# Kinetics

> Edexcel International A-Level Chemistry · IAL Chemistry Unit 4 (WCH14)
> Source: https://www.owlsprep.com/study/edexcel-ial-chemistry-u4-kinetics/

This guide covers all Edexcel IAL Chemistry Unit 4 Kinetics content, from core term definitions to experimental technique selection, rate equation deduction, mechanism analysis and Arrhenius calculation skills required for exam success.

**Prerequisites:** [Basic collision theory from Unit 2 Reaction Rates](https://www.owlsprep.com/study/edexcel-ial-chemistry-u2-reaction-rates/); [Foundational organic reaction mechanisms from Unit 3](https://www.owlsprep.com/study/edexcel-ial-chemistry-u3-organic-mechanisms/)

## Learning objectives

- Define core kinetics terms including rate, order, rate constant, half-life and RDS
- Deduce reaction order from concentration-time, rate-concentration and initial rate experimental data
- Relate rate equations to reaction mechanisms and identify the rate-determining step
- Explain the difference between SN1 and SN2 halogenoalkane hydrolysis using rate evidence
- Calculate activation energy from Arrhenius graph data
- Select appropriate experimental techniques for rate measurement for different reaction types

## Core Kinetics Terms & Rate Equation Fundamentals

**Rate Equation** — Mathematical relationship between reaction rate and reactant concentrations, where m and n are experimentally determined orders of reaction with respect to each reactant, independent of reaction stoichiometry.

*Notation:* rate = k[A]^m[B]^n

The overall order of a reaction is the sum of the individual orders (m + n). Units of the rate constant k are derived by rearranging the rate equation and substituting units for rate (mol dm⁻³ s⁻¹) and concentration (mol dm⁻³). For first order reactions, k has units s⁻¹; for second order reactions, k has units mol⁻¹ dm³ s⁻¹.

> **warning**
>
> Never take reaction orders directly from the balanced stoichiometric equation unless explicitly told the reaction is single-step. Orders are always derived from experimental data.

**Worked example:** A reaction has the rate equation $rate = k[X][Y]^2$. What is the order with respect to X, Y, and overall order? State the units of k.

1. Order with respect to X is 1, order with respect to Y is 2. Overall order = 1 + 2 = 3.
2. Rearrange the rate equation to solve for k: $k = \frac{rate}{[X][Y]^2}$
3. Substitute units: $\frac{mol\ dm^{-3}\ s^{-1}}{(mol\ dm^{-3}) \times (mol\ dm^{-3})^2} = mol^{-2}\ dm^6\ s^{-1}$

## Determining Reaction Order from Experimental Data

Reaction order is determined using two experimental approaches: continuous monitoring (recording concentration of a reactant or product over time to produce concentration-time graphs) and initial rate methods (measuring rate at the start of the reaction for different starting concentrations, often using clock reactions).

**Half-life** — Time taken for the concentration of a reactant to reduce to half its initial value. A constant half-life across the reaction confirms a first order reaction.

- 0 order: Concentration-time graph is a straight line with negative gradient, rate is constant regardless of reactant concentration
- 1st order: Concentration-time graph is an exponential curve, half-life is constant
- 2nd order: Concentration-time graph is a steep exponential curve, half-life increases as concentration falls

**Worked example:** Concentration of reactant A is measured at time intervals: t=0s [A]=0.80 mol dm⁻³, t=20s [A]=0.40 mol dm⁻³, t=40s [A]=0.20 mol dm⁻³, t=60s [A]=0.10 mol dm⁻³. Deduce the order of reaction with respect to A.

1. Calculate first half-life: time for [A] to fall from 0.80 to 0.40 mol dm⁻³ = 20s
2. Calculate second half-life: time for [A] to fall from 0.40 to 0.20 mol dm⁻³ = 20s
3. Half-life is constant across the reaction, so the reaction is first order with respect to A.

## Rate Equations and Reaction Mechanisms

**Rate-Determining Step (RDS)** — The slowest step in a multi-step reaction mechanism. The species and their stoichiometric counts in the RDS exactly match the orders in the experimentally determined rate equation.

**Worked example:** The rate equation for acid-catalysed iodination of propanone is $rate = k[CH_3COCH_3][H^+]$. Deduce which species are in the RDS, and state the order with respect to I₂.

1. Species present in the rate equation are propanone and H⁺ ions, so both appear in the RDS with order 1 each.
2. I₂ is not present in the rate equation, so the order with respect to I₂ is 0, meaning it only participates in fast steps after the RDS.

Rate equation evidence is used to distinguish between SN1 and SN2 mechanisms for halogenoalkane hydrolysis: 3° halogenoalkanes have a rate equation $rate = k[RX]$, so the RDS only involves the halogenoalkane (SN1 mechanism via a carbocation intermediate). 1° halogenoalkanes have a rate equation $rate = k[RX][OH^-]$, so the RDS involves both reactants (SN2 concerted mechanism).

## Arrhenius Equation and Activation Energy Calculation

$$ln\thinspace k = ln\thinspace A - \frac{E_a}{RT}$$

Plotting a graph of $ln\thinspace k$ on the y-axis against $1/T$ (in K⁻¹) on the x-axis gives a straight line with gradient equal to $-E_a/R$. R is given as 8.31 J mol⁻¹ K⁻¹ in the data booklet.

> **tip**
>
> Always convert temperature from Celsius to Kelvin (add 273) before calculating 1/T values, and convert your final Ea value from J mol⁻¹ to kJ mol⁻¹ as standard for exam answers.

**Worked example:** An Arrhenius plot of $ln\thinspace k$ against $1/T$ gives a gradient of -6000 K. Calculate the activation energy Ea in kJ mol⁻¹.

1. Use the relationship $gradient = -E_a / R$, rearranged to $E_a = -gradient \times R$
2. Substitute values: $E_a = -(-6000\ K) \times 8.31\ J\ mol^{-1}\ K^{-1} = 49860\ J\ mol^{-1}$
3. Convert to kJ mol⁻¹: $49860 / 1000 = 49.9\ kJ\ mol^{-1}$ (3 significant figures)

## Experimental Techniques for Rate Measurement

- Titration: Quench reaction at intervals then titrate remaining reactant, suitable for reactions with acid/base or redox reactants
- Colorimetry: Measure absorbance of coloured species (e.g. I₂) over time, non-invasive and continuous
- Mass change: Measure mass of reaction mixture as gaseous product escapes, suitable for reactions producing dense gases like CO₂
- Gas volume: Collect gaseous product in a gas syringe over time, suitable for reactions producing any insoluble gas

**Worked example:** Justify the most appropriate technique to measure the rate of reaction between calcium carbonate and hydrochloric acid, which produces CO₂ gas.

1. The reaction produces a gaseous product (CO₂) that leaves the reaction mixture.
2. Gas volume measurement via a gas syringe is appropriate, as the volume of CO₂ produced per unit time directly correlates to reaction rate.
3. Mass change measurement is also valid, as the system loses mass as CO₂ escapes to the atmosphere.

## Common pitfalls

- **Wrong:** Taking reaction order directly from the balanced stoichiometric equation
  - Why it fails: Orders are experimentally determined, stoichiometry reflects the overall reaction not the individual mechanism steps
  - Correct: Only use experimental data (graphs, initial rate tables) to deduce reaction order
- **Wrong:** Using Celsius temperature in Arrhenius 1/T calculations
  - Why it fails: The gas constant R uses absolute temperature in Kelvin, so Celsius values give incorrect gradient and Ea results
  - Correct: Always add 273 to Celsius temperature to convert to Kelvin before calculating 1/T values
- **Wrong:** Assuming I₂ is in the RDS for iodination of propanone
  - Why it fails: The rate equation does not include I₂, so it is not involved in the slow step of the reaction
  - Correct: Match species in the rate equation exactly to species in the RDS, with counts matching the order values
- **Wrong:** Mixing up SN1 and SN2 rate equations
  - Why it fails: SN1 has only the halogenoalkane in the rate equation (first order) while SN2 has both halogenoalkane and nucleophile (second order)
  - Correct: If rate depends only on [RX] it is SN1, if it depends on both [RX] and [Nu] it is SN2
- **Wrong:** Using incorrect units for the rate constant k
  - Why it fails: Units of k depend directly on the overall order of the reaction, using wrong units loses exam marks
  - Correct: Rearrange the rate equation to solve for k, substitute units of rate and concentration to derive k units for each reaction

## Cheatsheet

| Concept | Key Rule | Exam Check |
| --- | --- | --- |
| Reaction Order | 0: rate independent of [A]; 1: constant $t_{1/2}$; 2: rate ∝ [A]² | Confirm order only from experimental data, not stoichiometry |
| Rate Constant Units | 1st order: s⁻¹; 2nd order: mol⁻¹ dm³ s⁻¹; derive for higher orders | Always derive units by rearranging the rate equation |
| RDS Matching | Species in RDS = species in rate equation; count = order | Any species not in rate equation appears only in fast steps after RDS |
| SN1 vs SN2 | SN1: $rate = k[RX]$ (3° haloalkanes); SN2: $rate = k[RX][Nu]$ (1° haloalkanes) | Carbocation intermediate for SN1, concerted mechanism for SN2 |
| Arrhenius Calculation | Gradient of $ln\thinspace k$ vs $1/T$ = $-E_a/R$; $R = 8.31\ J\ mol^{-1}\ K^{-1}$ | Convert Ea from J to kJ mol⁻¹ for final answer, use Kelvin for T |

## What's next

Now that you have mastered Kinetics for Edexcel IAL Chemistry Unit 4, you are ready to move on to the next core topics in the Rates, Equilibria and Further Organic Chemistry unit. Kinetics principles form the foundation for understanding how reaction conditions affect yield, which you will explore in depth in the chemical equilibria topic. You will also apply rate equation evidence to further organic reaction mechanisms in later organic chemistry units. Practice past paper questions on Kinetics to consolidate your skills, focusing on 6-mark rate deduction and Arrhenius calculation questions which are frequently tested and carry high mark weight. Make sure you can justify experimental technique choices for different reaction types, as these are common low-stakes mark questions that are easy to score if you learn the key rules.

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