Kinetics
 

 

Factors affecting the Speed-Rates of Chemical Reactions

 

The effect of Concentration (see also graphs 4.6, 4.7)

  • If the concentration of any reactant in a solution is increased, the rate of reaction is increased

    • Increasing the concentration, increases the probability of a collision between reactant particles because there are more of them in the same volume and so increases the chance of a fruitful collision forming products.

    • e.g. Increasing the concentration of acid molecules increases the frequency or chance at which they hit the surface of marble chips to dissolve them (slower => faster, illustrated below)

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  • In general, increasing the concentration of reactant A or B will increase the chance or frequency of a successful collision between them and increase the speed of product formation (slower => faster, illustrated below).

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  • Increasing the concentration of reactant A or B will increase the chance or frequency of collision between them and increase the speed of product formation (slower => faster).

  • See also graphs 4.6, 4.7 for a numerical-quantitative data interpretation.

    (c) doc bGraph 4.6 shows the effect of increasing concentration, which decreases the reaction time, as the speed increases because the greater the concentration the greater the chance of fruitful collision. See the notes on rate in the graph 4.3

     

    (c) doc bGraph 4.3 shows the decrease in reaction time with increase in temperature as the reaction speeds up. The reaction time can represent how long it takes to form a fixed amount of gas in e.g. in the first few minutes of a metal/carbonate - acid reaction, or the time it takes for so much sulfur to form in the sodium thiosulphate - hydrochloric acid reaction. The time can be in minutes or seconds, as long as you stick to the same unit for a set of results e.g. a set of experiments varying the concentration of one of the reactants.

    (c) doc bGraph 4.7 shows the rate/speed of reaction is often proportional to the concentration of one particular reactant. This is due to the chance of a fruitful collision forming products being proportional to the concentration. The initial gradient of the product-time graph e.g. for gas in cm3/min (or s for timing the speed/rate) gives an accurate measure of how fast the gaseous product is being formed.  The reciprocal of the reaction time, 1/time, can also be used as a measure of the speed of a reaction. The time can e.g. represent how long it takes to make a fixed amount of gas, or the time it takes for so much sulphur to form in the sodium thiosulphate - hydrochloric acid reaction. The time can be in minutes or seconds, as long as you stick to the same unit for a set of results for a set of experiments varying the concentration or mass of one of the reactants.

      


The effect of Pressure
  • If one or more of the reactants is a gas then increasing pressure will effectively increase the concentration of the reactant molecules and speed up the reaction.

  • The A and B particle diagrams above could represent lower/higher pressure , resulting in lesser è greater concentration and so slower è faster reaction all because of the increased chance of a 'fruitful' collision.


 The effect of Stirring
  • In doing rate experiments with a solid and solution reactant e.g. marble chips-acid solution or a solid catalyst like manganese(IV) oxide catalysing the decomposition of hydrogen peroxide solution, it is sometimes forgotten that stirring the mixture is an important rate factor.

  • If the reacting mixture is not stirred ‘evenly’, the reactant concentration in solution becomes much less near the solid, which tends to settle out at the bottom of the flask.

  • Therefore, at the bottom of the flask the reaction prematurely slows down distorting the overall rate measurement and making the results uneven and therefore inaccurate. The 'unevenness' of the results is even more evident by giving the reaction mixture the 'odd stir'! You get jumps in the graph!!!

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 The effect of Surface Area - particle size of a solid reactant

  • If a solid reactant or a solid catalyst is broken down into smaller pieces the rate of reaction increases.

  • The speed increase happens because smaller pieces of the same mass of solid have a greater surface area compared to larger pieces of the solid.

  • Therefore, there is more chance that a reactant particle will hit the solid surface and react.

  • The diagrams below illustrate the acid–marble chip reaction (slower => faster, but they could also represent a solid catalyst mixed with a solution of reactants.

  • See also graphs 4.1 for a numerical-quantitative data interpretation.

    (c) doc bGraph 4.1 shows the decrease in the amount of a solid reactant with time. The graph is curved, becoming less steep as the gradient decreases because the reactants are being used up, so the speed decreases. Here the gradient is a measure of the rate of the reaction. In the first few minutes the graph will (i) decline less steeply for larger 'lumps' and (ii) decline more steeply with a fine powder i.e. (i) less surface area gives slower reaction and (ii) more surface area a faster reaction.
     

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The effect of Temperature (see also graphs 4.3, 4.4)

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

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

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  • Most molecular collisions do not result in chemical change.

  • 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 reaction speed.

  • Now heated molecules have a greater average kinetic energy, and so at higher temperatures, a greater proportion of them have the required activation energy to react.

  • This means that the increased chance of 'fruitful' higher energy collision 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.

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    • See also graphs 4.3 and 4.4 for a numerical-quantitative data interpretation

       

  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 heat means that a much greater proportion of the reactant molecules now has the minimum or activation energy to react. 
  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.
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    (c) doc bGraph 4.3 shows the decrease in reaction time with increase in temperature as the reaction speeds up. The reaction time can represent how long it takes to form a fixed amount of gas in e.g. in the first few minutes of a metal/carbonate - acid reaction, or the time it takes for so much sulphur to form in the sodium thiosulphate - hydrochloric acid reaction. The time can be in minutes or seconds, as long as you stick to the same unit for a set of results e.g. a set of experiments varying the concentration of one of the reactants.
    (c) doc b Graph 4.4 shows the increase in speed of a reaction with increase in temperature as the particles have more and more kinetic energy. The rate of reaction is proportional to 1/t where t = the reaction time. 
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The effect of a Catalyst  (see also the graph 4.8 at the bottom of the page)
  • I was once asked "what is the opposite of a catalyst? There is no real opposite to a catalyst, other than the uncatalysed reaction!

  • 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.

  • 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 or other reactive species).

  • 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 is reduced, irrespective of whether its an exothermic or endothermic reaction (see diagrams below).

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  • Therefore at the same temperature, 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!

  • 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 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.

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

    • 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

    • zeolites 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.

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