INTRODUCTION TO REACTION RATES

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

1. What do we mean by the speed or rate of a chemical reaction?

Why is it important to know how fast reactions occur?

How can we measure the speed or rate of a chemical reaction?

INTRODUCTION to methods of measuring how fast a reaction is going!

Factors affecting the Speed-Rates of Chemical Reactions

How do we define the rate of a reaction? What sort of apparatus to we use to follow the speed of a reaction? How do we go about a rates of reaction laboratory investigation ('labs')? What results can we obtain from rates of reaction experiments? What experimental methods are available for doing rate experiments in schools and colleges? How do we process and interpret the results? These revision notes are suitable for GCSE IGCSE O Level KS4 science chemistry students studying how we can follow the rate of a chemical reaction. The descriptions of experiment apparatus to do with rates of reactions and theoretical explanations should help with homework, coursework assignments, laboratory experiments 'labs' on the factors controlling the speed of a chemical reaction. These notes on slow/fast reaction rate examples, definition of rate of reaction, experimental investigation methods, 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 what to mean by reaction rate?, how do we measure reaction rate?, why is it important to know how fast chemical reactions go?, should prove useful for the new AQA GCSE chemistry, Edexcel GCSE chemistry & OCR GCSE chemistry (Gateway & 21st Century) GCSE (9–1), (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

Introduction to this INTRODUCTION !

Why do we want to know about the speeds (rate) of Chemical Reactions?

• Chemical reactions occur at vastly different rates over periods of a fraction of a second to days or years

• Although the reactivity of chemical reactants is a significant factor in how fast chemical reactions proceed, there are many other variables that can be controlled in order to speed them up or slow them down.

• Chemical reactions may also be reversible and therefore the effect of different variables needs to be understood and controlled in order to maximise the yield of desired product e.g. very important for the efficient economics of a chemical process in industry.

• Understanding energy changes that accompany chemical reactions is another important factor of this process.

• In industry, chemists and chemical engineers determine the effect of different variables on reaction rate to maximise the yield of product.

• Sometimes compromises have to be made in order to optimise the chemical processes in the chemical industry to ensure that enough product is produced within a sufficient time, and in an energy efficient way too.

• 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 or how slow is the reaction’ or 'the speed of the reaction'. It can be measured as the 'rate of formation of product' (e.g. collecting a gaseous product in a syringe) or the 'rate of removal of reactant'. The speeds of reactions vary greatly, with some pretty extreme situations of both slow and fast!

• The chemical weathering of rocks is an extremely very slow reaction and like fossil formation by mineralisation processes, it might take hundreds of thousands or millions of years!

• Rusting is a very ‘slow’ reaction, you hardly see any change looking at it, but after a few weeks or months you notice the difference in a piece of exposed iron or steel! However, its a lot slower with aluminium, so your greenhouse should last a few years!

• Corrosion of stonework, especially limestone buildings, by acid rain is slow, but over the last few hundred years of the industrial revolution many a medieval building has suffered from its effects!

• 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 and hydrogen. These reactions take a few minutes.

• 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 sodium dissolve in water or magnesium dissolves in acid, BUT only because they have formed a water soluble compound e.g. sodium hydroxide or magnesium chloride.

• A much faster would be magnesium burning in air, over in seconds!

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

• Explosive reactions would be described as ‘very fast’, in fact they are some of the fastest reactions known 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.

• Explosions are the fastest chemical reactions you will encounter and are over in a fraction of second, but a bunsen burner flames is a sort of 'controlled explosion' and the chemical reactions of the fuel combustion are very rapid indeed! but the rapidly burned fuel is continuously replaced by the gas flow.

• Explosive materials like dynamite or TNT produce powerful explosions in a fraction of a second.

• More on 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 to make as much product in a given time.

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

• Even just raising the temperature can have a dramatic effect on the speed of a reaction, but raising the temperature costs money.

• Rate of reaction knowledge is essential to operate a chemical plant safely and efficiently so sometimes a compromise must be reached i.e. economic optimum conditions.

• The Haber Synthesis of ammonia is a good example to study in terms of various economic factors.

• Health and Safety Issues:

• The faster a reaction, the more likely there might be health and safety issues to contend with.

• The physical state or concentration of a material can be an important factor.

• Fine powders in flour mills can be easily ignited and can explode as if the flour was a gas.

• Similarly, fine coal dust in a mine poses the same threat.

• Mixtures of flammable gases in air present an explosion hazard (gas reactions like this are amongst the fastest reactions known). If the flammable/explosive gas is in low concentration, there may be no risk, but you need to know the safe limits!

• 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, custard powder production lines!

• 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, if a reactant is in excess, then some of it will be left over.
• To measure the ‘speed of a chemical reaction’ or the ‘rate’ of a reaction depends on what the reaction is, and what is formed that can be measured as the reaction proceeds.
• How do we define the rate or speed of a chemical reaction?
• Mean rate of reaction = amount of product formed / time interval involved
• OR
• Mean rate of reaction = amount of reactant used up / time interval involved
• The use of the word mean is important.
• The rate of a reaction is constantly changing, in fact it is always slowing down as the reactants are being used up.
• Therefore, the best you can do is get an average value at a particular stage in the reaction and preferably as near as possible to the initial rate of the reaction - the speed at the start, after time = zero!
• This sort of definition will become pretty obvious when analysing graphs of date e.g. volume of gas formed at various time intervals, so the rate might be defined as ..
• Mean rate of reaction = cm3 gas formed / time taken for gas to form
• so the rate units might be cm3/s or cm3/min
• If the method of following the rate of reaction involves weigh the reactants or products e.g. ..
• Mean rate of reaction = mass loss / time taken
• so the rate units might be g/s or g/min
• The measured gas volumes (e.g. cm3) or mass losses (e.g. g) can be converted into moles so the rate can be expressed as mol/s or mol/min.
• EXAMPLES of experimental set-ups to follow the rate of a chemical reaction are described below.
• Examples of results data, graphs and their interpretation are also given.
• MEASURING GAS VOLUMES - following the formation of a gaseous product of a reaction by collecting the gas
• When a gas is formed from a solid reacting with a solution, it can be collected in a gas syringe (see the apparatus diagram below and the graph of results).
• Using e.g. a gas syringe system (illustrated below), you measure the volume of gas given off at regular time intervals.
• From the neat table of results, you then plot the results on graph paper (or with a software package like excel) making time the horizontal x-axis and the volume gas of gas the vertical y-axis.
• 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 e.g. in a metal carbonate - acid reaction giving carbon dioxide, or metal - acid reaction giving hydrogen gas.
• It is best to measure the initial rate because the reaction mixture is becoming depleted in reactants so changing the initial concentrations or surface area if there is a solid reactant dissolving.
• You can measure the gas formed every e.g. 30 or 60 seconds (0.5 or 1.0 minutes is typical) 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 gets slower and slower as the reactant amounts/concentrations are decreasing - so don't use this as a method of measuring reaction speed.
• Results - using graph paper
• This graph of volume of gas formed versus time represents a typical set of results, this particular graph is discussed in more details in More examples of graphs and their interpretation.
• The steeper the initial slope/gradient, the faster the initial reaction - its that simple!
• The gradient of the graph become less steep because the reaction will naturally slow down due to the cumulative loss of reactant molecules.
• For a series of experiments the initial gradient is proportional to the relative rate and you measure it by marking and measuring the gradient of the line from 0,0 tangentially over the first few minutes of the reaction when the graph line is usually reasonably linear.
• Three diagrams showing how to measure the initial rate of reaction e.g. gas formation or mass loss.
• When the gradient first becomes zero, i.e. the graph line becomes horizontal, one or more of the reactants has been used up and the reaction has then stopped.
• e.g. in the case of experiment run E, at 4 mins, at a maximum volume of 20 cm3 of gas formed.
• In terms of initial rate, the 'speed order' is W > X > E > Y > Z, with decreasing initial graph gradient.
• Graphs X, E, Y and Z can be the result of decreasing concentration from X to Z.
• The same pattern might be seen by decreasing the temperature from X to Z.
• Graph W might be the result of doubling the amount or concentration of one of the reactants compared to E, steeper initial gradient (faster) and twice the amount of gaseous product.
• Graphs X and E could represent the effect of adding a catalyst to X.
• Graphs Z to X could be produced by progressively crushing the same amount of solid reactant e.g. lumps < granules < powder etc.
• Note: Complying with 'a fair test' experiment design is discussed in detail for the experiments in each of the sections for the effect of 3a. concentration, 3c surface area, 3d temperature and 3e. catalyst.
• Apart from the initial gradient, you can use as a measure of the rate of reaction, how much gas is formed for a constant time e.g. product gas volume after 2 minutes.
• (*) 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 after the 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 reaction 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.
• however, I think the initial gradient method to calculate the initial rate is more accurate because after a 'fixed time' the reaction is slowing down with the natural depletion of reactants in forming the products.
• For more details for A Level students see ...
• Examples of reactions involving gas formation - DIAGRAM of APPARATUS and graph shown below
• You can follow the speed/rate of a reaction that produces a gaseous product by collecting the gas and measuring the total volume at suitable time intervals.
• (i) metals dissolving in acid ==> hydrogen gas, (test is lit splint => pop!),
• e.g. magnesium + sulphuric acid ==> magnesium sulphate + hydrogen
• Mg(s) + H2SO4(aq) ==> MgSO4(aq) +  H2(g)
• You can measure the volume of hydrogen gas formed with a gas syringe.
• You can do a similar experiment with hydrochloric acid.
• Mg(s) + 2HCl(aq) ==> MgCl2(aq) +  H2(g)
• (ii) carbonates dissolving in acids => carbon dioxide gas, (test is limewater => cloudy),
• calcium carbonate (marble chips)  + hydrochloric acid ==> calcium chloride + water + carbon dioxide
• CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l)  +  CO2(g)
• You can measure the volume of carbon dioxide gas formed with a gas syringe.
• (iii) the manganese(IV) oxide catalysed decomposition of hydrogen peroxide (oxygen gas, test is glowing splint => relights)
• hydrogen peroxide ==> water + oxygen
• 2H2O2(aq) ==> 2H2O(l)   +  O2(g)
• You can measure the volume of oxygen gas formed with a gas syringe.
• All of these gas forming reactions can all be followed with the gas syringe method suitable for use in schools and colleges.
• You can do all sorts of investigations to look at the effects of varying ...
• (a) the reactant solution concentration, and in industry pressure is varied, which is effectively varying the concentration of gaseous reactants.
• (b) the temperature of the reactants (solids plus solutions etc.),
• (c) the size of the solid reactant particles (surface area effect),
• You could also vary the particle size of a solid catalyst to vary its surface area.
• (d) the effectiveness of a catalyst on hydrogen peroxide decomposition.
• all of which are controlling factors on the speed or rate of a chemical reaction.
• BUT don't forget the forgotten factor if there is a solid reactant like limestone or a solid catalyst - stirring.
• The mixture should be gently and steadily stirred, preferably with a magnetic stirrer system.
• If you don't, the bottom layers of the solution become depleted more than the upper layers as the solid reactant will sediment out.
• This leads to erratic and inaccurate results.

• The shape of the graph is quite characteristic of the progress of a chemical reaction.
• See the 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
• (iii) you can estimate the time it takes to form a particular volume of gas.
• (i) 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.
• Results using Excel
• You can use application software like Excel to process and portray your data in a graphical format, or you can just plot your data, but accurately, on ordinary graph paper.
• 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.
• Extra note on graph interpretation
• How to derive other rate data from the same graphs as above.
• Using graph W: From the two blue crosses marked at 1 and 3 minutes respectively you can calculate the average speed/rate of the reaction W between the 1st and 3rd minutes of the reaction.
• Average rate = (volume at 3 mins - volume at 1 min) / time elapsed
• Average rate = (37 - 22) / 2 = 9/2 = 4.5 cm3 gas/min
• Using graph E: From the tangent drawn at 2 minutes you can get the rate of reaction E at 2 minutes.
• Without thinking, its pure coincidence I've used the 1st and 3rd minute again!
• From the tangent line: volume at 1 min = 12 cm3, volume at 3 min = 20 cm3
• Rate at 2 minutes = (volume at 3 mins - volume at 1 min) / time elapsed
• Rate at 2 minutes = (20 - 12) / 2 = 8/2 = 4.0 cm3 gas/min
• GAS MASS LOSS - following the mass loss due to the escape of a gaseous product
• 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 (limestone + acid), but I don't think it is as good as the gas syringe method for several reasons e.g.
• I don't consider this as accurate as the gas syringe method because the mass of gas is small, particularly in the case of hydrogen (molecular mass 2), it might be just ok for carbon dioxide (molecular mass 44) with a very accurate balance.
• Another problem is stirring, when dealing with a solid reactant and solution, the mixture, ideally, should be gently stirred to avoid local depletion of reactants, so how can you possible do it on a sensitive balance!
• 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).
• The total mass of the flask and contents are recorded at selected suitable time intervals and subtracted from the initial mass, therefore you have then recorded the actual mass loss which equals the mass of product formed.

• All results should be neatly tabulated and appropriate graphs drawn (see diagram below).

• The initial gradient of mass loss gives the rate of the reaction (you can ignore the negative sign, since technically its a negative gradient).

• The speed of the reaction would be measured as g/min or g/s, mass loss of reactant per unit of time, or mass of product formed per unit of time.

• This graphs represents the sort of results you might get.

• The faster the reaction, the greater the mass loss in a given time and the negative gradient of the graph is steeper.

• You can investigate different concentrations of soluble reactants or different sized pieces of a solid reactant or catalysts.

• The steeper the initial gradient, the faster the reaction.

• In terms of rate of reaction, X > E > Y > Z, which you see for a series of increasing temperatures or a series of increasing reactant concentration.

• PRECIPITATION - following the amount of a precipitate formed
• 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 precipitate 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 reaction can be followed in this way because both solutions, and after mixing, are colourless and more importantly, they are clear and transparent.
• Its the cloudiness that causes the cross to 'disappear'!
• The X is eventually obscured by the sulphur precipitate and the time noted.
• All results should be neatly tabulated. example?

• sodium thiosulfate + hydrochloric acid ==> sodium chloride + sulfur dioxide + water + sulfur
• Na2S2O3(aq) + 2HCl(aq) ==> 2NaCl(aq) + SO2(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 reactant solution mixture.
• 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.
• This would be a superior method than the naked eye because direct observation is subjective, different students might have different perceptions as to when the cross disappears!
• It is important you take the same total volume of solution to give the same depth of liquid you are viewing the cross through.
• Note that unlike the other methods described on this page, you do not get lots of data to plot to get the rate for a particular experiment, its a sort of 'one of' measurement for a particular concentration or temperature.
• So you get one graph when investigating one factor! e.g. for the effect of temperature the results might look like this ...
•  temperature (oC) 20 25 30 35 40 45 time for X to be obscured (s) 240 220 190 150 100 40
• ... from which you get one graph like or .
• Further examples of graphs that may be obtained from the different methods.
• For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"

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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