Reaction Rate Concepts

Overview

This lecture covers key concepts and methods for measuring and calculating reaction rates, determining rate equations, and understanding the Arrhenius equation for the OCR A-Level Chemistry "How Fast" (Kinetics) topic.

Measuring Rate of Reaction

• Rate can be measured by changes in pH, mass loss, gas volume, or color change during a reaction.
• Use a pH meter for reactions involving H+ ions, and a top pan balance or gas syringe for gaseous products.
• A colorimeter measures the absorbance of colored species to track concentration changes.

Rate Calculations and Graphs

• The rate of reaction is found from the gradient (change in y/change in x) of a concentration vs. time graph.
• For curves, draw a tangent to find instantaneous rate at a specific time.
• The initial rate is the gradient at t=0 (start of the reaction).
• Units for rate are typically (amount of product/reactant)/(time).

Clock Experiments & Initial Rate Method

• Clock reactions time until a visible endpoint (e.g., color change) occurs, estimating the initial rate.
• Three assumptions: constant temperature, negligible change in reactant concentration, and endpoint is not too far into the reaction.
• The iodine clock uses sodium thiosulfate and starch to indicate the endpoint via a sudden color change.

Rate Equations & Orders of Reaction

• Rate equation: rate = k[A]^m[B]^n, where m and n are reaction orders.
• Zero order: concentration change has no effect on rate.
• First order: rate ∝ concentration.
• Second order: rate ∝ concentration squared.
• Reaction orders must be determined experimentally.

Calculating and Rearranging Rate Equations

• To solve for rate, substitute values into the rate equation.
• To find k (rate constant), rearrange: k = rate / ([A]^m[B]^n).
• Units for k vary with overall reaction order.

Graphs and Orders from Experiment

• Zero order: flat line on rate vs. concentration graph.
• First order: straight diagonal line.
• Second order: upward curve.
• Half-life for first-order reactions is constant and used to calculate k: k = ln(2) / t₁/₂.

Deducing Rate Equations from Data

• Change concentrations of reactants one at a time and observe how the initial rate changes.
• Use proportional changes to determine reaction order with respect to each reactant.
• For complex tables, add columns to compare theoretical vs. actual rates.

Rate-Determining Step and Mechanisms

• The slowest step in a multi-step reaction is the rate-determining step.
• Only species in the rate equation appear in the rate-determining step.
• Catalysts can appear in the rate equation if they influence the rate.
• Use the stoichiometry and experimental data to deduce which step is rate-determining.

Arrhenius Equation

• Arrhenius equation: k = A * e^(-Ea/RT).
• Higher temperature or lower activation energy increases k (rate constant).
• Use natural logarithms to linearize and plot ln(k) vs 1/T; gradient = -Ea/R.
• Calculate activation energy using the rearranged (ln) form of Arrhenius equation.*

Key Terms & Definitions

• Rate of Reaction — Speed at which reactants convert to products.
• Colorimeter — Device measuring absorbance of colored solutions.
• Initial Rate — Rate right at the start (t=0) of a reaction.
• Order of Reaction — Power to which reactant concentration is raised in the rate equation.
• Rate Constant (k) — Proportionality constant in rate equations; depends on temperature.
• Rate-Determining Step — Slowest step in a multi-step reaction.
• Arrhenius Equation — Relates rate constant to activation energy and temperature.
• Half-life (t₁/₂) — Time for reactant concentration to halve in a first-order reaction.

Action Items / Next Steps

• Practice drawing and analyzing concentration-time and rate-concentration graphs.
• Complete calculations using the rate law and Arrhenius equations.
• Review and attempt iodine clock and other clock reaction experiments.
• Memorize and apply key terms and formulae.