Notes from CHE2FCB Lecture on Catalysis

Catalysis Overview

• Catalysts speed up reactions without undergoing permanent changes.
• They provide alternative pathways with lower activation energy.

Types of Catalysis

• Homogeneous Catalysis: Same phase as reactants (e.g., gas-phase catalysts).
• Heterogeneous Catalysis: Different phase, typically solid catalysts where reactions occur on surfaces (e.g., catalytic converters).
• Enzymes: Biological catalysts that enhance reaction rates significantly.

Key Concepts

• Activation Energy (Ea): Energy required to initiate a reaction. Catalysts lower Ea.
• Turnover Frequency (kcat): Number of substrate molecules converted to product per enzyme molecule per second.
• Michaelis-Menten Mechanism: Describes enzyme kinetics:
  $S + E ⇌ ES → P + E$
  $K<em>M = \frac{k</em>{-1} + k<em>{cat}}{k</em>1}$

Reaction Pathway

• Catalyzed reactions have a lower activation energy compared to uncatalyzed reactions.
• The graph of energy vs. reaction coordinate displays reactants, products, and transition states (TS).

Rate Laws and Catalysts

• Rate laws incorporate catalysts but typically not intermediates.
• Example: For bromide-catalyzed decomposition of peroxide:
  $ext{Rate} = k[H<em>2O</em>2][H_3O^+][Br^-]$

Autocatalysis

• Reaction where the product catalyzes its own formation leading to an increase in reaction rate.

Negative Catalysts (Inhibitors)

• Slow down reactions, e.g., Tetraethyl lead (TEL).
• Work by changing the pathway or removing reactive intermediates instead of just providing an alternative pathway.

Summary

• Catalysts do not alter the equilibrium constant (K). They accelerate both forward and reverse reactions equally.
• The overall reaction remains unchanged, only the rate changes due to the catalyst present.

Key Equation

• Enzyme kinetics from Michaelis-Menten:
  $<br />\nu<em>{max} = \frac{k</em>{cat} [E]}{K_M + [S]}$