Topic 11: Kinetics

Rates and Rate Equations

• Reaction Rate: The speed at which reactants convert into products.
  • Depends on reactant concentrations and the rate constant.
  • Described by the rate equation: [ A + B \rightarrow C ]
  • [ \text{Rate} = k[A]^m[B]^n ]
• Order of Reaction:
  • Constants (m) and (n) represent reaction order for each reactant.
  • Can be zero, first, or second order.
  • Total order: Sum of individual orders (m + n).

Rate Constant (k)

• Constant at a fixed temperature.
• Relates concentrations affecting the rate.
• Calculation: Rearrange the rate equation.
• Unit varies with reaction orders.

Orders of Reaction

• Zero Order:
  • Rate = k
  • Concentration does not affect the rate.
  • Horizontal line on rate-concentration graph.
• First Order:
  • Rate (= k[A])
  • Directly proportional to concentration.
  • Doubling concentration doubles the rate.
• Second Order:
  • Rate (= k[A]^2)
  • Rate proportional to concentration squared.
  • Doubling concentration quadruples the rate.

Concentration-Time Graphs

• Used to deduce reaction orders.
• Zero Order: Linear graph.
• First/Second Order: Curved graph.

Initial Rates Method

• Determines reaction order by varying reactant concentrations.
• Effects of doubling concentration:
  • Zero Order: No change.
  • First Order: Rate doubles.
  • Second Order: Rate quadruples.
• Example:
  • Given data deduces orders of reactants A, B, C.

Half-life

• Definition: Time for concentration to halve.
• First Order: Constant half-life throughout reaction.

Experimental Techniques

• Measuring change in reactant/product mass/concentration over time.
• Mass Change: Decrease in gas-producing reactions.
• Volume of Gas Evolved: Use gas syringe to measure volume/time.
• Titration: Determining concentration at regular intervals.
• Colorimetry: Uses light absorption to measure concentration.
  • Calibration curve used for concentration-time graph.

The Rate Determining Step

• Slowest step determines overall reaction rate.
• Rate equation includes species in the rate-determining step.
• Helps predict reaction mechanism.

The Arrhenius Equation

• Relates rate constant and temperature exponentially.
• Graphical Form: ln(k) vs. 1/T; gradient = -Ea/R.

SN1 and SN2 Reactions

• SN1: Two-step; rate depends on halogenoalkane.
  • Tertiary halogenoalkanes fastest due to carbocation stability.
• SN2: Single step; rate depends on both reactants.
  • Primary halogenoalkanes fastest due to less steric hindrance.

Catalysts

• Lower activation energy by providing alternative pathways.
• Homogeneous Catalysts: Same phase as reactants.
  • Example: Iron ions in redox reactions.
• Heterogeneous Catalysts: Different phase.
  • Example: Iron in the Haber process.
• Adsorption: Involves reactants binding to catalyst surface, increasing reaction rate.