1. Concentration of Reactants

Higher concentration → more collisions per unit time → higher rate. This is expressed quantitatively in the rate law: rate = k[A]^m[B]^n.

2. Temperature

A 10°C rise approximately doubles or triples the rate (temperature coefficient ≈ 2–3). Quantitatively expressed by the Arrhenius equation. Higher T → greater fraction of molecules with energy ≥ Ea.

3. Pressure (for gaseous reactions)

Higher pressure → higher concentration of gas molecules → more collisions → higher rate. Only significant for gas-phase reactions.

4. Catalyst

A catalyst provides an alternative reaction pathway with lower activation energy. Increases rate of both forward and backward reactions equally. Does NOT change equilibrium constant K or $\Delta H$.

5. Nature of Reactants

Reactions involving bond breaking require energy (Ea). Ionic reactions in solution (e.g., acid-base) are very fast (near zero Ea); reactions requiring bond breaking (e.g., covalent) are slower.

6. Surface Area

For heterogeneous reactions (solid-liquid or solid-gas), finely divided solids expose more surface area to the reactant → higher rate. Example: powdered Mg burns faster than Mg ribbon.

7. Light (Photochemical Reactions)

Light provides energy to overcome activation barrier. Examples: photosynthesis, photography (AgBr decomposition), $H_{2}$ + $Cl_{2}$ in light (explosive).