MCAT General Chemistry · Lesson 5

Chemical Kinetics

3 min read2 sectionsUpdated January 3, 2026

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5.1Chemical Kinetics
5.2Reaction Rates

2 sections

5.1 Chemical Kinetics

This section explains how reaction rates are determined, the concept of reaction mechanisms, transition states, collision theory, and factors affecting reaction rates including temperature, concentration, medium, and catalysts.

∆G and Spontaneity
- ∆G determines if a reaction is spontaneous.
- Even if $-\Delta G$ , reaction may not proceed quickly.
Reaction Mechanisms
- Most reactions occur in multiple steps; the sum of steps = overall reaction.
- Example: $A_2 + 2B → 2AB$ $A_{2} + 2 B \to 2 A B$
  - Step 1: $A_2 + B → A_2B$ (slow, rate-determining)
  - Step 2: $A_2B + B → 2AB$ (fast)
- Intermediate: species formed then consumed (e.g., A₂B).
- Rate-determining step: bottleneck, controls reaction speed.
Molecular Basis
- Collision theory: rate ∝ number of collisions/sec.
- Effective collisions require correct orientation & energy.
- Minimum energy required = activation energy.
- $\text{Rate} = Z × f$ $Rate = Z \times f$
  - Z = total collisions/sec
  - f = fraction of effective collisions
Arrhenius Equation
- $k = Ae^{-E_a/RT}$ $k = A e^{- E_{a} / RT}$
  - A: frequency factor (1/s)
  - $E_a$ : activation energy
  - R: gas constant
  - T: temperature
- Raising T raises rate (smaller exponent).

Khan Academy Video: Arrhenius Equation

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Transition State Theory
- Molecules form a high-energy transition state.
- Old bonds weaken, new bonds form along reaction coordinate.
- Activation energy = energy needed to reach transition state.
Theoretical Constructs
- Free energy change $\Delta G$ = energy difference between products & reactants.
- Example: $H_2 + Cl_2 → 2HCl$ (exergonic).
Factors Affecting Reaction Rate
- Concentration: ↑[reactants] → ↑collisions → faster rate.
- Temperature: ↑T → ↑rate (roughly doubles per 10°C); very high T can denature reactants.
- Medium: Polar solvents stabilize dipoles, weaken bonds, favor reactions.
- Catalysts: Increase rate without being consumed.
  - Homogeneous: same phase as reactants.
  - Heterogeneous: different phase.
  - Lower activation energy for forward and reverse reactions.

5.2 Reaction Rates

Describes how to measure reaction rates, determine rate laws, reaction orders, and analyze complex or mixed-order reactions.

Khan Academy Video: Reaction Rates

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Definition of Rate
- Rate measured as disappearance of reactants or appearance of products:
  - Reactant A: $-\frac{\Delta [A]}{\Delta t}$
  - Product C: $\frac{\Delta [C]}{\Delta t}$
- Example: $2A + B → C$ (2 moles A consumed per 1 mole B)
Rate Law
- $\text{Rate} = k[A]^x[B]^y$ $Rate = k [A]^{x} [B]^{y}$
  - k = rate constant
  - x, y = experimentally determined orders (not stoichiometric coefficients)
- Do not confuse with Keq:
  - $K_{eq} = \frac{k_{forward}}{k_{reverse}}$
Experimental Determination
- Vary one reactant, keep others constant → observe rate changes to determine exponents.
Reaction Orders
- Zero Order:
  - $\text{Rate} = k[A]^0[B]^0 = k$
  - K units: M/s
  - Independent of [A]; slope = k.
- First Order:
  - $\text{Rate} = k[A/B]^1$
  - Example: radioactive decay
    - $[A]_t = [A]_0 e^{-kt}$
  - Plot [A] → curved; ln[A] → straight line.
- Second Order:
  - $\text{Rate} = k[A]^1[B]^1$ or $\text{Rate} = k[A/B]^2$
  - Suggests collisions between reactants.
  - Plot [A] → nonlinear; 1/[A] → straight line, slope = k.
- Higher-Order Reactions:
  - Rare for 3+ particles to collide simultaneously with correct orientation & energy.
- Mixed/Non-Integer Orders:
  - Rate orders can change over time; described as broken/mixed order.
  - Example: $\text{Rate} = \frac{k_1[C][A]^2}{k_2 + k_3[A]}$ $Rate = \frac{k _{1} [ C ] [ A ] ^{2}}{k _{2} + k _{3} [ A ]}$
    - Early: $k_3[A] >> k_2$ → first-order in A
    - Late: $k_2 >> k_3[A]$ → second-order in A

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