Chemical kinetics — Eduqas A-Level Chemistry

Chemical kinetics is the study of reaction rates and the factors that influence them. For WJEC A-Level Chemistry, this topic explores how fast reactions occur, why some reactions are instantaneous while others take years, and how we can control reaction speed. Understanding kinetics is essential for predicting reaction behaviour in industrial processes, such as the Haber process for ammonia production, where optimising rate is critical for efficiency.

The topic builds on GCSE ideas of collision theory and activation energy, but introduces quantitative methods like rate equations, orders of reaction, and the Arrhenius equation. You'll learn to determine rate laws from experimental data, interpret concentration-time and rate-concentration graphs, and explain how temperature, concentration, and catalysts affect rate at a molecular level. This knowledge is foundational for further study in chemical equilibrium and reaction mechanisms.

Kinetics also connects to real-world applications: from drug metabolism in the body to the degradation of materials. Mastering this topic requires a blend of mathematical manipulation and conceptual understanding, making it a key area for exam success. It typically appears in both multiple-choice and long-answer questions, often requiring graph analysis and calculation.

• →Rate of reaction: defined as change in concentration per unit time (mol dm⁻³ s⁻¹), measured using initial rates or continuous monitoring.
• →Orders of reaction: zero, first, and second order; determined from rate-concentration graphs or initial rate data. The overall order is the sum of individual orders.
• →Rate equation: rate = k[A]ᵐ[B]ⁿ, where k is the rate constant (units vary with order). The rate constant is temperature-dependent via the Arrhenius equation.
• →Arrhenius equation: k = Ae⁻ᴱᵃ/ᴿᵀ, linking rate constant to activation energy (Ea) and temperature. A plot of ln k against 1/T gives a straight line with slope -Ea/R.
• →Catalysts: provide an alternative pathway with lower activation energy, increasing rate without being consumed. Homogeneous and heterogeneous catalysts are distinguished.

• Principles of measuring reaction rate by sampling and quenching
• Determination of reaction order from experimental results
• Application of the general rate equation
• Concept and identification of the rate-determining step
• Link between reaction kinetics and reaction mechanism
• Use of the Arrhenius equation to find activation energy and frequency factor
• Effect of temperature and catalysts on the rate constant

• 💡Ensure you can derive the units for the rate constant k for any given order of reaction
• 💡Practice calculating activation energy from the Arrhenius equation using logarithmic forms
• 💡Be prepared to interpret concentration-time and rate-concentration graphs to determine reaction order
• 💡Always link the rate equation to the mechanism: the rate-determining step involves the species present in the rate equation
• 💡When determining orders from initial rate data, compare experiments where only one concentration changes. Use the ratio of rates to find the order (e.g., if rate doubles when concentration doubles, it's first order).
• 💡For Arrhenius plots, ensure you plot ln k (y-axis) against 1/T (x-axis) in Kelvin. The gradient is -Ea/R, so Ea = -gradient × R (8.314 J mol⁻¹ K⁻¹). Watch units: Ea in J mol⁻¹, not kJ.
• 💡In long-answer questions, always define the rate in terms of a reactant or product. Use the correct units for rate (mol dm⁻³ s⁻¹) and for k (e.g., s⁻¹ for first order, dm³ mol⁻¹ s⁻¹ for second order).

• Confusing the rate-determining step with the overall reaction stoichiometry
• Incorrectly applying units to rate constants for different orders of reaction
• Misinterpreting the Arrhenius plot (slope and intercept)
• Failing to correctly identify the rate-determining step from a proposed mechanism
• Misconception: The rate constant k changes with concentration. Correction: k is constant at a fixed temperature; only temperature (and catalysts) alter k.
• Misconception: Doubling concentration always doubles the rate. Correction: This is only true for first-order reactions. For second-order, doubling concentration quadruples the rate; for zero-order, rate is unchanged.
• Misconception: Activation energy is the energy barrier that must be overcome for a reaction to occur. Correction: Activation energy is the minimum energy required for a successful collision; molecules must collide with energy ≥ Ea and correct orientation.