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8.4–8.5 The Effect of Concentration, Pressure and Catalysts

GCSE Chemistry · AQA · Chapter 8: Rates and Equilibrium

A free GCSE Chemistry lesson for AQA (8462) from OBA School: how crowding particles together speeds up reactions, and how catalysts give reactions an easier route. Watch the video, read the explanation and worked examples, then try the revision questions. You can work through the full lesson, with every video, practice quizzes and printable worksheets with mark schemes, in our free online course.

What you will learn

  • Explain how increasing the concentration of a solution or the pressure of a gas increases the rate
  • Interpret data showing how rate depends on concentration, including proportional relationships
  • Describe what a catalyst is and explain how it works using a reaction profile
  • Explain why catalysts are important in industry, and recall that enzymes are biological catalysts
  • Find and compare initial rates by drawing tangents at t = 0 (Higher tier)

Key words

  • Concentration: The amount of a substance dissolved in a given volume of solution, e.g. in mol/dm³ or g/dm³.
  • Pressure (of a gas): Caused by gas particles hitting the walls of their container. More particles in the same volume means a higher pressure.
  • Catalyst: A substance that changes the rate of a chemical reaction but is not used up in the reaction.
  • Enzyme: A protein that acts as a biological catalyst.
  • Alternative reaction pathway: A different route from reactants to products, with a lower activation energy.

Concentration and pressure

In a more concentrated solution there are more reactant particles in the same volume. For gases, a higher pressure squeezes the same particles into a smaller volume, which has the same effect. Either way the particles are more crowded, so they collide more frequently, and the rate increases.

Key fact: Concentration and pressure change the frequency of collisions, not their energy. The proportion of collisions that are successful stays the same.

Concentration data and proportionality

Zinc granules were added to sulfuric acid of different concentrations. The time to collect 30 cm³ of hydrogen was measured.

  • Time to collect 30 cm³ (s); 0.25 200; 0.50 100; 1.00 50
  • Mean rate (cm³/s); 0.25 0.15; 0.50 0.30; 1.00 0.60

Worked example: spotting a proportional relationship

mean rate = 30 ÷ 100 = 0.30 cm³/s at 0.50 mol/dm³, and 30 ÷ 50 = 0.60 cm³/s at 1.00 mol/dm³.
Doubling the concentration (0.50 → 1.00) doubles the rate (0.30 → 0.60).

Answer: rate is directly proportional to concentration (rate ∝ concentration)

Exam tip: A graph of rate against concentration for a proportional relationship is a straight line through the origin. Use the word “proportional” only if both conditions are true.

Comparing initial rates (Higher tier)

The reaction is fastest at the very start, when the concentration of reactants is highest. To compare experiments fairly, find the initial rate: draw a tangent to each curve at t = 0 (it starts at the origin) and work out its gradient.

Worked example: initial rate from a tangent

1.0 mol/dm³: the tangent goes from (0, 0) to (40 s, 60 cm³). Initial rate = 60 ÷ 40 = 1.5 cm³/s.
0.5 mol/dm³: the tangent goes from (0, 0) to (80 s, 60 cm³). Initial rate = 60 ÷ 80 = 0.75 cm³/s.
Doubling the concentration doubles the initial rate: twice as many H⁺ ions in the same volume, so twice as many collisions per second with the zinc surface.

Answer: initial rate is directly proportional to the acid concentration

Both curves level off at the same volume because the zinc is the limiting reactant and the same mass was used. If the acid were the limiting reactant, halving its concentration would also halve the final volume.

Catalysts and how they work

A catalyst increases the rate of a reaction but is not used up: the same mass is left at the end, so a small amount can be used again and again. Catalysts are not included in the balanced equation, but you can write the catalyst above the arrow: 2H₂O₂ → 2H₂O + O₂ with “MnO₂” written over the arrow. Different reactions need different catalysts. Many industrial catalysts are transition metals or their compounds.

A catalyst provides a different pathway for the reaction that has a lower activation energy. At the same temperature, a greater proportion of collisions now have enough energy to react, so the frequency of successful collisions goes up.

Worked example: testing a catalyst

Hydrogen peroxide decomposes slowly: 2H₂O₂(aq) → 2H₂O(l) + O₂(g).
Adding 0.50 g of manganese(IV) oxide powder makes it fizz quickly. 84 cm³ of oxygen is collected in 60 s, so mean rate = 84 ÷ 60 = 1.4 cm³/s.
Afterwards the solid is filtered off, washed, dried and weighed: 0.50 g.

Answer: MnO₂ speeds up the reaction and is not used up, so it is a catalyst

Catalysts in industry and in living things

  • Iron; Used for Making ammonia from nitrogen and hydrogen (the Haber process).
  • Platinum (with rhodium) gauze; Used for Oxidising ammonia in the first stage of making nitric acid (for fertilisers and explosives).
  • Platinum, rhodium, palladium; Used for Catalytic converters in cars: change carbon monoxide, nitrogen oxides and unburnt fuel into less harmful gases. The metal is a thin coating on a honeycomb to give a huge surface area.
  • Nickel; Used for Adding hydrogen to vegetable oils to harden them (e.g. for spreads).
  • Enzymes; Used for Biological catalysts: digestion, biological washing powders, making bread and yoghurt.

Solid catalysts are used as fine powders, small pellets (often with holes through them) or thin wire gauzes. These shapes give a large surface area to volume ratio, so gas molecules collide with the catalyst surface more often and less of the expensive metal is needed.

Why industry uses catalysts

• Products are made faster, so the plant makes more in a day.

• Reactions can run at lower temperatures, saving energy and money and cutting carbon dioxide emissions from burning fuels.

• They are not used up, so they can be reused. Drawbacks: some are expensive or toxic, and impurities can “poison” them: an impurity in the reactants binds to the catalyst surface so it stops working and must be cleaned or replaced.

Common mistakes

  • Saying a catalyst gives the particles more energy. It lowers the activation energy by providing a different pathway.
  • Saying a catalyst is “not involved” in the reaction. It takes part, but it is regenerated, so it is not used up.
  • Saying a catalyst increases the amount of product. It only makes the same amount faster.
  • Saying higher concentration means “more collisions”. Say more frequent collisions (more per second).
  • Using “proportional” when the graph is a straight line that does not pass through the origin.

Revision questions

Try these, then check your answers below each one.

Question 1. What is a catalyst? (Choose from: A substance that is used up to speed up a reaction, A substance that changes the rate of a reaction without being used up, A substance that increases the amount of product, A substance that gives particles more energy.)

Answer: A substance that changes the rate of a reaction without being used up

Working: A catalyst speeds up a reaction but is not used up, so it can be reused.

Question 2. True or false? A catalyst is written as a reactant in the balanced equation for a reaction.

Answer: False

Working: False. Catalysts are not used up, so they are not included in the balanced equation (they may be written above the arrow).

Question 3. What name is given to biological catalysts?

Answer: enzymes

Working: Enzymes are biological catalysts, e.g. amylase.

Related lessons

Continue in the full course

This lesson is section 8.4–8.5 (AGCCH0804) of our free GCSE Chemistry course for AQA. In the full course you get every video for this section, a self-marking practice quiz, a printable worksheet with a mark scheme, chapter tests, and your progress saved. Go to the full GCSE Chemistry course.

 

 

 

 

 

 

 

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