Rates of Reaction

What you'll learn

This revision guide covers everything you need to know about rates of reaction for WJEC GCSE Chemistry. You'll understand how to measure reaction rates, apply collision theory to explain rate changes, and analyse the factors that speed up or slow down chemical reactions. These concepts form a core part of both the quantitative and qualitative chemistry assessed in your exams.

Key terms and definitions

Rate of reaction — the speed at which reactants are converted into products, measured as the change in amount of reactant or product per unit time.

Collision theory — the scientific explanation stating that particles must collide with sufficient energy and correct orientation for a reaction to occur.

Activation energy — the minimum energy that colliding particles must possess for a collision to result in a chemical reaction.

Catalyst — a substance that increases the rate of a reaction without being chemically changed or used up in the process.

Surface area to volume ratio — a measure comparing the exposed surface of a solid to its total volume; higher ratios mean faster reaction rates.

Concentration — the amount of solute dissolved in a given volume of solution, typically measured in g/dm³ or mol/dm³.

Successful collision — a collision between particles that results in a chemical reaction because the particles have sufficient energy (at or above activation energy) and correct orientation.

Reversible reaction — a chemical reaction where the products can react to reform the original reactants.

Core concepts

Measuring rates of reaction

There are three main methods for measuring reaction rates, each suitable for different reaction types:

Volume of gas produced

• Use a gas syringe or upturned measuring cylinder filled with water
• Record the volume of gas at regular time intervals (e.g., every 30 seconds)
• Plot volume against time to create a rate curve
• Suitable for reactions producing gases like CO₂, H₂, or O₂

Change in mass

• Place the reaction vessel on a digital balance
• Record mass at regular intervals as gas escapes
• Mass decreases as gaseous products leave the system
• More accurate for reactions producing heavier gases
• Cotton wool in the flask neck prevents liquid spray while allowing gas escape

Time for a precipitate to form

• Mix reactants that produce a cloudy precipitate
• Place a cross or mark beneath the reaction vessel
• Time how long it takes for the cross to disappear from view
• Suitable for reactions like sodium thiosulfate with hydrochloric acid
• Note: This method is subjective as different observers judge disappearance differently

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 measure (e.g., cm³/s for gas volume, g/s for mass change).

Collision theory and successful collisions

Collision theory explains why reaction conditions affect rates. For a reaction to occur:

1. Particles must collide with each other
2. They must possess sufficient energy (at least equal to the activation energy)
3. They must have the correct orientation relative to each other

Not all collisions lead to reactions. Most collisions are unsuccessful because particles either lack sufficient energy or approach from the wrong angle. Only successful collisions break bonds and form new products.

The activation energy acts as an energy barrier. Particles with kinetic energy below this threshold simply bounce off each other. Reactions with lower activation energies proceed faster because a greater proportion of collisions are successful.

Factors affecting reaction rates

Temperature

Increasing temperature increases reaction rate because:

• Particles gain kinetic energy and move faster
• Collision frequency increases (more collisions per second)
• More particles possess energy equal to or greater than the activation energy
• A higher proportion of collisions are successful

The energy increase is the more significant factor. Even a 10°C rise can double reaction rates in many cases.

Concentration (solutions) and pressure (gases)

Increasing concentration or pressure increases reaction rate because:

• More particles occupy the same volume
• Particles are closer together
• Collision frequency increases
• More successful collisions occur per unit time

Doubling the concentration of a reactant often doubles the initial rate of reaction. This relationship is particularly important for understanding reaction mechanisms.

Surface area

Increasing surface area of solid reactants increases reaction rate because:

• More particles are exposed to react
• Collision frequency with other reactants increases
• Breaking a large piece into smaller pieces dramatically increases surface area

For example, powdered calcium carbonate reacts much faster with hydrochloric acid than large marble chips of the same mass. This principle explains why dust in flour mills or coal mines poses explosion risks—fine particles have enormous surface areas.

Catalysts

Catalysts increase reaction rates by:

• Providing an alternative reaction pathway
• Lowering the activation energy
• Allowing more collisions to be successful at the same temperature
• Remaining chemically unchanged at the end of the reaction

Different catalysts work for different reactions:

• Iron in the Haber process (nitrogen + hydrogen → ammonia)
• Manganese(IV) oxide for decomposing hydrogen peroxide
• Nickel for hydrogenating vegetable oils
• Platinum/palladium/rhodium in catalytic converters

Biological catalysts (enzymes) follow the same principles but are specific to particular reactions and can be denatured by high temperatures or wrong pH values.

Interpreting rate of reaction graphs

Graphs of amount of product or amount of reactant against time have characteristic shapes:

Steep initial gradient — fast initial rate because:

• High concentration of reactants
• Maximum collision frequency
• Most reactant particles available

Gradient decreases — rate slows because:

• Reactants are being used up
• Concentration decreases
• Fewer collisions per second

Horizontal line (plateau) — reaction complete because:

• At least one reactant is completely used up (limiting reagent)
• No further product can form
• The amount of product remains constant

Comparing curves on the same axes:

• Steeper initial gradient = faster initial rate
• Curve reaching plateau earlier = faster overall reaction
• Same final height = same amount of product (if amount of limiting reactant is the same)
• Higher final plateau = more product formed (different amounts of limiting reactant)

Practical investigations

WJEC requires familiarity with practical methods for investigating rates. Common experiments include:

Marble chips and hydrochloric acid

• CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
• Measure CO₂ volume with gas syringe
• Investigate effect of concentration, temperature, or surface area (chip size)
• Control variables carefully (e.g., same mass of CaCO₃, same volume of HCl)

Magnesium ribbon and hydrochloric acid

• Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
• Measure H₂ volume or mass loss
• Fast reaction—good for temperature investigations

Sodium thiosulfate and hydrochloric acid

• Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + S(s) + H₂O(l)
• Sulfur precipitate makes solution cloudy
• Time disappearance of cross underneath flask
• Rate = 1 ÷ time
• Investigate concentration or temperature effects

Worked examples

Example 1: Calculating rate from graph data

Question: A student reacts zinc with sulfuric acid and measures the volume of hydrogen gas produced. After 20 seconds, 40 cm³ of gas has been produced. After 60 seconds, 80 cm³ of gas has been produced. Calculate the mean rate of reaction between 20 and 60 seconds. [3 marks]

Solution:

• Change in volume = 80 - 40 = 40 cm³ ✓
• Change in time = 60 - 20 = 40 s ✓
• Rate = 40 cm³ ÷ 40 s = 1 cm³/s ✓

Example 2: Explaining the effect of concentration

Question: Explain, in terms of particles, why increasing the concentration of hydrochloric acid increases the rate of reaction with magnesium ribbon. [3 marks]

Solution:

• Higher concentration means more acid particles in the same volume ✓
• This increases the frequency of collisions (between acid particles and magnesium atoms) ✓
• Therefore more successful collisions occur per second / rate increases ✓

Mark scheme note: Must refer to particles/collision frequency for full marks. Simply stating "more particles" without explaining collision frequency would lose a mark.

Example 3: Catalyst function

Question: A catalyst is added to a reaction mixture. Describe how a catalyst increases the rate of reaction and state one property of the catalyst at the end of the reaction. [3 marks]

Solution:

• The catalyst provides an alternative reaction pathway ✓
• This lowers the activation energy (of the reaction) ✓
• The catalyst is chemically unchanged / not used up ✓

Alternative acceptable answers: "more particles have sufficient energy to react" or "increases proportion of successful collisions"

Example 4: Graph interpretation

Question: A student investigated how surface area affects the rate of reaction between calcium carbonate and hydrochloric acid. Two experiments were performed using large chips and small chips. Both used the same mass of calcium carbonate and same volume of acid.

Sketch two curves on the same axes showing how the volume of carbon dioxide produced changes with time for: (a) large chips, (b) small chips. Label your curves L and S. [3 marks]

Solution: A correct sketch should show:

• Two curves starting at origin, curving upward then leveling off ✓
• Curve S (small chips) steeper initially than curve L ✓
• Both curves reaching the same final height / plateau ✓

Common mistakes and how to avoid them

• Confusing "faster reaction" with "more product" — Changing temperature or adding a catalyst makes reactions faster but doesn't change the final amount of product. The final amount depends on the quantities of reactants. Only compare final heights on graphs when the same amount of limiting reactant is used.

• Stating particles "have more energy to react" without mentioning collisions — Always link energy to collision theory. State that more particles have energy ≥ activation energy AND this increases successful collisions. Don't forget the collision aspect.

• Writing "particles move around more" instead of using precise language — Use scientific terminology: particles have greater kinetic energy, move faster, collision frequency increases. Avoid vague phrases like "move around more" or "bump into each other more."

• Forgetting that catalysts are not used up — A catalyst is unchanged at the end of the reaction. Don't say it "helps the reaction" without explaining HOW (lowering activation energy via alternative pathway). Don't confuse catalysts with reactants.

• Incorrect units or no units for rate — Always include units. For rate calculations, units depend on measurements: cm³/s, g/s, mol/s, or even s⁻¹ when calculating rate as 1/time. Match units to the data given.

• Not controlling variables in practical investigations — When investigating one factor (e.g., temperature), all other factors must remain constant (same concentration, same surface area, same volume, same mass). Identify control variables explicitly in method descriptions.

Exam technique for "Rates of Reaction"

• "Explain" questions require particle theory — When asked to explain rate changes, you must refer to particles, collisions, and energy. Structure answers: state what happens to particles → effect on collisions → effect on successful collisions/rate. Expect 2-3 marks for these explanations.

• Graph questions test multiple skills — Read axes carefully (is it mass loss or mass remaining? Volume of gas or volume of reactant?). Calculate gradients for rates at specific times. Compare curves systematically: initial gradient, overall time, final amount. Show working for gradient calculations: change in y ÷ change in x.

• Practical questions assess method and variables — Describe how you'd measure the dependent variable precisely (e.g., "use a gas syringe to measure volume of hydrogen every 30 seconds"). State controlled variables explicitly. Justify apparatus choices (e.g., "gas syringe more accurate than inverted measuring cylinder because...").

• Command words matter — "State" = simple answer, 1 mark. "Describe" = give characteristics or how to do something, 2 marks. "Explain" = give reasons using scientific knowledge, typically 2-3 marks. "Calculate" = show method and working, marks for process and answer.

Quick revision summary

Rate of reaction measures how quickly reactants form products. Collision theory explains that particles must collide with sufficient energy (≥ activation energy) and correct orientation. Temperature, concentration/pressure, surface area, and catalysts all increase rate by increasing collision frequency or successful collision proportion. Measure rates using gas volume, mass change, or precipitation time. Graph gradients show rate; steeper = faster. Catalysts provide alternative pathways with lower activation energy but remain unchanged. Always link explanations to particles and collisions for exam marks.