GCSE level chemistry notes on 3e. How does using a catalyst affect the speed or rate of a reaction? How does a catalyst work?

Collision theory, reaction profile and lowering activation energy

GCSE level Chemistry Notes: The effect of a catalyst on reaction rate (speed)

3. The Factors affecting the Rate of Chemical Reactions

REACTION RATE and the action of CATALYSTS

Using a CATALYST

3e The effect of a Catalyst (see also light effect and graph 4.8)

(1) Experimental methods for investigating the effect of a catalyst on the rate of a chemical reaction

Parts of the sections of 1. Introduction and 2. collision theory are repeated here, but with extra experimental methods and theoretical details applied to experiments and theories linked to the effect of using a catalyst on the rate of a chemical reaction

• The apparatus can be used to investigate the how the speed of the decomposition of hydrogen peroxide varies with different catalysts.
  • The flask and gas syringe system for measuring the rate of a chemical reaction.
  • Oxygen gas is given off which can be collected in the gas syringe and its volume is used to measure how fast the reaction is going. The grey 'blobs' could represent the solid insoluble catalyst.
  • In the diagram above, the white 'blobs' represent oxygen gas being evolved and the grey lumps the catalyst powder.
  • hydrogen peroxide == catalyst ==> water + oxygen
  • 2H₂O_2(aq) ===> 2H₂O_(l) + O_2(g)
  • MnO₂ Manganese(IV) oxide (manganese dioxide'), is a very effective catalyst, but can also try other transition metal oxides as catalysts like CuO copper(II) oxide.
  • The variables to be kept constant are - the concentration of the hydrogen peroxide solution, the volume of the hydrogen peroxide, the same amount of catalyst (ideally of the same particle size) and the temperature of the reaction mixture.
  • You must also swirl the flask gently to ensure a good mixing as the reaction proceeds.
  • The insoluble particles of the catalyst should be of the same size to give the same surface area of reactant, BUT, this is difficult to achieve in practice at school/college level.
  • More details of laboratory investigations ('labs') involving 'rates of reaction' i.e. experimental methods for observing the speed of a reaction and including the effect of a catalyst are given in the INTRODUCTION

• In this case, measuring the initial rate of gas formation (see left and below diagrams) gives a reasonably accurate measure of how fast the reaction is for that concentration.
• 
• The initial gradient, giving the initial rate of reaction, is the best method i.e. the best straight line covering several results at the start of the reaction by drawing the gradient line using the slope of the tangent from time = 0, where the graph is nearly linear.
• Examples of graph data for two experiments where one of the reactants is completely used up - all reacted.
• The two graph lines represent two typical sets of results to explain how the rate of reaction data can be processed.
• Graph A (for a faster reaction) could represent using a catalyst, but not in Graph B (a slower reaction).

• 

  • The set of graphs (above) shows you some typical results.
  • The rate of reaction order is X > E > Y > Z,
    • X is an effective catalyst,  substances Y and Z seem to slow the reaction, acting as inhibitors.
    • The more effective the catalysts, steeper the initial gradient, the faster the reaction.
    • The more effective the catalyst, the more it lowers the activation energy, the more chance of a successful 'fruitful' collision.
  • For the effect of a catalyst on the rate of reaction, under some circumstances graph W could represent the result of taking twice the mass of solid catalyst or twice the concentration (same volume) of a soluble reactant, BUT it does depend on which reactant is in excess, so take care in this particular graph interpretation.
• More details of laboratory investigations ('labs') involving 'rates of reaction' i.e. experimental methods for observing the speed of a reaction and including the effect of a catalyst are given in the INTRODUCTION
• How can you be sure the 'catalyst' was a true catalyst?
  • The definition of catalyst:
  • A catalyst is a substance that increases the rate of reaction without being chemically changed at the end of the reaction.
  • If you recover the catalyst at the end of a reaction, and then purify it e.g. wash all the reactants off with pure water, after this it should still work as a catalyst.

(2) Theoretical interpretation of the results of the effect of catalyst on the rate of a chemical reaction

For each factor I've presented several particle diagrams to help you follow the text explaining how the particle collision theory accounts for your observations of reaction rate varying with a catalyst (some 'work' better than others!)

A picture of a particles of gaseous molecules or molecules/ions in solution undergoing a chemical changes on the surface of a catalyst - look for the fruitful collision on the catalyst surface.

• WHAT IS A CATALYST?

• HOW DOES A CATALYST AFFECT THE SPEED OF A CHEMICAL REACTION?

• HOW DOES A CATALYST WORK?

• Why does a catalyst speed up a reaction?

• I was once asked "what is the opposite of a catalyst? There is no real opposite to a catalyst, other than the uncatalysed reaction! However, catalysed reactions are very important in industry and most biochemistry involves enzyme catalysts.

• The word catalyst means an added substance, in contact with the reactants, that changes the rate of a reaction without itself being chemically changed in the end. The catalyst may temporarily changed though!

• There are the two phrases you may come across:

  • a 'positive catalyst' meaning speeding up the reaction (plenty of examples in most chemistry courses)

  • OR a 'negative catalyst' slowing down a reaction (rarely mentioned at GCSE, sometimes at AS-A2 level, e.g. adding a chemical that 'mops up' free radicals or other reactive species).

  • One other point, never say anything remotely like 'the catalyst doesn't take part in the reaction', it just speeds it up. It does take a crucial part in the reaction and may be temporarily chemically changed, but the original catalyst does reform to perform its task again i.e. speed up the reaction again!

• Catalysts increase the rate of a reaction by helping break chemical bonds in reactant molecules and provide a 'different pathway' for the reaction.

• This effectively means the Activation Energy E_a is reduced, irrespective of whether its an exothermic or endothermic reaction (see diagrams below for the smaller 'humps' of the lowered activations energy produced by using a catalyst).

• Therefore at the same temperature, using a catalyst, more reactant molecules have enough kinetic energy to react compared to the uncatalysed situation - this increases the frequency of fruitful collisions leading to product formation - more successful collisions per unit time.

  • The purple arrow indicates the greater activation energy for the uncatalysed reaction.

  • The green hump and green arrow indicates the lower activation energy of the catalysed reaction - less particle energy required due to the action of the catalyst.

  • The catalyst does NOT increase the energy of the reactant molecules!

  • Neither does a catalyst increase the frequency of reactant particle collisions.

  • BUT, the catalyst does increase the frequency of 'fruitful' collisions because less kinetic energy is needed by the particles to break bonds and undergo chemical change.

  • Many solution or gaseous catalysed reactions involve a solid catalyst.

    • The reactant molecules are adsorbed onto the surface, and this 'sticking on to the surface' enables the bonds of the reactant molecules to be more easily broken.

    • This is actually what 'lowering of the activation energy' means at the molecular level, how easy is it to break bonds, so we are getting a bit technical here!

• Although a true catalyst does take part in the reaction and may change chemically temporarily, but it does not get used up and can be reused/regenerated with more reactants. It does not change chemically or get used up in the end.

  • Black manganese(IV) oxide (manganese dioxide) catalyses the decomposition of hydrogen peroxide.
  • hydrogen peroxide ===> water + oxygen
    • 2H₂O_2(aq) ===> 2H₂O_(l) + O_2(g)
    • Note that the catalyst is NOT included in the equation.
    • This is because overall the catalyst doesn't change chemically.

  • The manganese dioxide is chemically the same at the end of the reaction but it may change a little physically if its a solid.

  • In the hydrogen peroxide solution decomposition by the solid black catalyst manganese dioxide, the solid can be filtered off when reaction stops 'fizzing' i.e. all of the hydrogen peroxide has reacted-decomposed.

  • After washing with water, the catalyst can be collected and added to fresh colourless hydrogen peroxide solution and the oxygen production 'fizzing' is instantaneous! In other words the catalyst hasn't changed chemically and is as effective as it was fresh from the bottle!

    • Note: At the end of the experiment the solution is sometimes stained brown from minute manganese dioxide particles. The reaction is exothermic and the heat has probably caused some disintegration of the catalyst into much finer particles which appear to be (but not) dissolved. In other words the catalyst has changed physically BUT NOT chemically.

    • You can try salts of transition metals (notable for catalytic effects) e.g. copper sulfate, iron sulfate or cobalt sulfate to see if they have any effect on the rate of decomposition of hydrogen peroxide. Its best to use sulfate salts rather than chlorides because the hydrogen peroxide may oxidise the chloride ion and give rise to further unintentioned reactions.

  • You can investigate the rate of reaction of hydrochloric acid with zinc and find that adding a few granules of copper or a few drops of copper sulfate solution speeds of the reaction, which you can measure the rate of reaction by following the rate of evolution of hydrogen gas.

• Using an effective catalyst can reduce costs in the chemical and food industries by increasing the rate of reaction (more efficient) and lowering the energy requirements if the process can be done at lower temperatures.
  • Increasing the rate of reaction saves time and operating at a lower temperature saves energy and therefore saves money.
  • However, catalysts can be very specialised and expensive to produce.
  • They also get contaminated ('poisoned') and become less efficient and might have to be extracted and cleaned up, but if a true catalyst, this' refurbishment' should enable the catalyst to be reused (theoretically catalysts take part in the reaction, but are not consumed in the reaction).
    • Sulfur compounds poison the iron catalyst used in the Haber synthesis of ammonia.
    • One way of minimising the poisoning-contamination of catalysts is to purify the reactant molecules before they enter the chemical reactor chamber.
  • See 'chemical economics' for other commercial aspects of chemical production.

• Different reactions need different catalysts and they are all extremely important in industry:

  • Catalysts make chemical industrial processes much more efficient and economic e.g.

  • Nickel catalyses the hydrogenation of unsaturated fats to margarine

  • iron catalyses the combination of unreactive nitrogen and hydrogen to form ammonia in the Haber Synthesis.

  • Zeolite minerals catalyse the cracking of big hydrocarbon molecules into smaller ones

  • Most polymer making reactions require a catalyst surface or additive in contact with or mixed with the monomer molecules.

  • Catalysts in car exhausts (catalytic converters) change harmful carbon monoxide and nitrogen monoxide into harmless carbon dioxide and nitrogen, they also convert unburned potentially carcinogenic hydrocarbons into carbon dioxide and water.

• Enzymes are biochemical catalysts are dealt with on other pages

  • e.g. enzymes in yeast convert sugar into alcohol and are important in digestion chemistry

  • They have the advantage of bringing about reactions at normal temperatures and pressures which would otherwise need more expensive and energy-demanding equipment.

  • Enzymes have ability to reduce the activation energy needed (See rates of reaction notes)

  •  

  • The reaction profile diagrams illustrate the point.

  • By lowering the activation energy, more molecules have enough kinetic energy to react - bonds are more easily broken - so greater probability of fruitful collision forming the products of the reaction.

  • Activation energy and reaction profiles  and for enzymes - natural biochemical catalysts

  • Enzymes - biological catalysts, optimum pH or temperature, use in biotechnology (GCSE chemistry)

  • Enzymes - structure, function, optimum conditions, investigation experiments (GCSE biology notes)

  • These provide full explanations of the key and lock mechanism illustrated below.

Enzyme catalysed synthesis of a larger molecule from two smaller ones.

Enzyme catalysed decomposition a larger molecules into two smaller molecules.

For more advanced details on catalysis see

Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

(3) APPENDIX - activation energy, catalysts and reaction profiles

(4) Key revision points about catalysts and the rate of chemical reactions

A structured set of brief revision notes on rates of reaction – effect of catalysts, aligned to the major GCSE and IGCSE chemistry exam boards.

They include fair test design, collision theory with activation energy, analysis of results, and exam tips/misconceptions to sharpen exam preparation.

Rates of Reaction – Effect of Catalysts

1. Core Concept

• Catalyst: A substance that increases the rate of a chemical reaction without being used up.
• Catalysts provide an alternative reaction pathway with a lower activation energy.
• They do not change the yield or equilibrium position, only the speed.

2. Experimental Investigation

Common Practical Examples

• Decomposition of hydrogen peroxide:
  • With manganese(IV) oxide catalyst → oxygen gas produced faster.
  • Measure volume of oxygen over time.
• Reaction of zinc with hydrochloric acid:
  • Compare rate with and without copper(II) sulfate catalyst.
• Iodine clock reaction:
  • Compare time taken for colour change with and without catalyst.

Fair Test Design

• Keep concentration, volume, temperature, surface area constant.
• Only change presence/type of catalyst.
• Use same apparatus and same measurement method.
• Repeat experiments for reliability and calculate averages.

3. Analysing Results

• Plot rate versus catalyst presence.
• Expect faster reaction when catalyst is present (steeper curve, shorter time).
• Rate can be calculated as: Rate = change in quantity/time
• Compare curves: catalysed reaction reaches completion faster but same final product amount.

4. Collision Theory & Activation Energy

• Collision theory: Particles must collide with sufficient energy and correct orientation.
• Catalyst effect:
  • Provides alternative pathway with lower activation energy.
  • More collisions have energy ≥ new (lower) activation energy.
  • Reaction rate increases even though collision frequency may be unchanged.

5. Typical Exam Board Specification Content

• Required practicals include catalysts (hydrogen peroxide decomposition).
• Emphasis on activation energy diagrams and catalyst role.
• Focus on fair test rates investigation design and interpreting catalysed versus uncatalysed curves.
• Application of collision theory and activation energy to catalyst experiments.
• Graphical analysis and practical design catalyst rate experiment investigations.
• Explanation using collision theory and the idea of activation energy.

6. Student Exam Tips

• Always mention controlled variables when describing experiments.
• Use activation energy language: “Catalyst lowers activation energy, so more particles can react.”
• Draw energy profile diagrams: show lower activation energy pathway.
• Quote data when analysing graphs (e.g., “With catalyst, reaction finished in 30s compared to 90s without”).
• Units matter: seconds (s), cm³/s, etc.
• Remember catalysts are not used up – they remain chemically unchanged at the end.

7. Typical Misconceptions

•  Thinking catalysts increase collision frequency (they lower activation energy, not speed particles).
•  Believing catalysts are consumed in the reaction (they are not).
•  Assuming catalysts change yield or equilibrium (they only affect rate).
•  Forgetting to control other variables (e.g., concentration, temperature).
•  Confusing “rate of reaction” with “amount of product formed.”

8. Quick Revision Table

GCSE level 'Rates of Reaction' multiple choice quiz