Chemical Kinetics

Chemical Kinetics (Introduction)

Study of rate of chemical reactions

Deals with:

Speed of reaction

Mechanism (steps involved)

Factors affecting rate:

Concentration

Temperature

Catalyst

Surface area

Rate Equation (Rate Law)

General form: Rate=k[A]^m[B]^n

k = rate constant

m, n = order w.r.t reactants

Determined experimentally

Not from balanced equation

Order of Reaction

Sum of powers in rate law: Order=m+n

Zero order → rate independent of concentration

First order → rate ∝ concentration

Second order → rate ∝ square

Integral Rate Equations

Zero Order: [A]=[A]₀−kt

First Order: ln[A]=ln[A]₀−kt

Second Order: 1/[A]=1/[A]₀+kt

Half-Life of Reaction (t₁/₂)

Zero Order: t₁/₂=[A]₀/2k

First Order: t₁/₂=0.693/k

Second Order: t₁/₂=1/k[A]₀

First order half-life is independent of concentration

Molecularity

Number of molecules involved in single elementary step

Always whole number (1, 2, 3)

Never zero or fractional

Unimolecular → 1 molecule

Bimolecular → 2 molecules

Termolecular → 3 molecules

Order → experimental

Molecularity → theoretical (mechanism-based)

Collision Theory

Reaction occurs when molecules collide effectively

Conditions:

Proper orientation

Sufficient energy (≥ activation energy)

Rate ∝ Number of effective collisions

Effect of Temperature

Increase in temperature → increase in rate

Arrhenius Equation: k=Ae−Ea/RT

Higher T → more molecules cross activation energy

Rate roughly doubles for every 10°C rise

Effect of Catalyst

Catalyst lowers activation energy

Increases rate

Does NOT change equilibrium

Does NOT get consumed

Provides alternative reaction pathway

Numericals (How to Solve)

Find rate constant (k) using rate law or integrated equations

Find order by comparing experiments: R1/R2=([A]1/[A]2)^m

Half-life problems: Direct formula substitution

Time-based concentration: Use integrated equations