# Introduction to Kinetics and Equilibria

> Edexcel International A-Level Chemistry · IAL Chemistry U2
> Source: https://www.owlsprep.com/study/edexcel-ial-chemistry-u2-introduction-to-kinetics-and-equilibria/

This guide covers all qualitative content for Edexcel IAL Chemistry Unit 2 Topic 9, including collision theory, Maxwell-Boltzmann distributions, dynamic equilibrium, qualitative Le Chatelier’s principle, and industrial reaction compromises.

**Prerequisites:** [Basic atomic and molecular collision concepts](https://www.owlsprep.com/study/edexcel-ial-chemistry-u1-structure-of-the-atom/); [Understanding of exothermic and endothermic reactions](https://www.owlsprep.com/study/edexcel-ial-chemistry-u2-energetics/)

## Learning objectives

- Apply collision theory to explain effects of concentration, temperature, pressure and surface area on reaction rate
- Interpret Maxwell-Boltzmann distribution curves to explain effects of temperature and catalysts on reaction rate
- Calculate reaction rate using 1/time and tangent gradient methods
- Define dynamic equilibrium and apply Le Chatelier's principle to predict equilibrium position shifts
- Justify industrial compromise conditions balancing reaction yield and rate

## Collision Theory and Reaction Rate Calculation

**Collision Theory** — A model explaining reaction rate: reactions occur when reactant particles collide with sufficient energy (≥ activation energy, $E_a$) and correct orientation to break bonds and form products.

Four key factors affect the frequency of successful collisions and thus reaction rate: concentration/pressure (higher values = more particles per unit volume = more collisions), temperature (higher = particles move faster, more high-energy collisions), surface area (larger = more exposed reactant particles available for collision), and presence of a catalyst.

**Worked example:** A student investigates the rate of reaction between magnesium ribbon and dilute hydrochloric acid. They time how long it takes for 10 cm³ of hydrogen gas to be produced, recording a time of 25 s. Calculate the relative rate of reaction, and state the effect on rate if the acid is replaced with double the concentration.

1. 1. Use the rate = 1/time formula for relative rate calculation:
2. $$r = \frac{1}{t} = \frac{1}{25} = 0.04 \, s^{-1}$$
3. 2. Doubling the concentration of hydrochloric acid increases the number of H+ ions per unit volume. This increases the frequency of successful collisions between Mg and H+ particles.
4. 3. The reaction rate will approximately double.

**Check your understanding**

1. Why does increasing the surface area of a solid reactant increase reaction rate?

   - It increases the activation energy of the reaction
   - It increases the number of exposed reactant particles available for collision
   - It increases the average kinetic energy of reactant particles
   - It changes the orientation of collisions

   *Answer:* It increases the number of exposed reactant particles available for collision

   *Why:* Incorrect: Surface area does not affect activation energy or particle kinetic energy. It only exposes more particles to collisions.

## Maxwell-Boltzmann Distribution and Temperature Effects

**Maxwell-Boltzmann Distribution** — A graph showing the distribution of kinetic energies of gas (or solution) particles at a fixed temperature. The area under the curve equals the total number of particles, which remains constant for a closed system.

Key features of the curve: the peak represents the most probable kinetic energy, the mean energy is to the right of the peak, and only the fraction of particles with energy ≥ $E_a$ can react successfully. When temperature increases: the curve shifts right, the peak becomes lower and broader, the area under the curve stays the same, and the fraction of particles with energy ≥ $E_a$ increases significantly, leading to a faster reaction rate.

**Worked example:** Sketch and label a comparison of Maxwell-Boltzmann curves for a reaction at 298 K and 310 K, marking the activation energy $E_a$ and explaining the effect of the temperature increase on reaction rate.

1. 1. Draw the 298 K curve: higher, narrower peak at lower kinetic energy.
2. 2. Draw the 310 K curve: lower, broader peak shifted to higher kinetic energy, both curves start at the origin and asymptote to the x-axis at high energy.
3. 3. Mark $E_a$ as a vertical line to the right of both peaks.
4. 4. The area under the curve to the right of $E_a$ is much larger for 310 K than 298 K: a greater fraction of particles have sufficient energy for successful collisions, so rate increases.

> **Exam tip:** Always state that the area under the M-B curve is constant (equal to total number of particles) in your exam answers, this is a common mark point.

## Catalysts and Reaction Profiles

**Catalyst** — A substance that increases the rate of a chemical reaction without being used up in the reaction. It works by providing an alternative reaction route with a lower activation energy $E_a$.

Reaction profiles (energy level diagrams) show the change in energy of reactants as they form products. For an uncatalysed reaction, the peak energy equals the activation energy for the original route. For a catalysed reaction, the peak is lower, and may show an intermediate energy level if the catalyst forms a temporary intermediate with reactants. Catalysts do not affect the enthalpy change of the reaction, or the position of equilibrium: they only increase the rate at which equilibrium is reached.

**Worked example:** Draw a reaction profile for an exothermic reaction, showing both catalysed and uncatalysed routes, and label the activation energy for both.

1. 1. Draw a horizontal line for reactant energy, a lower horizontal line for product energy (since reaction is exothermic, ΔH is negative).
2. 2. Draw a high curved peak from reactants to products for the uncatalysed route: label the vertical difference between reactant energy and the peak as $E_a$(uncatalysed).
3. 3. Draw a lower curved peak (or two smaller peaks showing an intermediate) for the catalysed route: label the vertical difference between reactant energy and the lower peak as $E_a$(catalysed).
4. 4. Label the enthalpy change ΔH as the vertical difference between reactant and product energy.

> **info**
>
> Industrial catalysts reduce the temperature and pressure required for reactions, cutting energy use, greenhouse gas emissions and operational costs, improving process sustainability.

## Dynamic Equilibrium and Le Chatelier's Principle

**Dynamic Equilibrium** — A state in a closed reversible reaction system where the rate of the forward reaction equals the rate of the reverse reaction, so the concentrations of reactants and products remain constant (they are not equal, just unchanging).

Le Chatelier's Principle states that if a change is applied to a system at dynamic equilibrium, the position of equilibrium will shift to oppose the change. You only need to apply this qualitatively for homogeneous systems (all reactants and products in the same state) for this topic: 1. Concentration: increase reactant concentration → shift right to use up extra reactants; increase product concentration → shift left to use up extra products. 2. Pressure (for gaseous systems only): increase pressure → shift to the side with fewer moles of gas to reduce pressure; decrease pressure → shift to the side with more moles of gas. 3. Temperature: increase temperature → shift in the endothermic direction to absorb extra heat; decrease temperature → shift in the exothermic direction to release heat.

**Worked example:** Consider the homogeneous gaseous equilibrium: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \quad \Delta H = -92 \, kJ \, mol^{-1}$. Predict and justify the effect of increasing pressure on the position of equilibrium.

1. 1. Count the moles of gas on each side: left side = 1 + 3 = 4 moles of gas, right side = 2 moles of gas.
2. 2. Increasing pressure shifts equilibrium to the side with fewer moles of gas to oppose the pressure increase.
3. 3. Equilibrium shifts to the right, increasing the yield of ammonia.

**Exam command terms**

Key command terms for equilibrium questions:

- **Predict** — State the direction of the equilibrium shift (left/right/no change) *(Predict the effect of increasing temperature on the equilibrium yield of product.)*

- **Justify** — Explain the shift using Le Chatelier's principle, referencing the change applied and how the shift opposes it. *(Justify your prediction of the equilibrium shift when pressure is increased.)*

## Industrial Compromise Conditions

Industrial chemical processes balance three key factors: maximum yield, fast reaction rate, and low operational cost. The optimal conditions are often a compromise between these factors, as conditions that increase rate may reduce yield, or vice versa. For example, the Haber process (ammonia production) uses moderate temperature, high pressure, and an iron catalyst to balance yield, rate and cost.

**Worked example:** For the Haber process equilibrium (exothermic forward reaction), explain why a moderate temperature of around 450°C is used instead of a much lower temperature, even though lower temperature gives a higher equilibrium yield of ammonia.

1. 1. A lower temperature would shift equilibrium right to give higher ammonia yield, but reaction rate would be very slow, as fewer particles have energy ≥ $E_a$.
2. 2. A very high temperature would give a fast rate, but shift equilibrium left, reducing ammonia yield significantly.
3. 3. The moderate 450°C temperature is a compromise: it gives a fast enough reaction rate while still producing an acceptable yield of ammonia, especially when combined with an iron catalyst to further increase rate.

## Common pitfalls

- **Wrong:** Stating that a catalyst shifts the position of equilibrium or increases product yield.
  - Why it fails: Catalysts only increase the rate of forward and reverse reactions equally, so they do not affect equilibrium position or yield, only the time taken to reach equilibrium.
  - Correct: State that catalysts reduce activation energy, increasing reaction rate, but have no effect on equilibrium yield or position.
- **Wrong:** Claiming that increasing temperature increases the total number of particles with energy ≥ $E_a$ by increasing the total number of particles.
  - Why it fails: The total number of particles in a closed system is constant. Higher temperature increases the *fraction* of particles with energy ≥ $E_a$, not the total number of particles.
  - Correct: Explain that higher temperature shifts the Maxwell-Boltzmann distribution right, increasing the fraction of particles with energy ≥ $E_a$, leading to more successful collisions.
- **Wrong:** Stating that at dynamic equilibrium, the concentrations of reactants and products are equal.
  - Why it fails: At dynamic equilibrium, concentrations are *constant* (unchanging), not equal. The ratio of concentrations depends on the reaction and conditions.
  - Correct: Define dynamic equilibrium as equal forward and reverse reaction rates, leading to constant reactant and product concentrations.
- **Wrong:** Applying pressure change effects to equilibrium systems containing only solids or liquids.
  - Why it fails: Pressure changes only affect the volume (and thus concentration) of gaseous particles; solids and liquids have negligible volume changes with pressure.
  - Correct: Only apply pressure shift rules to homogeneous gaseous equilibrium systems, counting moles of gas only on each side of the equation.
- **Wrong:** Calculating reaction rate from time as rate = time, instead of rate = 1/time for relative rate.
  - Why it fails: Faster reactions take less time, so rate is inversely proportional to time taken for a fixed amount of product to form.
  - Correct: Calculate relative rate as the inverse of time taken, with units of $s^{-1}$ if time is measured in seconds.

## Cheatsheet

| Concept | Key Exam Fact |
| --- | --- |
| Collision Theory | Successful collisions require ≥ Eₐ and correct orientation |
| Rate calculation | Relative rate = 1/time; rate = gradient of tangent to conc-time graph |
| Maxwell-Boltzmann (higher T) | Curve shifts right, lower peak, higher fraction of particles ≥ Eₐ |
| Catalyst effect | Lowers Eₐ, no change to M-B curve, no effect on equilibrium yield |
| Dynamic Equilibrium | Forward rate = reverse rate, constant (not equal) reactant/product concentrations |
| Le Chatelier: Pressure increase | Shift to side with fewer moles of gas (gaseous systems only) |
| Le Chatelier: Temperature increase | Shift in endothermic direction to absorb heat |
| Industrial compromise | Balance yield, rate, and operational cost for optimal conditions |

## What's next

Now that you have mastered the qualitative fundamentals of kinetics and equilibria for Edexcel IAL Chemistry Unit 2, you are ready to move on to more advanced chemistry topics in the unit, including group chemistry and halogenoalkane reactions. This content also forms the foundation for quantitative kinetics and equilibrium calculations you will encounter in Unit 4 of the IAL Chemistry course, where you will learn to use rate equations, rate constants, and equilibrium constants to solve numerical problems. Ensure you practice past paper questions on this qualitative content to familiarize yourself with common exam phrasing and mark scheme requirements.

---

From [OwlsPrep](https://www.owlsprep.com) — free study guides for A-Level, IB, AP and IGCSE, written against the official syllabus. Canonical page: https://www.owlsprep.com/study/edexcel-ial-chemistry-u2-introduction-to-kinetics-and-equilibria/