Reaction Rates

Reaction rates measure how quickly reactants turn into products in a chemical reaction, a core concept in kinetics. For example, rusting (\( \ce{4Fe + 3O2 -> 2Fe2O3} \)) is slow, while combustion is fast. Rates are expressed as the change in concentration over time:

\[ \text{Rate} = -\frac{\Delta [\ce{A}]}{\Delta t} = \frac{\Delta [\ce{B}]}{\Delta t} \]

This article explores factors influencing rates, mathematical descriptions, and their significance.

Factors Affecting Rates

Several factors alter reaction speed:

  • Concentration: Higher reactant concentrations increase collision frequency (e.g., \( \ce{H2 + I2 -> 2HI} \)).
  • Temperature: Raises kinetic energy, speeding reactions (Arrhenius effect).
  • Catalysts: Lower activation energy, e.g., enzymes in biology.
  • Surface Area: More area (e.g., powdered vs. solid) speeds heterogenous reactions.

Rate Laws

Rate laws relate rate to reactant concentrations:

\[ \text{Rate} = k [\ce{A}]^m [\ce{B}]^n \]

\( k \): rate constant; \( m, n \): reaction orders (determined experimentally). Example: \( \ce{2N2O5 -> 4NO2 + O2} \), Rate = \( k [\ce{N2O5}] \) (first-order).

Problem: If \( k = 0.1 \, \text{s}^{-1} \) and \( [\ce{N2O5}] = 0.2 \, \text{M} \), Rate = \( 0.1 \times 0.2 = 0.02 \, \text{M/s} \).

Activation Energy

Activation energy (\( E_a \)) is the energy barrier reactants must overcome. The Arrhenius equation relates \( k \) to \( E_a \):

\[ k = A e^{-\frac{E_a}{RT}} \]

\( A \): pre-exponential factor; \( R \): gas constant; \( T \): temperature. Catalysts reduce \( E_a \), speeding reactions without being consumed.

Applications

Kinetics is key in:

  • Industry: Optimizing reaction conditions (e.g., Haber process).
  • Medicine: Drug decomposition rates.
  • Environment: Ozone depletion rates (\( \ce{O3 + Cl -> O2 + ClO} \)).

It informs process efficiency and safety.