(c) doc bDoc Brown's Chemistry KS4 science GCSE/IGCSE/O Level Revision Notes

REACTION RATE and CATALYSTS

Factors affecting the Speed-Rates of Chemical Reactions

3e. How does using a catalyst affect the speed or rate of a reaction?

How can we investigate the effect of a catalyst on the rate of a reaction? What is a catalyst? Why can a catalyst change the speed of a chemical reaction? How can we investigate the effect of concentration on the rate of a chemical reaction? What apparatus do we need to investigate the effect of a catalyst on the speed of a reaction? How do we process the results from catalysed experiments?  These revision notes are suitable for GCSE IGCSE O Level KS4 science chemistry students studying the effects of catalysts and how they work. The descriptions of experiments to do with catalysed reactions and theoretical explanations should help with homework, coursework assignments, laboratory experiment investigations 'labs' on catalytic effects on the speed or rate of a chemical reaction. These notes on the effects of catalysts on reaction rates and the experimental methods and investigations involved, are designed to meet the highest standards of knowledge and understanding required for students/pupils doing GCSE chemistry, IGCSE chemistry, O Level chemistry, KS4 science courses and can be useful primer for A Level chemistry courses. These revision notes on the effect of catalysts on the rate of a reaction speed should prove useful for the new AQA GCSE chemistry, Edexcel GCSE chemistry & OCR GCSE chemistry (Gateway & 21st Century) GCSE (91), (9-5) & (5-1) science courses.


RATES of REACTION INDEX: 1. What do we mean by rate/speed of reaction? and how can we measure the speed?  2. Collision theory of chemical reactions  3. Factors affecting the speed of a chemical reaction: 3a Effect of changing concentration  3b Effect of changing pressure  3c Effect of changing particle size/surface area & stirring of a solid reactant  3d Effect of changing temperature  3e Effect of using a catalyst  3f Light initiated reactions 4. More examples of graphs and their interpretation  *  A Level pages on rates - chemical kinetics


3. The Factors affecting the Rate of Chemical Reactions

Using a CATALYST

(c) doc b3e The effect of a Catalyst (see also light effect and graph 4.8)


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

  • Factors affecting the rates of Reaction - theory and methods of measuring the speed of a reaction (c) Doc Brown
  • 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
    • 2H2O2(aq) ====> 2H2O(l) + O2(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).
  • (c) doc bThe set of graphs (left) shows you some typical results.
    • The rate of reaction order is X > E > Y > Z, and could represent four increasingly effective catalysts.
      • 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

 


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

  • 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 Ea 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).

Factors affecting the rates of Reaction - particle collision theory model (c) Doc Brown Factors affecting the rates of Reaction - particle collision theory model (c) Doc Brown

  • Therefore at the same temperature, using a catalyst, more reactant molecules have enough kinetic energy to react compared to the uncatalysed situation.

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

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

    • 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
      • 2H2O2(aq) ==> 2H2O(l) + O2(g)
    • The manganese dioxide is chemically the same at the end of the reaction but it may change a little physically if its a solid e.g.

    • 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.

  • 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

    • Enzymes in yeast convert sugar into alcohol

    • 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 another page - (c) doc b enzymes and biotechnology

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

  • For more details on catalysis see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

APPENDIX - activation energy, catalysts and reaction profiles

 

Reaction profile diagram

Comments related to the reaction activation energy and use of a catalyst
An exothermic reaction with a small activation energy

The reaction may go very well without a catalyst at a practical temperature, perhaps even at room temperature

An exothermic reaction with a moderately high activation energy.

This reaction might benefit from using a catalyst if a suitable one is available

An endothermic reaction with a big activation energy

This reaction would benefit from using a catalyst e.g. to avoid using an excessively high temperature, catalysts used to crack crude oil into useful fractions

 


RATES of REACTION INDEX: 1. What do we mean by rate/speed of reaction? and how can we measure the speed?  2. Collision theory of chemical reactions  3. Factors affecting the speed of a chemical reaction: 3a Effect of changing concentration  3b Effect of changing pressure  3c Effect of changing particle size/surface area & stirring of a solid reactant  3d Effect of changing temperature  3e Effect of using a catalyst  3f Light initiated reactions 4. More examples of graphs and their interpretation  *  A Level pages on rates - chemical kinetics



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