Summary of Chemical Kinetics

General Chemistry Overview

  • Focus on Chemical Kinetics: study of reaction rates.

Kinetics vs Thermodynamics

  • Kinetics: speed of reactions, mechanisms.
  • Thermodynamics: energetics, favorability of products vs reactants.

Reaction Rate

  • Defined as Rate=−Δ[reactant]Δt\text{Rate} = -\frac{\Delta [\text{reactant}]}{\Delta t}
  • Instantaneous vs Average Rate. Instantaneous: slope of tangent. Average: over a time interval.

Reaction Order

  • Defined by the powers in the rate law, sums to overall order.
  • Example: Rate=k[A]n[B]m\text{Rate} = k[A]^n[B]^m
  • Experimental determination needed. Types: Zero (no change), First (linear), Second (quadratic).

Integrated Rate Laws

  • Zero Order: [A]<em>t=−kt+[A]</em>0[A]<em>t = -kt + [A]</em>0
  • First Order: ln⁡[A]<em>t=−kt+ln⁡[A]</em>0\ln[A]<em>t = -kt + \ln[A]</em>0
  • Second Order: 1[A]<em>t=kt+1[A]</em>0\frac{1}{[A]<em>t} = kt + \frac{1}{[A]</em>0}

Half-Life Expressions

  • Zero: t<em>1/2=[A]</em>02kt<em>{1/2} = \frac{[A]</em>0}{2k}
  • First: independent of concentration. t1/2=ln⁡(2)kt_{1/2} = \frac{\ln(2)}{k}
  • Second: t<em>1/2=1k[A]</em>0t<em>{1/2} = \frac{1}{k[ A]</em>0}

Activation Energy & Temperature

  • Kinetics influenced by collision theory and activation energy.
  • Arrhenius equation: k=Ae−EaRTk = Ae^{-\frac{E_a}{RT}}
  • Higher temperatures increase collision energy and reaction rates.

Reaction Mechanisms

  • Steps describing how reactions occur at a molecular level. Determined by:
    • Elementary steps cannot be simplified further.
    • Rate-determining step governs the overall rate law.
  • Catalysts provide alternative pathways, lowering activation energy without being consumed.