹ or kJ mol⁻¹)
- R = Gas constant (8.314 J K⁻¹ mol⁻¹)
- T = Absolute temperature (in Kelvin)
- Logarithmic forms:
- ln k = ln A - Ea / RT
- log k = log A - Ea / (2.303 RT)
- Graphical Determination: A plot of log k vs. 1/T gives a straight line with:
- Slope = - Ea / (2.303 R)
- Intercept = log A
- Calculating Ea from rate constants at two temperatures (T₁ and T₂):
- log (k₂ / k₁) = (Ea / 2.303 R) [ (T₂ - T₁) / (T₁ T₂) ]
9. Collision Theory of Chemical Reactions
- Basis: Reactant molecules are assumed to be hard spheres, and reactions occur when these spheres collide.
- Key Postulates:
- The rate of reaction is proportional to the collision frequency (Z) - the number of collisions per second per unit volume.
- Not all collisions lead to product formation. Only effective collisions result in a reaction.
- For a collision to be effective, two conditions must be met:
- Energy Barrier: Colliding molecules must possess a minimum energy called threshold energy (which is related to activation energy). Ea = Threshold Energy - Average KE of reactants.
- Orientation Barrier: Colliding molecules must have proper orientation relative to each other at the time of collision so that old bonds can break and new bonds can form.
- Rate Equation based on Collision Theory: Rate = P * Z<0xE2><0x82><0x9AB> * e⁻ᴱᵃ/ᴿᵀ
- Z<0xE2><0x82><0x9AB> = Collision frequency of reactants A and B.
- e⁻ᴱᵃ/ᴿᵀ = Fraction of molecules with energy ≥ Ea.
- P = Probability or steric factor (accounts for proper orientation). It is related to the Arrhenius factor A (A = P * Z<0xE2><0x82><0x9AB>).
10. Effect of Catalyst
- Definition: A substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change.
- Mechanism:
- Provides an alternative reaction pathway or mechanism with a lower activation energy (Ea).
- It does not alter the Gibbs energy change (ΔG) or equilibrium constant (K) of the reaction. It only helps attain equilibrium faster.
- Forms temporary bonds with reactants, forming an intermediate complex, which then decomposes to give products and regenerate the catalyst.
- Graphical Representation: A catalyst lowers the peak of the energy profile diagram.
11. Relevance to Lab Manual Experiments
- Experiment: Study the effect of concentration on the rate of reaction between Sodium thiosulphate (Na₂S₂O₃) and Hydrochloric acid (HCl).
- Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g) + S(s)
- Rate is studied by monitoring the time taken for a fixed amount of sulphur (S) to precipitate and obscure a mark on paper. Varying [Na₂S₂O₃] while keeping [HCl] and T constant shows rate ∝ [Na₂S₂O₃].
- Experiment: Study the effect of temperature on the rate of reaction between Sodium thiosulphate and HCl.
- Same reaction as above. Time taken for precipitation is measured at different temperatures. Rate increases significantly with temperature (roughly doubles for 10°C rise), allowing calculation of Ea using Arrhenius equation.
- Experiment: Study the rate of reaction between Potassium iodate (KIO₃) and Sodium sulphite (Na₂SO₃) - Iodine Clock Reaction.
- IO₃⁻ + 3SO₃²⁻ → I⁻ + 3SO₄²⁻ (slow) followed by IO₃⁻ + 5I⁻ + 6H⁺ → 3I₂ + 3H₂O (fast) and I₂ + SO₃²⁻ + H₂O → 2I⁻ + SO₄²⁻ + 2H⁺ (very fast, consumes I₂ as it forms). Starch indicator is used. When SO₃²⁻ is consumed, I₂ accumulates and gives blue colour with starch. Time taken for blue colour appearance is measured. Varying initial concentrations helps determine the rate law.
- Experiment: Study the rate of hydrolysis of an ester (e.g., ethyl acetate) catalyzed by an acid.
- CH₃COOC₂H₅ + H₂O --[H⁺]--> CH₃COOH + C₂H₅OH
- Progress is monitored by titrating the acetic acid produced against a standard NaOH solution at different time intervals. Demonstrates pseudo-first-order kinetics. Effect of acid catalyst concentration or temperature can also be studied.
Multiple Choice Questions (MCQs)
-
The unit of rate constant for a zero-order reaction is:
a) s⁻¹
b) L mol⁻¹ s⁻¹
c) mol L⁻¹ s⁻¹
d) L² mol⁻² s⁻¹ -
For a first-order reaction, the half-life period (t½) is related to the rate constant (k) by the expression:
a) t½ = k / 0.693
b) t½ = 0.693 / k
c) t½ = [R]₀ / 2k
d) t½ = k / [R]₀ -
The rate law for the reaction A + 2B → C is found to be Rate = k[A][B]. If the concentration of reactant 'B' is doubled, keeping the concentration of 'A' constant, the value of the rate constant will be:
a) Doubled
b) Halved
c) Remain unchanged
d) Quadrupled -
Which of the following statements is incorrect about the order of a reaction?
a) It is determined experimentally.
b) It can be zero or fractional.
c) It is the sum of powers of concentration terms in the rate law.
d) It is always equal to the sum of stoichiometric coefficients of reactants in the balanced equation. -
The role of a catalyst is to change the:
a) Gibbs energy of reaction
b) Enthalpy of reaction
c) Activation energy of reaction
d) Equilibrium constant of reaction -
Consider the Arrhenius equation k = A e⁻ᴱᵃ/ᴿᵀ. The factor 'A' represents:
a) Activation Energy
b) Fraction of molecules with energy ≥ Ea
c) Pre-exponential factor or Frequency factor
d) Boltzmann constant -
Molecularity of a reaction:
a) Can be determined from the rate law.
b) Can be zero.
c) Is applicable only to elementary reactions.
d) Is always equal to the overall order of reaction. -
Hydrolysis of ethyl acetate in acidic medium is an example of:
a) Zero-order reaction
b) First-order reaction
c) Second-order reaction
d) Pseudo first-order reaction -
For a reaction 2A + B → Products, the rate law is Rate = k[A]²[B]. What is the overall order of the reaction?
a) 1
b) 2
c) 3
d) 0 -
An increase in the temperature of a reaction generally increases the reaction rate primarily because:
a) The activation energy is lowered.
b) The collision frequency increases drastically.
c) The fraction of collisions with sufficient energy increases significantly.
d) The pressure of reactants increases.
Answer Key for MCQs:
- c) mol L⁻¹ s⁻¹
- b) t½ = 0.693 / k
- c) Remain unchanged (Rate constant 'k' depends only on temperature, not concentration)
- d) It is always equal to the sum of stoichiometric coefficients of reactants in the balanced equation.
- c) Activation energy of reaction
- c) Pre-exponential factor or Frequency factor
- c) Is applicable only to elementary reactions.
- d) Pseudo first-order reaction
- c) 3 (Order = 2 + 1 = 3)
- c) The fraction of collisions with sufficient energy increases significantly.
Make sure you understand these concepts thoroughly. Refer back to your NCERT textbook and lab manual for further details and experimental procedures. Good luck with your preparation!