What is the rate of reaction in Chemistry?

The rate of reaction in Chemistry refers to the speed at which reactants are converted into products in a chemical reaction. It is typically measured by the change in concentration of a reactant or product over time. Understanding the rate of reaction is crucial for controlling industrial processes and is a key concept in the Cambridge IGCSE Chemistry curriculum.

How can temperature affect the rate of reaction?

Temperature can significantly affect the rate of reaction. Generally, increasing the temperature increases the rate of reaction. This is because higher temperatures provide reactant particles with more energy, leading to more frequent and energetic collisions, which are necessary for reactions to occur. This concept is important for students studying the Cambridge IGCSE Chemistry syllabus.

What role does concentration play in the rate of reaction?

Concentration plays a crucial role in the rate of reaction. A higher concentration of reactants typically leads to a faster reaction rate because there are more particles available to collide and react. This principle is fundamental in understanding chemical kinetics, a topic covered in the Cambridge IGCSE Chemistry curriculum.

Why do catalysts increase the rate of reaction?

Catalysts increase the rate of reaction by providing an alternative reaction pathway with a lower activation energy. This means that more reactant particles have enough energy to react, leading to an increased rate of reaction. Catalysts are not consumed in the reaction, making them efficient for repeated use. This concept is essential for students learning about chemical reactions in Cambridge IGCSE Chemistry.

How does surface area affect the rate of reaction?

Surface area affects the rate of reaction by influencing the number of particles available for collisions. A larger surface area allows more particles to be exposed and available to react, increasing the rate of reaction. This is particularly relevant in reactions involving solids, where breaking a solid into smaller pieces increases its surface area. Understanding this concept is important for students studying the rate of reaction in the Cambridge IGCSE Chemistry course.

Why is "Saying 'particles collide more' as the only explanation" a common mistake in Rate of Reaction?

Why it happens: Imprecise. How to avoid it: Use 'frequency of SUCCESSFUL collisions' AND distinguish FREQUENCY (concentration, surface area) from ENERGY (temperature, catalyst). Source: 0620/42 — every series

Why is "Saying catalysts are consumed" a common mistake in Rate of Reaction?

Why it happens: Confusing with reactants. How to avoid it: Catalysts are recovered chemically unchanged at the end.

Why is "Saying temperature only increases collision frequency" a common mistake in Rate of Reaction?

Why it happens: Forgetting the bigger effect. How to avoid it: Bigger effect is the increased FRACTION of particles with E E_a.

Why is "Saying a catalyst changes the position of equilibrium" a common mistake in Rate of Reaction?

Why it happens: Generalising 'speeds up'. How to avoid it: A catalyst speeds up forward AND reverse equally → position of equilibrium UNCHANGED.

Chemical ReactionsCambridge IGCSE Chemistry (0620)

Rate of Reaction

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Rate of Reaction — Cambridge IGCSE 0620 Chemistry Extended (2026)

Five factors that change reaction speed (concentration, temperature, surface area, pressure, catalyst). Plus the collision theory that explains them all.

What you’ll learn

Mapped to the Cambridge IGCSE 0620 syllabus (2026-2028).

7.2 — State the factors that affect rate of reaction.
7.2 — Explain rate effects using collision theory.
7.2 — Describe the use of catalysts and give industrial examples.
7.2 — Sketch and interpret rate-time graphs.

Collision theory — the underlying explanation

Reactions happen when particles collide with enough energy AND right orientation. More effective collisions → faster rate.

Collision theory. For two particles to react, three things must happen:

They must COLLIDE.
They must collide with enough energy (at least the activation energy, $E_a$ ).
They must be ORIENTED correctly (for many reactions).

Effective collision = a collision that meets all three conditions and leads to a reaction.

Increasing rate means increasing the rate of effective collisions. Each of the five factors below works through this lens.

Worked qualitative. Why doesn't every collision produce a reaction? Most collisions are too gentle (below $E_a$ ) or have the wrong orientation. Only a small fraction of collisions are effective.

Rate = effective collisions per second.
Need: collision + enough energy + right orientation.
Most collisions don't react.

Concentration and pressure

More particles in the same volume → more collisions per second → faster.

Concentration (solutions). Higher concentration = more particles per dm³.

More particles in a given volume.
More likely they'll collide.
More effective collisions per second → FASTER rate.

Worked qualitative. Doubling the concentration of HCl in a reaction with marble chips roughly DOUBLES the rate (for first-order reactions). The reaction lasts about half the time.

Pressure (gases). Higher pressure = compressing the gas = more particles per volume.

Same effect as concentration: more collisions per second → faster.

Worked qualitative. Increasing pressure on a gas-phase reaction (e.g. Haber process N₂ + H₂ → NH₃) increases the rate of forward reaction.

Pressure on liquid/solid reactions. Negligible effect — particles already close together.

More concentrated → more particles per volume.
More pressure → same idea for gases.
More collisions per second → faster rate.
Pressure barely affects condensed-phase reactions.

Temperature

Hotter particles move faster AND have more energy → many more effective collisions.

Two effects of higher temperature.

Particles move FASTER → collide MORE OFTEN.
Particles have MORE energy on average → MORE collisions exceed the activation energy.

The second effect is much bigger. A 10°C rise typically DOUBLES the rate.

Worked qualitative. Why does food spoil more slowly in a fridge? Lower temp → enzymes (catalysts) and microbial reactions are slowed down → food lasts longer.

Why heating works two ways:

Counts collisions per second (slight increase).
Counts the FRACTION of collisions with enough energy (big increase).

Cambridge tip. When asked "why does temperature increase rate?", give BOTH reasons: more frequent collisions AND a higher fraction of collisions with sufficient energy.

Higher T → faster particles → more collisions.
Higher T → more energy per particle → bigger fraction above $E_a$ .
10°C rise: typically doubles the rate.

Surface area (for solids)

Smaller pieces or powdered = more surface exposed → more collision sites → faster.

Surface area effect. When a solid reacts with a liquid or gas, the reaction can only happen at the SURFACE.

A large lump has small surface-area-to-volume ratio.
Powder has huge SA:V — much more surface exposed.
Powder reacts MUCH faster than a lump of the same mass.

Worked qualitative. Why does a flour mill have explosion safety rules? Powdered flour has enormous surface area. Mixed with air, the slightest spark can trigger rapid combustion of the entire dust cloud — flour-dust explosions are real industrial hazards.

Practical experiment. Marble chips in HCl: vary chip size (small → medium → large). Plot CO₂ volume vs time. Smallest chips give the steepest initial slope.

Cambridge tip. Always describe surface area in terms of SOLIDS reacting with gases or liquids. Two liquids or two gases don't have "surface area" in this sense (they mix freely).

Smaller solid pieces = bigger SA:V ratio.
More surface = more collision sites.
Faster rate.
Industrial hazard: dust explosions.

Catalysts

Lower the activation energy → more collisions are effective → faster, without being used up.

Catalyst. A substance that increases the rate of a reaction without being chemically used up itself. It can be reused indefinitely.

Mechanism. Catalysts provide an ALTERNATIVE PATHWAY with a LOWER activation energy.

More particles have enough energy to overcome $E_a$ .
More effective collisions.
Faster rate.

Catalyst does NOT change $\Delta H$ . Same products, same energy difference between reactants and products. Just a different (faster) route.

Industrial examples.

Iron in the Haber process ( $\text{N}_{2} + 3\text{H}_{2} \to 2\text{NH}_{3}$ ).
Vanadium(V) oxide in the Contact process (manufacture of sulfuric acid).
Nickel in the hydrogenation of vegetable oils to make margarine.
Platinum, rhodium, palladium in catalytic converters (cars).
Enzymes in biology — biological catalysts that speed up reactions in living cells.

Worked qualitative. A factory using a catalyst can run at LOWER temperatures while keeping the same rate. Saves energy → cheaper to operate → less environmental impact.

Catalyst lowers $E_a$ .
More collisions become effective.
Catalyst not used up.
$\Delta H$ unchanged.
Industrial: Haber, Contact, hydrogenation, catalytic converters.

Rate-time graphs

Plot product (or reactant) over time. Steepest at start; flattens as reactants run out.

Typical graph. Plot the volume of gas produced (or mass lost, etc.) on the y-axis vs time on the x-axis.

Curve rises steeply at first (highest rate when reactants most concentrated).
Slope decreases as reactants are used up.
Flattens out when reaction completes (reactants exhausted).

Reading the rate. Rate at any point = gradient of the curve. Use a tangent at that point.

Comparing conditions. Two reactions starting from the same amount but at different temperatures (say) give two curves. The hotter one rises faster — steeper initial slope. Both reach the same FINAL volume (if the reaction goes to completion) — same yield, just different rate.

Worked. Volume of CO₂ from $1\,\text{g}$ marble + 50 cm³ HCl after 1 minute: 50 cm³. After 2 minutes: 80 cm³. After 3 minutes: 95 cm³. After 5 minutes: 100 cm³ (constant).

Initial rate $\approx 50\,\text{cm}^{3}/\text{min}$ .
Average rate over 5 min: $100/5 = 20\,\text{cm}^{3}/\text{min}$ .
Reaction effectively complete at 5 min.

Steepest at start; flattens as reactants run out.
Rate at any moment = gradient of tangent.
Higher T / catalyst: same final yield, faster reach.
Same starting amount: same final volume regardless of conditions.

How it’s examined

Rate of reaction is examined every Paper 4 (8-12 marks): factors, collision theory explanations, rate-time graph. Examiner reports flag students explaining temperature with only 'particles move faster' (Cambridge wants BOTH 'more frequent' AND 'higher fraction with $E_a$ ').

Sources: Cambridge IGCSE Chemistry 0620 syllabus 2026-2028 (7.2); 0620/42 Oct/Nov 2024 — Q11 (rate, catalyst); 0620 Examiner Reports 2022-2024. Last reviewed 2026-05-06.

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Step-by-step worked examples — Rate of Reaction

Step-by-step solutions to past-paper-style questions on rate of reaction, written exactly the way a tutor would explain them at the board.

1Rate from a gas-volume graph
Extended• Adapted from 0620/42 May/Jun 2024 Q15• graph
Question
$60\,\text{cm}^{3}$ of gas is produced in $30\,\text{s}$ . Find the average rate.
Step-by-step solution
1. Step 1
  Rate = quantity / time.
  $\text{rate} = \dfrac{60}{30} = 2.0\,\text{cm}^{3}/\text{s}$
Answer
$2.0\,\text{cm}^{3}/\text{s}$
2Why higher concentration speeds up the reaction
Extended• concentration
Question
Explain in particle terms why doubling the concentration of one reactant doubles the rate.
Step-by-step solution
1. Step 1
  Higher concentration → more reactant particles per unit volume.
2. Step 2
  More frequent collisions per unit time → more SUCCESSFUL collisions per second.
Answer
More particles per volume → more frequent successful collisions → higher rate.
3Effect of temperature
Extended• temperature
Question
Explain why a $10\degree\text{C}$ rise typically doubles the rate.
Step-by-step solution
1. Step 1
  Higher temperature → particles move FASTER → more frequent collisions.
2. Step 2
  MORE IMPORTANTLY: a much greater fraction have energy ≥ activation energy → many more collisions are SUCCESSFUL.
Answer
Faster motion + greater fraction with $E \ge E_{a}$ → far more successful collisions.
Examiner tip
The bigger effect is on the FRACTION above $E_{a}$ , not just collision frequency.
4Effect of a catalyst
Extended• catalyst
Question
Explain how a catalyst speeds up a reaction.
Step-by-step solution
1. Step 1
  Provides an alternative reaction path with LOWER activation energy.
2. Step 2
  More particles have $E \ge E_{a}$ → more successful collisions per second.
3. Step 3
  Catalyst is not consumed; can be reused.
Answer
Lower $E_{a}$ → more particles have enough energy → faster reaction.

Key Formulae — Rate of Reaction

The formulae you need to memorise for rate of reaction on the Cambridge IGCSE 0620 paper, with every variable defined in plain English and a note on when to use it.

Rate of reaction
$\text{rate} = \dfrac{\text{change in amount of reactant or product}}{\text{time}}$
When to use
Computing average or instantaneous rate.

Key Definitions and Keywords — Rate of Reaction

Definitions to memorise and the exact keywords mark schemes credit for rate of reaction answers — sharpened from recent examiner reports for the 2026 0620 sitting.

Collision theory
Examiner keyword
Reactions occur when reactant particles collide with sufficient energy ( $\ge E_{a}$ ) and the correct orientation.
Activation energy ( $E_{a}$ )
Examiner keyword
Minimum kinetic energy that colliding particles need for a successful reaction.
Catalyst
Examiner keyword
Substance that increases the rate by providing a route with LOWER activation energy. NOT used up.
Surface area
Examiner keyword
Larger surface area (smaller particles) → more sites for collisions → higher rate.

Common Mistakes and Misconceptions — Rate of Reaction

The traps other students keep falling into on rate of reaction questions — taken from recent Cambridge IGCSE 0620 examiner reports and mark schemes — and how to avoid them.

✕Saying 'particles collide more' as the only explanation
0620/42 — every series
Why it happens
Imprecise.
How to avoid it
Use 'frequency of SUCCESSFUL collisions' AND distinguish FREQUENCY (concentration, surface area) from ENERGY (temperature, catalyst).
✕Saying catalysts are consumed
Why it happens
Confusing with reactants.
How to avoid it
Catalysts are recovered chemically unchanged at the end.
✕Saying temperature only increases collision frequency
Why it happens
Forgetting the bigger effect.
How to avoid it
Bigger effect is the increased FRACTION of particles with $E \ge E_{a}$ .
✕Saying a catalyst changes the position of equilibrium
Why it happens
Generalising 'speeds up'.
How to avoid it
A catalyst speeds up forward AND reverse equally → position of equilibrium UNCHANGED.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

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Rate of Reaction — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Rate of Reaction

Short Study Notes in the form of Slides

Detailed Study Notes

What you’ll learn

How it’s examined

1Rate from a gas-volume graph

2Why higher concentration speeds up the reaction

3Effect of temperature

4Effect of a catalyst

Rate of reaction

Collision theory

Activation energy (EaE_{a}Ea​)

Catalyst

Surface area

✕Saying 'particles collide more' as the only explanation

✕Saying catalysts are consumed

✕Saying temperature only increases collision frequency

✕Saying a catalyst changes the position of equilibrium

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Rate of Reaction — frequently asked questions

Activation energy ( $E_{a}$ )