Doc Brown's Chemistry KS4 science GCSE/IGCSE Revision Notes

Factors affecting the Speed-Rates of Chemical Reactions

This page describes the factors controlling the speeds of chemical reactions and the collision theory behind it discussed. The factors affecting the speed of reaction are also presented using particle models to give a theoretical basis to the rules on the effects of concentration, pressure, temperature, solid reactant particle size (surface area), stirring, catalysts and light. Methods of how to collect data are also described and graphical treatment of the observations and how to draw conclusions. How do we graphically interpret tables of results and what graphs are useful and why. How to we calculate the speed of a reaction?

SECTIONS on this page: 1. What do mean by rate/speed of reaction and its measurement? * 2. Collision theory of reaction * 3. Factors: 3a concentration, 3b pressure, 3c stirring, 3d particle size/surface area, 3e temperature, 3f catalyst, 3g light * 4. Examples of graphs

KEY WORDS-PHRASES in alphabetical order for this rates web page: hydrochloric/sulphuric acid-metal e.g. Mg/carbonate reaction * hydrochloric acid-sodium thiosulphate reaction * Activation energy * Catalysts * Concentration effect * Graphs-gas collection * Graphs-examples * hydrogen peroxide decomposition * How reactions happen * Interpreting results * Light (catalyst) effect *  Methods of measuring rate * Pressure effect * Rate of reaction * Reaction profiles *  Stirring effect * Surface area/size of solid particle reactant effect * Temperature effect

See also Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction

1. What do we mean by Rate and how is it measured?

• WHAT DO WE MEAN BY SPEED OR RATE IN THE CONTEXT OF A CHEMICAL REACTION?

• IS IT TO FAST OR TO SLOW TO MEASURE THE SPEED?

• WHAT SORT OF WAYS CAN WE MEASURE THE SPEED OF A CHEMICAL REACTION?

• The phrase ‘rate of reaction’ means ‘how fast is the reaction’ or 'the speed of the reaction'. It can be measured as the 'rate of formation of product' (e.g. collecting gaseous product in a syringe) or the 'rate of removal of reactant'. The speeds of reactions are very varied.

  • Rusting is a ‘slow’ reaction, you hardly see any change looking at it!

  • The weathering of rocks is an extremely very slow reaction.

  • The fermentation of sugar to alcohol is quite slow but you can see the carbon dioxide bubbles forming in the 'froth' in a laboratory experiment or beer making in industry!

  • A faster reaction example is magnesium reacting with hydrochloric acid to form magnesium chloride and hydrogen or the even faster reaction between sodium and water to form sodium hydroxide.

  • Combustion reactions e.g. when a fuel burns in air or oxygen, is a very fast reaction.

    • A 'use of words' revision note: Reacting and/or dissolving? Chemical or physical change?

      • If you take the solids magnesium chloride or sodium hydroxide and mix them with water they dissolve to form a solution, but no chemical reaction to form new substances takes place i.e. dissolving on its own is basically a physical change.

      • However, the two substances mentioned above are formed in a chemical reaction change, where the word 'dissolving' on its own is inadequate. The phrases reaction with ... or reaction between ... are much more appropriate, but there is no denying that the magnesium/sodium dissolve in acid/water, BUT only because they have formed a water soluble compound.

  • Explosive reactions would be described as ‘very fast’ e.g. the pop of a hydrogen-air mixture on applying a lit splint or the production of a gas to inflate the air bags safety feature of many cars.

• The importance of "Rates of Reaction knowledge":

  • Time is money in industry, the faster the reaction can be done, the more economic it is.

    • You need to know how long reactions are likely to take.

    • Hence the great importance of catalysts e.g. transition metals or enzymes which reduce time and save money.

  • Health and Safety Issues:

    • Mixtures of flammable gases in air present an explosion hazard (gas reactions like this are amongst the fastest reactions known).

      • e.g. Methane gas in mines, petrol vapour etc. are all potentially dangerous situations so knowledge of 'explosion/ignition threshold concentrations', ignition temperatures and activation energies are all important knowledge to help design systems of operation to minimise risks.

      • Flammable fine dust powders can be easily ignited e.g. coal dust in mines, flour in mills.

        • Fine powders have a large surface area which greatly increases the reaction rate causing an explosion. Any spark from friction is enough to initiate the reaction!

• A reaction will continue until one of the reactants is used up.
• To measure the ‘speed’ or ‘rate’ of a reaction depends on what the reaction is, and can what is formed be measured as the reaction proceeds? Two examples are outlined below.
• When a gas is formed from a solid reacting with a solution, it can be collected in a gas syringe (see diagram below and the graph).
  • The initial gradient of the graph e.g. in cm³/min (speed or rate) gives an accurate measure of how fast a gaseous product is being formed in  metal/carbonate - acid reaction (forming H₂/CO₂ respectively). You can measure the gas formed every e.g. 30 seconds and plot the graph and measure the initial gradient in e.g. cm³/min or cm³/sec.
  • The most accurate measurements are made early on in the reaction when the gas volume versus time is almost linear. You can take a series of measurements and draw the graph (origin 0,0) to get the rate from the gradient (e.g. cm³/min) or measure the time to make a fixed volume of gas (* see below).
  • If the reaction is allowed to go on, you can measure the final maximum volume of gas and the time at which the reaction stops, though this a very poor measure of rate, because the reaction just goes slower and slower as the reactant amounts/concentrations are decreasing - so don't use this as a method of measuring reaction speed.
  • (*) The reciprocal of the reaction time, 1/time, can also be used as a measure of the speed of a reaction. The time can represent how long it takes to form a fixed amount of gas first few minutes of a metal/carbonate - acid reaction, or the time it takes for so much sulphur to form to obscure the X 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.
  • For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"
• Examples of reactions involving gas formation
  • (i) metals dissolving in acid ==> hydrogen gas, (test is lit splint => pop!),
    • e.g. magnesium + sulphuric acid ==> magnesium sulphate + hydrogen
      • Mg_(s) + H₂SO_4(aq) ==> MgSO_4(aq) + H_2(g)
  • (ii) carbonates dissolving in acids => carbon dioxide gas, (test is limewater => cloudy),
    • calcium carbonate (marble chips)  + hydrochloric acid ==> calcium chloride + water + carbon dioxide
      • CaCO_3(s) + 2HCl_(aq) ==> CaCl_2(aq) + H₂O_(l) + CO_2(g)
  • and (iii) the manganese(IV) oxide catalysed decomposition of hydrogen peroxide (oxygen gas, test is glowing splint => relights)
    • hydrogen peroxide ==> water + oxygen
      • 2H₂O_2(aq) ==> 2H₂O_(l) + O_2(g)
        • can all be followed with the gas syringe method.
  • You can do all sorts of investigations to look at the effects of 
    • (a) the solution concentration,
    • (b) the temperature of the reactants,
    • (c) the size of the solid particles (surface area effect),
    • (d) the effectiveness of a catalyst on hydrogen peroxide decomposition.

• The shape of the graph is quite characteristic (see diagram above and notes below).
  • The reaction is fastest at the start when the reactants are at a maximum (steepest gradient in cm³/min).
  • The gradient becomes progressively less as reactants are used up and the reaction slows down.
  • Finally the graph levels out when one of the reactants is used up and the reaction stops.
  • The amount of product depends on the amount of reactants used.
  • The initial rate of reaction is obtained by measuring the gradient at the start of the reaction. A tangent line is drawn through the first part of the graph, which is usually reasonably linear from the x,y origin 0,0.
    • This gives you an initial rate of reaction in cm³ gas/minute,
    • Typical results from a gas producing reaction are shown below, for different amounts or concentrations of reactants. How to calculate the reaction rate is explained below.
    • e.g. for run q [ ], after 2 mins, 20 cm³ of gas formed, so the rate of reaction is 20/2 = 10 cm³/min.
    • From the graph of results you can measure the relative rate of reaction from (i) the initial gradient in cm³/min (see on diagram above), (ii) you can estimate from the graph the volume of gas formed after a particular time e.g. 3 minutes or (iii) you can estimate the time it takes to form a particular volume of gas. (i) is the best method i.e. the best straight line covering several results at the start of the reaction.
    • 
    • Keeping the temperature constant is really important for a 'fair test' if you are investigating speed of reaction/rate of reaction factors such as concentration of a soluble reactant or the particle size/surface area of a solid reactant. On the advanced gas calculations page, temperature sources of error and their correction are discussed in calculation example Q4b.3, although the calculation is above GCSE level, the ideas on sources of errors are legitimate for GCSE level.
      • Note that if the temperature of a rates experiment was too low compared to all the other experiments, the 'double error' would occur again, but this time the measured gas volume and the calculated speed/rate of reaction would be lower than expected.

• The rate of a reaction that produces a gas can also be measured by following the mass loss as the gas is formed and escapes from the reaction flask.
  • The method is ok for reactions producing carbon dioxide or oxygen,
  • but not very accurate for reactions giving hydrogen (too low a mass loss for accuracy).
  • The reaction rate is expressed as the rate of loss in mass from the flask in e.g. g/min based on the initial gradient (see graph below).

• When sodium thiosulphate reacts with an acid, a yellow precipitate of sulphur is formed and forms the basis of a good project for assessment.
  • To follow this reaction in your investigation you can measure how long it takes for a certain amount of sulphur to form.
  • You do this by observing the reaction down through a conical flask, viewing a black cross on white paper (see diagram below).
  • The X is eventually obscured by the sulphur precipitate and the time noted.
  • sodium thiosulfate + hydrochloric acid ==> sodium chloride + sulfur dioxide + water + sulfur
  • Na₂S₂O_3(aq) + 2HCl_(aq) ==> 2NaCl_(aq) + SO_2(aq) + H₂O_(l) + S_(s)
    • Note: You do not see gas bubbles because the very nasty sulphur dioxide gas is very soluble in water so take care you do not inhale any of the air near the flask when you are doing the experiment or washing out the apparatus afterwards.

mix => ongoing => watch stopped =>

• By using the same flask and paper X  you can obtain a relative measure of the speed of the reaction in forming the same amount of sulphur.
• The speed or rate of reaction can expressed as 'x amount of sulphur'/time, so the rate is proportional to 1/time for a particular run of the experiment.  In other words since you don't know the absolute mass of sulphur formed, the reciprocal of the time is taken as a measure of the relative rate of reaction.
  • You can investigate the effects of
    • (a) the hydrochloric acid or sodium thiosulphate concentration
    • (b) the temperature of the reactants.
  • to show the effects of changing one of the variables you can plot  graphs of e.g.
    • reaction time versus temperature or concentration,
    • or rate of reaction (1/reaction time) versus temperature or concentration.
• You can also measure the speed of this reaction by using a light gate to detect the precipitate formation. The system consists of a light beam emitter and sensor connected to computer and the reaction vessel is placed between the emitter and sensor. The light reading falls as the sulphur precipitate forms.
• Further examples of graphs that may be obtained from the different methods.
• For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

• WHAT IS THE EFFECT OF CHANGING THE CONCENTRATION OF A REACTANT?

• AND WHY IS THE REACTION SPEED CHANGED?

• Why does increase in concentration speed up a reaction?

• 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)

• 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).

• 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 and 4.8 for a numerical-quantitative data interpretation.

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

• WHAT IS THE EFFECT OF CHANGING PRESSURE ON THE SPEED OF A REACTION?

• DOES INCREASING THE PRESSURE ALWAYS HAVE AN EFFECT?

• Why does an increase in pressure speed up a reaction with a gaseous reactant?

• 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 particles are, on average, closer together and collisions between the particles will occur more frequently.

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

• Solid reactants and solutions are NOT affected by change in pressure, there concentration is unchanged.

• CAN STIRRING AFFECT THE RATE OF A REACTION?

• DOES STIRRING AFFECT THE SPEED OF THE REACTION BETWEEN A SOLID AND A SOLUTION?

• Why does stirring speed up a reaction between a solid and a solution?

• 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!!!

• Stirring cannot affect a completely mixed up solution at the particle level i.e. two solutions of soluble substance that react together are unaffected by stirring.

• WHAT HAPPENS TO THE SPEED OF A REACTION IF WE CHANGE THE PARTICLE SIZE OF A REACTING SOLID?

• WHAT DOES BREAKING UP A SOLID REACTANT INTO FINER PIECES DO TO IT IN TERMS OF HOW IT REACTS?

• 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 and 4.8(iii) for a numerical-quantitative data interpretation.

• DOES TEMPERATURE AFFECT THE SPEED OF A CHEMICAL REACTION?

• IF SO, HOW AND WHY?

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

• 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!

• 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.

• 

  • See also graphs 4.3, 4.4 and 4.8 for a numerical-quantitative data interpretation.

  Trying to resolve an apparent confusion for GCSE students!

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.
4. This is usually only fully discussed at AS-A2 level, but it may impress the teacher for GCSE coursework if you look up the Maxwell-Boltzmann distribution of kinetic energies, 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.
5. For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

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

• WHAT IS A CATALYST?

• HOW DOES IT 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!

• 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 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).

• 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
    • 2H₂O_2(aq) ==> 2H₂O_(l) + O_2(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.

• Enzymes are biochemical catalysts are dealt with on another page - 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"

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3g The Effect of Light

• CAN LIGHT AFFECT THE SPEED OF ANY REACTIONS?

• IF IT DOES, HOW DOES CHANGE THE SPEED OF A CHEMICAL REACTION?

• Why does increasing light intensity sometimes increase the speed of a reaction?

• Light energy (uv or visible radiation) can initiate or catalyse particular chemical reactions.

  • As well as acting as an electromagnetic wave, light can be considered as an energy 'bullets' called photons and they have sufficient 'impact energy' to break chemical bonds, that is, enough energy to overcome the activation energy.

  • The greater the intensity of light (visible or ultra-violet) the more reactant molecules are likely to gain the energy react, so the reaction speed increases.

• Examples:

  • Silver salts are converted to silver in the chemistry of photographic exposure of the film.

    • Silver chloride (AgCl), silver bromide (AgBr) and silver iodide (AgI) are all sensitive to light ('photosensitive'), and all three are used in the production of various types of photographic film to detect visible light and beta and gamma radiation from radioactive materials.

    • Each silver halide salt has a different sensitivity to light.

    • When radiation hits the film the silver ions in the salt are reduced by electron gain to silver

      • Ag⁺ + e^- ==> Ag (X = halogen atom, Cl, Br or I)

      • and the halide ion is oxidised to the halogen molecule by electron loss

      • 2X^- ==> X₂ + 2e^-

      • so overall the change via light energy is: 2AgX ==> 2Ag + X₂

    • AgI is the least sensitive and used in X-ray radiography, AgCl is the most sensitive and used in 'fast' film for cameras.

  • Photosynthesis in green plants:

    • The conversion of water + carbon dioxide ==> glucose + oxygen

      • 6H₂O_(l) + 6CO_2(g) ==> C₆H₁₂O_6(aq) + 6O_2(g)

      • requires the input of sunlight energy and the green chlorophyll molecules absorb the photon energy packets of light and initiate the chemical changes summarised above.

  • Photochemical Smog:

    • This is very complex chemistry involving hydrocarbons, carbon monoxide, ozone, nitrogen oxides etc.

    • Many of the reactions to produce harmful chemicals are catalysed or promoted by light energy.

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4. More examples of interpreting graphical results ('graphing'!)

PLOTTING GRAPHS - PLOTS OF GRAPHS OF DATA AND HOW TO INTERPRET THEM

PLEASE Note

(i) rate of reaction = speed, (ii) see two other graphs and notes

(ii) Graphs 4.1, 4.2 and 4.5 just show the theoretical shape of a graph for a single particular experiment. Graphs 4.3 and 4.4 (temperature), 4.6 and 4.7 (concentration) and 4.8 (several factors illustrated) shows the effect of changing a variable on the rate of the reaction and hence the relative change in the curve-shape of the graph line.

(iii) The rate of reaction may be expressed as the reciprocal of the reaction time (1/time) e.g. for the

time for sulphur formation (to obscure the X)  in the sodium thiosulfate - hydrochloric acid reaction

or where a fixed volume of gas is formed, though in this can also be expressed as gas volume/time too as cm³/s or cm³/min even though the gas volume is the same for a given set of results of changing one variable whether it be concentration or temperature.

If you have detailed data e.g. multiple gas volume readings versus time, the best method for rate analysis is the initial rate method described on and below the diagram of the gas syringe gas collection system.

(iv) for detailed observations of gas versus time

Graph 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.

Graph 4.2 shows the increase in the amount of a solid product with time. The graph tends towards a maximum amount possible when all the solid reactant is used up and the graph becomes horizontal. This means the speed has become zero as the reaction has stopped. Here the gradient is a measure of the rate of the reaction.

Graph 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. Theory of temperature effect

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. See the notes on rate in the Graph 4.7 paragraph below and the theory of temperature effect.

Graph 4.5 shows the increase in the amount of a gas formed in a reaction with time. Here the gradient is a measure of the rate of the reaction. Again, the graph becomes horizontal as the reaction stops when one of the reactants is all used up! More on this type of graph.

Graph 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 paragraph above and the theory of concentration effect

Graph 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. Theory of concentration effect

Graph 4.8 A set of results for the same reaction (i) The graph lines W, X, original, Y and Z on the left diagram are typical of when a gaseous product is being collected. The middle graph might represent the original experiment 'recipe' and temperature. Then the experiment repeated with variations e.g. (ii) X could be the same recipe as the original experiment but a catalyst added, forming the same amount of product, but faster. (iii) Initially, the increasing order of rate of reaction represented on the graph by curves Z to W i.e. W > X > original > Y > Z might represent progressively increasing concentrations of reactant or progressively higher temperature of reaction or progressively smaller lumps-particle/increasing surface area of a solid reactant. All three trends in changing a reactant/reaction condition variable produce a progressively faster reaction shown by the increasing gradient in cm3/min which represents the rate/speed of the reaction. (iv) Z could represent taking half the amount of reactants or half a concentration. The reaction is slower and only half as much gas is formed. (v) W might represent taking double the quantity of reactants, forming twice as much gas e.g. same volume of reactant solution but doubling the concentration, so producing twice as much gas, initially at double the speed (gradient twice as steep).

See also graphs for enzymes showing effects of pH, temperature and concentration. For more details on concentration results analysis see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

This page should help with ideas about factors controlling the rates of chemical reactions for coursework projects, assignments, investigations etc.

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• AND on other associated web pages:

  • Enzymes and biotechnology (rates factors, temperature and pH effects, lock-key mechanism)

  • Ammonia synthesis (applying rates and equilibrium factors)

  • Uses of catalysts in ...

    • Transition Metals, Industrial Chemistry and Pollution control

  • Energy changes in chemistry

  • RATES QUIZ QUESTIONS and WORKSHEETS printout or online:

    • GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction

    • MULTI-WORD GAP-FILL QUIZ on Rates of Reaction

    • MATCHING PAIRS QUIZ on rates of reaction Q1 and Q2

    • CROSSWORD PUZZLE on rates of reaction and ANSWERS!

    • 10 JUMBLED SENTENCES to sort out!

  • See also the brainstorm of GCSE rates coursework-projects investigation ideas and two

  • Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS" ADVANCED LEVEL theory pages on 'Kinetics' for GCE-AS-A2-IB/US gr 11-12) and adventurous GCSE students!

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keywords-phrases: 1. What do mean by rate of reaction and its measurement? * 2. Collision theory of reaction * 3. Factors: 3a concentration, 3b pressure, 3c stirring, 3d particle size/surface area, 3e temperature, 3f catalyst, 3g light * 4. Examples of graphs, hydrochloric/sulphuric acid-metal magnesium or carbonate reaction * hydrochloric acid-sodium thiosulphate reaction * Activation energy * Catalysts * Concentration effect * Graphs-gas collection * Graphs-examples * hydrogen peroxide decomposition * How reactions happen * Interpreting results * Light (catalyst) effect * Methods of measuring rate * Pressure effect * Rate of reaction * Reaction profiles *  Stirring effect * Surface area/size of solid particle reactant effect * Temperature effect   GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction MULTI-WORD worksheets revision questions GAP-FILL QUIZ MATCHING PAIRS QUIZ Q1 and Q2 CROSSWORD PUZZLE and ANSWERS! 10 JUMBLED SENTENCES to sort out!

Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  WJEC gcse science chemistry CCEA/CEA gcse science chemistry O Level Chemistry (revise courses equal to US grade 8, grade 9 grade 10) science chemistry courses revision guides

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