Rates of Reaction and Energy — Edexcel GCSE Chemistry Revision Notes

What you'll learn

This topic combines two fundamental areas of Edexcel GCSE Chemistry: how quickly chemical reactions occur and the energy transfers involved. You'll encounter questions on rates of reaction in Paper 1 (Topic 6) and energy changes in Paper 2 (Topic 7), worth approximately 10-15% of your total marks. Understanding collision theory, activation energy, and how to interpret reaction profiles is essential for securing marks in both multiple-choice and extended-response questions.

Key terms and definitions

Rate of reaction — the speed at which reactants are converted into products, measured as the change in concentration of a reactant or product per unit time (typically g/s, cm³/s, or mol/s).

Activation energy — the minimum energy that colliding particles must possess for a chemical reaction to occur, represented by Ea on energy profile diagrams.

Catalyst — a substance that increases the rate of a chemical reaction without being chemically changed or used up in the reaction, by providing an alternative reaction pathway with lower activation energy.

Collision theory — the principle that chemical reactions occur only when particles collide with sufficient energy (equal to or greater than the activation energy) and with the correct orientation.

Exothermic reaction — a chemical reaction that transfers energy to the surroundings, usually as heat, resulting in a temperature increase of the surroundings. The energy of products is lower than the energy of reactants.

Endothermic reaction — a chemical reaction that takes in energy from the surroundings, usually as heat, resulting in a temperature decrease of the surroundings. The energy of products is higher than the energy of reactants.

Successful collision — a collision between reactant particles that has both sufficient energy (≥ activation energy) and correct geometric orientation to break bonds and form products.

Frequency of collisions — the number of collisions between reactant particles per unit time, which increases with higher concentration, pressure, temperature, or surface area.

Core concepts

Measuring rates of reaction

Edexcel GCSE Chemistry exams frequently test your ability to describe and analyse methods for measuring reaction rates. Three common experimental approaches appear regularly:

Measuring gas volume produced:

Use a gas syringe or inverted measuring cylinder filled with water
Record the volume of gas at regular time intervals (e.g., every 30 seconds)
Plot volume against time to produce a rate curve
Example reaction: marble chips (calcium carbonate) reacting with hydrochloric acid produces carbon dioxide
Equation: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Measuring mass loss:

Place the reaction vessel on a balance
Record the mass at regular intervals as gas escapes
The mass decreases as gas is released to the atmosphere
More accurate than gas collection for fast reactions
Example: magnesium ribbon reacting with sulfuric acid

Measuring colour/turbidity change:

Time how long it takes for a cross below a flask to disappear
Used when a precipitate forms that makes the solution cloudy
Example: sodium thiosulfate reacting with hydrochloric acid produces sulfur precipitate
Equation: Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g) + S(s)

The rate of reaction can be calculated using:

Rate = amount of reactant used / time OR Rate = amount of product formed / time

Units depend on what you're measuring: g/s, cm³/s, or mol/s are most common in Edexcel papers.

Factors affecting rates of reaction

Edexcel examiners expect you to explain rate changes using collision theory. Each factor must be linked to collision frequency and/or energy:

Temperature:

Increasing temperature increases the kinetic energy of particles
Particles move faster, colliding more frequently
More importantly, a greater proportion of collisions now exceed the activation energy
This dual effect makes temperature a particularly powerful factor
A 10°C increase typically doubles the reaction rate for many reactions

Concentration (solutions) and pressure (gases):

Higher concentration means more particles in the same volume
Particles are closer together, increasing collision frequency
More collisions per second leads to more successful collisions per second
Pressure has the same effect for gases by reducing the volume available
Example: 2.0 mol/dm³ hydrochloric acid reacts faster with zinc than 1.0 mol/dm³ acid

Surface area:

Breaking a solid into smaller pieces increases total surface area
More particles are exposed and available for collision
Powdered calcium carbonate reacts faster than large marble chips with the same mass
Industrial importance: catalytic converters use honeycomb structures to maximize surface area

Catalysts:

Provide an alternative reaction pathway with lower activation energy
More colliding particles now have sufficient energy to react
The catalyst participates in the reaction mechanism but is regenerated
Remains chemically unchanged at the end
Industrial examples: iron in the Haber process (N₂ + 3H₂ → 2NH₃), vanadium(V) oxide in the Contact process for sulfuric acid manufacture

Interpreting rate of reaction graphs

Edexcel papers regularly include graphs showing volume of gas or mass against time. You must be able to:

Identify key features:

Steepest gradient = fastest rate (usually at the start when concentration is highest)
Gradient decreases as reactants are used up
Horizontal line = reaction finished (one or more reactants completely used up)

Compare different conditions:

Same final volume/mass = same amount of reactants used
Steeper initial gradient = faster rate due to changed conditions (higher temperature, concentration, surface area, or catalyst present)
Horizontal line reached sooner = reaction completed more quickly
Different final volume/mass = different amounts of limiting reactant

Calculate rate at a specific time:

Draw a tangent to the curve at that time point
Calculate gradient = change in y / change in x
Include units in your answer

Energy changes in reactions

Every chemical reaction involves energy changes as bonds are broken and formed. Edexcel GCSE Chemistry requires you to understand energy profiles and calculate energy changes.

Bond breaking and bond making:

Breaking bonds is endothermic — requires energy input
Making bonds is exothermic — releases energy
Overall energy change = energy required to break bonds - energy released when bonds form

Exothermic reactions:

Energy released from forming new bonds exceeds energy required to break original bonds
Products have lower energy than reactants
ΔH is negative (by convention)
Temperature of surroundings increases
Examples: combustion, neutralization, oxidation, most displacement reactions
Real example: methane combustion CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = -890 kJ/mol

Endothermic reactions:

Energy required to break bonds exceeds energy released from forming new bonds
Products have higher energy than reactants
ΔH is positive
Temperature of surroundings decreases
Examples: thermal decomposition, photosynthesis, dissolving some salts (e.g., ammonium nitrate)
Real example: calcium carbonate decomposition CaCO₃(s) → CaO(s) + CO₂(g) ΔH = +178 kJ/mol

Reaction profiles

Reaction profile diagrams appear frequently in Edexcel papers. You must label and interpret them correctly.

Key features to identify and label:

Y-axis: Energy (sometimes labeled "Enthalpy")
X-axis: Progress of reaction or Reaction pathway
Reactants energy level (starting point on left)
Products energy level (end point on right)
Activation energy (Ea): height from reactants to the peak
Overall energy change (ΔH): difference between products and reactants levels

Exothermic profile:

Products lower than reactants
Energy released to surroundings
Downward arrow showing ΔH as negative

Endothermic profile:

Products higher than reactants
Energy absorbed from surroundings
Upward arrow showing ΔH as positive

Effect of catalysts on profiles:

Lower activation energy peak
Same starting reactants level
Same ending products level
Overall ΔH unchanged
Often shown as a dotted line for comparison

Calculating energy changes

Edexcel GCSE Chemistry includes calculations using bond energies. The method involves:

Step 1: Identify all bonds broken in reactants (energy IN)

Step 2: Identify all bonds formed in products (energy OUT)

Step 3: Calculate total energy in and total energy out using bond energy values from the data sheet

Step 4: Energy change = total energy in - total energy out

If the answer is negative, the reaction is exothermic. If positive, the reaction is endothermic.

You must show all working to access full marks. Edexcel mark schemes award method marks even if the final answer is incorrect.

Worked examples

Example 1: Rates of reaction practical

Question: A student investigates the reaction between magnesium ribbon and dilute hydrochloric acid. The student measures the volume of hydrogen gas produced every 30 seconds. Describe a method the student could use to collect and measure the gas accurately. (4 marks)

Answer:

Add a known volume of hydrochloric acid to a conical flask (1)
Add a known mass/length of magnesium ribbon and immediately attach a bung connected to a gas syringe/delivery tube and inverted measuring cylinder (1)
Record the volume of gas at regular time intervals, e.g., every 30 seconds, until no more gas is produced (1)
Plot a graph of volume against time (1)

Mark scheme notes: Accept alternatives such as measuring mass loss using a balance. Award marks for stating specific quantities and regular time intervals.

Example 2: Energy profile interpretation

Question: The diagram shows an energy profile for a reaction.

[Imagine a profile with products lower than reactants, activation energy labeled as 150 kJ/mol, and energy change of -80 kJ/mol]

(a) Is this reaction exothermic or endothermic? Explain your answer. (2 marks)

(b) What is the activation energy for this reaction? (1 mark)

Answers:

(a) Exothermic (1) because the products have lower energy than the reactants / energy is released to the surroundings (1)

(b) 150 kJ/mol (1)

Example 3: Bond energy calculation

Question: Hydrogen reacts with chlorine to form hydrogen chloride:

H₂(g) + Cl₂(g) → 2HCl(g)

Use the bond energies below to calculate the energy change for this reaction.

H-H: 436 kJ/mol, Cl-Cl: 242 kJ/mol, H-Cl: 431 kJ/mol (4 marks)

Answer:

Bonds broken: 1 × H-H = 436 kJ, 1 × Cl-Cl = 242 kJ

Total energy in = 436 + 242 = 678 kJ (1)

Bonds formed: 2 × H-Cl = 2 × 431 = 862 kJ (1)

Total energy out = 862 kJ (1)

Energy change = 678 - 862 = -184 kJ/mol (1)

The reaction is exothermic (accept without this statement for full marks)

Common mistakes and how to avoid them

Mistake: Stating that catalysts "speed up particles" or "give particles more energy."

Correction: Catalysts lower the activation energy by providing an alternative pathway. They do not change the energy or speed of particles. More particles now have sufficient energy to react when they collide.

Mistake: Confusing the overall energy change (ΔH) with activation energy (Ea) on reaction profiles.

Correction: Activation energy is always measured from the reactants level to the peak. The energy change is the difference between products and reactants levels. A reaction can be exothermic (negative ΔH) but still have a high activation energy.

Mistake: Saying concentration affects how hard particles collide.

Correction: Concentration affects collision frequency only (how often particles collide), not the energy of collisions. Temperature is the factor that affects collision energy.

Mistake: In bond energy calculations, subtracting the wrong way: energy out - energy in.

Correction: Always calculate: energy change = energy IN (bonds broken) - energy OUT (bonds formed). Breaking bonds requires energy, making bonds releases energy.

Mistake: Drawing catalyst lines on reaction profiles that change the reactants or products energy levels.

Correction: A catalyst only lowers the activation energy peak. The reactants and products energy levels remain identical. Draw a lower curve that starts and ends at exactly the same points.

Mistake: Claiming a steeper graph means "more gas produced" when comparing reactions with the same reactants.

Correction: A steeper gradient means a faster rate, not more product. If the same mass of reactants is used, the final volume/mass will be identical regardless of rate. The steeper curve simply reaches this final value sooner.

Exam technique for Rates of Reaction and Energy

Command word awareness: "Explain" questions about rates require you to link your answer to collision theory. State the factor change AND its effect on collision frequency or energy. For example: "Increasing temperature increases the kinetic energy of particles, so more collisions exceed the activation energy, leading to more successful collisions per second." This structure typically earns 2-3 marks.

Graph questions: When comparing curves, systematically describe (1) initial gradient/rate, (2) final volume/mass, and (3) time taken to complete. Edexcel mark schemes award one mark per valid comparison. Always refer to specific data from the graph.

Bond energy calculations: Show every step of your working. Examiners award method marks for correctly identifying bonds broken and bonds formed, even if arithmetic errors occur. Always include units (kJ/mol) and state whether the reaction is exothermic or endothermic if the question asks.

Practical method questions: Structure answers chronologically and include specific details: volumes, masses, time intervals, apparatus names. Questions worth 4-6 marks expect 4-6 distinct points. Safety precautions or control variables may also be required.

Quick revision summary

Rate of reaction measures how quickly reactants form products. Collision theory explains that reactions occur only when particles collide with sufficient energy (activation energy) and correct orientation. Rates increase with higher temperature, concentration, pressure, surface area, or by adding a catalyst. Exothermic reactions release energy (products lower energy than reactants, ΔH negative); endothermic reactions absorb energy (products higher energy, ΔH positive). Energy changes equal energy required to break bonds minus energy released forming bonds. Catalysts lower activation energy without being used up, providing an alternative reaction pathway.