Reaction rates and collision reasoning

Reaction Kinetics & Rates

Chemical kinetics studies how fast reactions proceed, rather than just thermodynamic favourability.

Rate law: $v = k\,[A]^{m}\,[B]^{n}$

Collision theory: Reactions require molecules to collide with sufficient kinetic energy $\ge E_a$ (activation energy) and correct orientation.

Arrhenius & Catalysts

Arrhenius equation: $k = A\,e^{-E_a / RT}$

• $A$: frequency factor
• $E_a$: activation energy
• $R$: $8.314\text{ J mol}^{-1}\text{K}^{-1}$

Linear form: $\ln k = \ln A - \frac{E_a}{R}\cdot\frac{1}{T}$. Plotting $\ln k$ vs $1/T$ gives slope $-E_a/R$.

Catalysts lower $E_a$ without being consumed.

Key insight: Doubling temperature dramatically increases rates due to exponential dependence.

Common pitfall: A catalyst speeds the journey but never moves the destination. It lowers forward and reverse barriers equally, changing rate, not equilibrium.