📘 Chemical Kinetics (Full Study Guide)Class XII CBSE Syllabus Comprehensive Summary

1. Introduction & Rate• Kinetics: Studies 'How Fast' (Speed)• Rate: Change in [Conc] / TimeRate = -1/a Δ[A]/Δt = 1/c Δ[C]/Δt(-) Reactant decrease | (+) Product increase• Unit: mol L⁻¹ s⁻¹

2. Factors Affecting Rate• Conc: Higher [Conc] → Faster Rate• Temp: 10°C rise ≈ 2x Rate increase• Catalyst: Lowers Ea (Activation Energy)• Surface Area: Powdered > Lump• Nature: Ionic (Fast) vs Covalent (Slow)

Order vs Molecularity<table style="width: 100%; border-collapse: collapse;"><tr><th style="border-bottom: 1px solid #CBD5E0">Feature</th><th style="border-bottom: 1px solid #CBD5E0">Order (Experimental)</th><th style="border-bottom: 1px solid #CBD5E0">Molecularity (Theoretical)</th></tr><tr><td>Value</td><td>0, fraction, negative possible</td><td>Always whole number (>0)</td></tr><tr><td>Applicability</td><td>Overall reaction</td><td>Elementary reactions only</td></tr></table>

Integrated Rate Law Summary<table style="width: 100%; border-collapse: collapse;"><tr style="background: #EDF2F7"><th>Order</th><th>Rate Law</th><th>Integrated Equation</th><th>Half-life (t₁/₂)</th><th>k Units</th></tr><tr><td>0 Order</td><td>Rate = k</td><td>[A] = [A]₀ - kt</td><td>[A]₀ / 2k</td><td>mol L⁻¹s⁻¹</td></tr><tr><td>1st Order</td><td>Rate = k[A]</td><td>k = (2.303/t) log([A]₀/[A])</td><td>0.693 / k</td><td>s⁻¹</td></tr><tr><td>2nd Order</td><td>Rate = k[A]²</td><td>1/[A] = kt + 1/[A]₀</td><td>1 / k[A]₀</td><td>L mol⁻¹s⁻¹</td></tr></table>• Graph Tips: - Zero Order: [A] vs t (Linear) - 1st Order: log[A] vs t (Linear) - 2nd Order: 1/[A] vs t (Linear)

4. Temperature & Arrhenius Equationk = A e^(-Ea/RT) → log k = log A - Ea / (2.303RT)• Graph: Plot log k vs 1/T yields a straight line.• Slope: -Ea / 2.303R• Collision Theory: Effective collisions = Energy (≥ Ea) + Proper Orientation.• Catalyst: Lowers peak of Energy Profile Diagram, speeds up both directions.

5. Quick Examples & PYQs• Pseudo First Order: Ester hydrolysis in excess H₂O (Rate depends on Ester only).• Zero Order: Decomposition of NH₃ on Pt catalyst.• Radioactive Decay: All follow First Order kinetics.• Note: 1st order half-life is independent of initial concentration!

Feature	Order (Experimental)	Molecularity (Theoretical)
Value	0, fraction, negative possible	Always whole number (>0)
Applicability	Overall reaction	Elementary reactions only

Order	Rate Law	Integrated Equation	Half-life (t₁/₂)	k Units
0 Order	Rate = k	[A] = [A]₀ - kt	[A]₀ / 2k	mol L⁻¹s⁻¹
1st Order	Rate = k[A]	k = (2.303/t) log([A]₀/[A])	0.693 / k	s⁻¹
2nd Order	Rate = k[A]²	1/[A] = kt + 1/[A]₀	1 / k[A]₀	L mol⁻¹s⁻¹

lkps