3d. What is the effect of temperature on the rate or speed of a chemical reaction?

Doc Brown's Chemistry KS4 science GCSE/IGCSE/O level Revision Notes - Factors affecting the Speed-Rates of Chemical Reactions - Doc Brown's chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE  chemistry, O level & ~US grades 8, 9 and 10 school science courses or equivalent for ~14-16 year old science students for national examinations in chemistry

Rates of reaction notes INDEX

3. The Factors affecting the Rate of Chemical Reactions

REACTION RATE and TEMPERATURE

Varying the TEMPERATURE of the reactants - increase or decrease

• Experimental methods for investigating the effect of temperature on the rate of a chemical reaction.

• 
• (i) The above diagram illustrates how you can investigate how varying the temperature affects the rate at which it reacts with a given quantity of limestone granules - but water bath not shown.
  • The apparatus set-up illustrated above is fine for an initial room temperature base-line experiment, but for higher temperatures, its a bit tricky - some ideas are described with another diagram further down the page.
  • calcium carbonate (marble chips)  + hydrochloric acid ==> calcium chloride + water + carbon dioxide
  • CaCO_3(s) + 2HCl_(aq) ===> CaCl_2(aq) + H₂O_(l) + CO_2(g)
  • In the diagram above, the white 'blobs' represent carbon dioxide gas being evolved and the grey lumps the limestone chips, granules or powder.
  • You follow the reaction by measuring the volume of carbon dioxide formed using the gas syringe system.
  • You must keep the following variables constant - the volume of hydrochloric acid, the mass of limestone AND its particle size, the concentration of the hydrochloric acid and TRY to keep a gentle constant stirring rate as you are noting down the time and volume of carbon dioxide gas formed.
  • Gentle stirring (swirling action) is important, if you don't, the bottom layers of acid become depleted in acid giving a falsely slow rate of reaction.
  • So, the only factor you should vary in the temperature of the reactants and flask.
  • You repeat the experiment at different temperatures to see the effect of temperature on the rate-speed of the reaction between hydrochloric acid and limestone/marble chips-powder.
  • 
  • Keeping the temperature constant, particularly at temperatures above room temperature (ambient temperature), is quite a problem which is best solved by using a thermostatically controlled large water bath.
    • The flask and gas syringe system for measuring the rate of a chemical reaction.
    • If you have access to a thermostated water bath in which to the different temperature experiments, that's great, if not, I've suggested a few ideas below.
    • Before you had the solid reactant (marble) or a solid catalyst, depending on the experiment, you must allow the conical flask of solution (acid, hydrogen peroxide etc.) to reach the ambient temperature of the water bath.
    • If no thermostat system is available, there are two simpler, but not as accurate alternatives.
      • If the experiment doesn't take too long, a large beaker or trough of water might do, its temperature monitored with a thermometer, but still allow time for the conical flask and contents to warm up to the same temperature. Since cooling is taking place all the time you need to warm up the water a few degrees above the desired reaction temperature. You can also warm up the conical flask solution to the same temperature independently. You should take the temperature at the start of the reaction and at the end and use the mean value for your final results table.
      • You can, least accurately of all, measure the temperature of the solution at the start of the reaction (before adding solid) and re-measure at the end of the experimental time allotted, and use the average temperature. Its not that accurate, but its better than nothing and you should still be able to derive the general trend of how temperature affects the speed-rate of a reaction.

• (ii) The same apparatus can be used to investigate the how the speed of the decomposition of hydrogen peroxide varies at different temperatures in the presence of a fixed amount of catalyst (e.g. manganese(IV) oxide, manganese dioxide).
  • hydrogen peroxide ===> water + oxygen
  • 2H₂O_2(aq) ===> 2H₂O_(l) + O_2(g)
  • You must keep the following variables constant - the volume of hydrogen peroxide solution, the concentration of the hydrogen peroxide, the mass of catalyst AND its particle size, and TRY to keep a gentle constant stirring rate as you are noting down the time and volume of carbon dioxide gas formed.
  • Gentle stirring is important, if you don't, the bottom layers of hydrogen peroxide become depleted in acid giving a falsely slow rate of reaction.
  • You follow the reaction by measuring the volume of oxygen gas formed.
  • You repeat the experiment at different temperatures using the same volume and concentration of hydrogen peroxide and mass of the same catalyst to see the effect of temperature on the rate-speed of the catalysed decomposition of hydrogen peroxide.

• 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 temperature are given in the INTRODUCTION
• In both these cases, 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 a higher temperature than 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, and could represent four increasing temperatures for fixed amounts of solid and concentration of reactants.
  • The greater the temperature, the steeper the initial gradient, the faster the reaction.
• For the effect of temperature on the rate of reaction, under some circumstances graph W could represent the result of taking twice the mass of solid reactant (e.g. double amount of marble chips) 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.
• The graph on the right shows how the initial rate varies with increase in temperature of the reaction mixture. The reciprocal of the reaction time can be taken as a measure of the speed of the reaction at that particular temperature.

• (iii) Investigating how varying the concentration of either sodium thiosulfate or hydrochloric acid affects the rate they react together to give a precipitate of sulfur

• mix => ongoing => watch stopped =>

  • You must keep the volumes of reactants constant, the concentrations of the hydrochloric acid and sodium thiosulfate constant, and the same person making all the observations with the same size cross on white paper.

  • You must pre-heat both solutions to the desired temperature, separately, before mixing them together.

  • It is important you take the same total volumes of reactant solutions to give the same depth of liquid you are viewing the cross through.

  • Everything should be mixed quickly and the clock started, but there is no need to stir the mixture once it is fully mixed.
  • You note the time when the cross first disappears.
  • To vary the temperature of the reactant solution mixture, you will have to pre-heat the solutions, mix them, start the clock AND take the temperature. The reaction mixture may cool a little, so you can re-take the temperature at the end and use the average, not perfect, but more accurate than either temperature reading.
  • You repeat the experiment at different temperatures to see their effect on the rate-speed of the acid promoted decomposition of sodium thiosulfate to form a sulfur precipitate. Typical results are shown below.

  • temperature (oC)	20	25	30	35	40	45
    time for X to be obscured (s)	240	220	190	150	100	40

  • The graphs on the left shows how the reaction time and rate varies to obscure the X with increase in temperature of the hydrochloric acid and sodium thiosulfate mixture. The reciprocal of the reaction time can be taken as a measure of the speed of the reaction at that particular temperature.
• 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 temperature are given in the INTRODUCTION

• See also graphs 4.6, 4.7 and 4.8 for a numerical-quantitative data interpretation AND the introduction page

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 the temperature of the reaction system (some 'work' better than others!)

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

• IF SO, HOW AND WHY?

• Why does a reaction go faster at a higher temperature?

  • The greater the temperature of the reactants, the greater the average kinetic energy of the particles.
  • Therefore, the more chance of a successful more energetic 'fruitful' collision between two particles with sufficient combined kinetic energy to overcome the activation energy barrier, break bonds and form the products.
  • The frequency of collision increases too, but this is the lesser of the two factors which both contribute to an increased rate of reaction on raising the temperature.

    • We are talking about an increased frequency of fruitful collisions, mainly because of the higher average kinetic energy of the molecules - leading to product formation.

    • Increased frequency means an increase in 'rate/speed of reaction'.

• When gases or liquids are heated the particles gain kinetic energy and on average move faster (see diagrams below).

• The increased speed increases the chance (frequency) of collision between reactant molecules and the rate of reaction increases.

• BUT this is NOT the main reason for the increased reaction speed, so be careful in your theory explanations if investigating the effect of temperature, so read on after the pictures!

  • The more important factor is the kinetic energy of the particles and therefore the higher energy collisions when the temperature is raised.

• Most molecular collisions do not result in chemical change.

• To further explain the effect of temperature on the speed of a reaction, we need to consider the reaction profile and the importance of the activation energy.

• 

• Before any change takes place on collision, the colliding molecules must have a minimum kinetic energy called the Activation Energy shown on the energy level diagrams below (sometimes called reaction profile/progress diagrams - shown below).

  • Going up and to the top 'hump' represents bond breaking on reacting particle collision.

    • The purple arrow up represents this minimum energy needed to break bonds to initiate the reaction, that is the activation energy.

  • Going down the other side represents the new bonds formed in the reaction products. The red arrow down represents the energy released - exothermic reaction.

• It does not matter whether the reaction is an exothermic or an endothermic in terms of energy change, its the activation energy which is the most important factor in terms of temperature and its effect on reaction speed.

• Now 'hotter' molecules have a greater average kinetic energy, and so at higher temperatures, a greater proportion of them have the required activation energy to react i.e. their combined kinetic energy on collision is sufficient to break open bonds and allow the reaction to proceed to product formation.

• This means that the increased chance (frequency) of 'fruitful' higher energy collisions in a given time greatly increases the speed of the reaction, depending on the fraction of molecules with enough energy to react.

• For this reason, generally speaking, and in the absence of catalysts or extra energy input, a low activation energy reaction is likely to be fast and a high activation energy reaction much slower, reflecting the trend that the lower the energy barrier to a reaction, the more molecules are likely to have sufficient energy to react on collision.

• See also graphs 4.3, 4.4 and 4.8 for a numerical-quantitative data interpretation for a numerical-quantitative data interpretation AND the introduction page

  • Industrial note on the effect of temperature on the rate of reaction

    • In industry you would try to run the reaction at the highest economic temperature.

    • BUT, the energy bill should not be economically demanding,

    • the speed of reaction at higher temperature must be under control and not too fast to be dangerous,

    • if the reaction is exothermic and an equilibrium is formed, operating at too high a temperature might significantly reduce the yield.

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  APPENDIX Trying to resolve an apparent confusion for a GCSE (or A Level) student!

1. With increase in temperature, there is an increased frequency (or chance) of collision due to the more 'energetic' situation - but this is the minor factor when considering why rate of a reaction increases with temperature.
2. The minimum energy needed for reaction, the activation energy (to break bonds on collision), stays the same on increasing temperature.
   • However, the average increase in particle kinetic energy caused by the absorbed thermal energy means that a much greater proportion of the reactant molecules now has the minimum or activation energy to react.
   • So, at a higher temperature, there are more particles with the higher kinetic energies.
   • Therefore there will be more particles colliding with enough energy to overcome the threshold activation energy.
3. It is this increased chance of a 'successful' or 'fruitful' higher energy collision leading to product formation, that is the major factor, and this effect increases more than the increased frequency of particle collision, for a similar rise in temperature.
4. This is usually only fully discussed at Advanced A level Chemistry, but it may impress the teacher for GCSE coursework if you look up the Maxwell-Boltzmann distribution of kinetic energies, though its quite difficult to get over some of these ideas without considering graphs of probability versus particle KE, but that's up to you!
   • There is also the Arrhenius Equation relating rate of reaction and temperature - but this involves advanced level mathematics.

Rates of reaction notes INDEX

GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction

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