Pre-test Results:
Although all of these concentrations are shown to have worked, I do not have enough time available to conduct results including each concentration. If I am going to repeat each concentration several times for more accurate results I need to investigate a smaller range of concentrations. I have chosen to exclude the two lowest concentrations (0.001 and 0.002mol/dm-3) when conducting my actual experiment. This is because they took the longest time for the mixture to turn colourless. Whilst conducting this pre-test I used a magnetic stirrer. However, I was not sure whether this was the most appropriate method to use. To solve this I conducted another pre-test, which is shown in the table below:
From my results a reliable trend can be seen fro the results obtained when stirring by hand and when using the magnetic stirrer. However, when the mixture is left without stirring anomalies arise, as is shown in the table. There is a reliable trend until the concentration of 0.005mol/dm-3, where the time taken for the mixture to turn colourless decreases rather than increasing. To ensure that I could see the colour chance easily, I placed the conical flask on a white tile. This meant that the colour change would be more obvious. I have decided to use the magnetic stirrer as it allows me to concentrate more on when the colour change occurs, rather than stirring the mixture. When I was stirring the mixture by hand, I found that I had a tendency to lose concentration and lose track of when the colour change occurred. Having a magnetic stirrer meant that all I needed to do was focus on the mixture and note the timing of the colour change. I believe this will help me in obtaining accurate results.
I also conducted a pre-test to trial my method for investigating how temperature change affects the reaction rate. I came across a few problems that need to be adjusted in my final experiment. I originally decided to just monitor the temperature of the water bath, but soon realised that this would most likely not share the same value as the solutions in the boiling tubes. To overcome this problem I placed a thermometer in each boiling tube, one in the tube containing potassium bromide and one in the solution of the other reactants. Once both had reached the correct temperature I was able to mix the two solutions, knowing that both were at the required temperature. I also discovered that when investigating temperatures above 40°C the reaction became too fast to record accurate timings. I would have lowered the concentration of potassium bromide, although I wanted to keep this constant to my previous experiments investigation how concentration affects rate of reaction. Instead, I chose the following range of temperatures: 10, 15, 20, 25, 30, 35 and 40°C. This range also makes it easier for me to keep the temperature of the reactants stable, as they do not fluctuate to far above or below room temperature.
For the concentrations of the other reactants used in this reaction I decided to alter them by the same amount and in the same range as I have chosen to alter the concentration of potassium bromide. This ensures that the experiment is as fair as possible, as all reactants are being varied by the same amount. It could be harder to work out a rate equation if I vary the concentrations of the reactants by different amounts to each other.
Method for Investigating Concentration Change
Apparatus:
Method
- Take 5 burettes and label them as “potassium bromate(V)”, “sulphuric acid”, “phenol”, “potassium bromide” and “distilled water”. Wash them out with water and then with the corresponding solution. Fill each burette with its corresponding solution.
-
Add 5cm3 of potassium bromate(V) solution, 5cm3 of sulphuric acid solution and 5cm3 of phenol solution to a beaker. Add 4 drops of methyl orange indicator to the beaker. These amounts should be kept constant throughout the experiment to ensure that the test is as fair as possible. The same solution should be used in each trial to keep the test fair, as different solutions may have very slight differences in concentration.
-
Add 5cm3 of potassium bromide solution of the required concentration to another beaker. For most trials the potassium bromide will need mixing with distilled water to reach the desired concentration.
- Place the first beaker onto a white tile and pour into it the potassium bromide solution. Start the stopwatch and continuously stir the mixture until it goes completely colourless. Record the time taken for this to happen.
-
Wash the beakers out with distilled water and repeat the experiment using concentrations of 0.003, 0.004, 0.005, 0.006, 0.008 and 0.01mol/dm-3 of potassium bromide. Repeat each concentration 5 times. All other factors such as temperature should be kept constant and the same apparatus should be used each time. Sulphuric acid concentration should be kept at 0.1mol/dm-3 when not being investigated, so that it is not so weak that the reaction takes place too slowly to be measured. All other solutions should be kept at 0.01mol/dm-3.
-
Once all concentrations of potassium bromide solution has been investigated and repeated, conduct experiments with respect to potassium bromate(V) and sulphuric acid. To do this, simply alter the concentration of the chosen reactant as was done when using potassium bromide. Concentration can be changed simply by taking a sample of the original solution and adding distilled water. Where M1 is the molarity of the solution before dilution, M2 is molarity after dilution, V1 is the volume of the solution before dilution and V2 is volume after dilution:
V1 = (M2 x V2)
M1
For example, to make a 250ml solution with concentration 0.008M of potassium bromide from a solution of 0.01M, you would do the following:
V1 = (0.008 x 250ml)
0.01
= 200ml
This means that you would need to use 200ml of the original 0.01M solution with 50ml of distilled water to make 250ml of 0.008M solution.
Remember to keep all other solutions at a constant concentration throughout. Investigate the same range of concentrations as was used for potassium bromide, to ensure that a rate equation can easily be found.
Making Required Solutions
0.01M potassium bromate(V):
KBrO3:
K = 1 x 39.1 = 39.1
Br = 1 x 79.9 = 79.9
O3 = 3 x 16.0 = 48.0
RFM = 167
grams needed = (0.01 x 167 x 250)
1000
= 0.4175g dissolved in enough distilled water to make 250ml of solution.
1M Sulphuric acid:
H2SO4:
H2 = 2 x 1.0 = 2
S = 1 x 32.1 = 32.1
O4 = 4 x 16.0 = 64.0
= 98.1
grams needed = (1 x 98.1 x 250)
1000
= 24.525g dissolved in enough distilled water to make 250ml of solution.
0.0001M phenol:
C6H5OH:
C = 6 x 12.0 = 72.0
H = 6 x 1.0 = 6.0
O = 1 x 16.0 = 16.0
= 94.0
grams needed = (0.0001 x 94 x 250)
1000
= 0.00235g dissolved in enough distilled water to make 250ml of solution.
0.01M potassium bromide:
KBr:
K = 1 x 39.1 = 39.1
Br = 1 x 79.9 = 79.9
= 119.0
grams needed = (0.01 x 119 x 250)
1000
= 0.2975g dissolved in enough distilled water to make 250ml of solution.
Risk Assessment ()
The following risk assessments are based on concentrated solutions of each substance. The concentrations I will be using are much more diluted that the solutions mentioned below to ensure that my experiment is safer. Although some of the acute hazards and symptoms associated with the concentrations I am using will not be as severe as those shown below, the risks are still very serious. For this reason I will still take appropriate precautions and will handle each substance with care.
Potassium bromate(V):
Sulphuric acid:
Phenol:
Potassium bromide:
The hazards for this substance are almost negligible, especially for the low concentrations that I will be using. However, protective goggles will still be worn to avoid contact with eyes, as this could cause irritation. This substance could also cause some discomfort if ingested, so I will take be very careful to ensure that this doesn’t happen.
Method for Investigating Temperature Change
Method
- Set the apparatus up as was done in the experiment investigating concentration change. However, this time use a heating plate rather than a magnetic stirrer so that temperature can be adjusted.
-
Mix 5cm3 of potassium bromate(V) solution, 5cm3 of sulphuric acid solution and 5cm3 of phenol solution in a boiling tube and add 4 drops of methyl orange indicator.
-
Add 5cm3 of potassium bromide to a separate boiling tube.
- Fill a large beaker with water and place it on top of the heating plate and set it to the required temperature. To achieve cooler temperatures put ice cubes into the beaker and monitor the temperature using a thermometer. Place the boiling tubes containing the correct solutions in the water bath so that they are also heated or cooled to this temperature.
- Mix the two solutions and start the stopwatch. Record the time taken for the solution to go colourless. Ensure that the temperature of the water is kept constant. This can be done by having a thermometer in the beaker. Although the higher temperatures should be easily maintained by the heating apparatus, lower temperatures must be kept constant by adding more ice cubes if required.
- Repeat 5 times with each temperature to ensure accurate and fair results.
Results With Respect to Potassium Bromate(V)
The table below shows my results when varying the concentration of potassium bromate(V):
Graph 1 on the next page shows the concentration of potassium bromate(V) plotted against the average time taken for the solution to turn colourless. From this graph I can see that there is a negative correlation because as concentration is increased, the time taken for the solution to turn colourless decreases. However, this graph does not provide enough information to work out the rate equation data for potassium bromate(V). For this, I had to draw up a graph of concentration against reaction rate. Reaction rate is worked out using the following equation:
Reaction Rate = 1
Time
This is presented in Graph 2. Graph 2 has a line of best fit which is a straight diagonal line, rather than a curve. It also shows that there is a positive correlation between concentration and reaction rate, as when concentration of potassium bromate(V) is increased, reaction rate increases. I could draw the line of best through the origin at point (0,0) as I know that if the concentration is 0M then the reaction will not occur – the reaction rate will be 0seconds-1 (this can be applied to all other graphs of concentration against reaction rate that I have drawn). I decided to draw the majority of my graphs using a computer to plot the axes and points, and drawing the line or curve of best fit by hand. The computer can provide better accuracy than I can as far as drawing the axes and plotting the points is concerned. However, I believe that I can provide more accuracy when drawing the line or curve of best fit.My line of best fit suggests that the reaction, with respect to potassium bromate(V), is first order. This means that that rate equation so far is as follows:
Rate = k[BrO3]
Results With Respect to Potassium Bromide
The table below shows my results when varying the concentration of potassium bromide:
Graph 3 on the next page shows the concentration of potassium bromide plotted against reaction rate. The line of best fit is once again a straight diagonal one. This shows that there is again a positive correlation between concentration and reaction rate, as when the concentration of potassium bromide is increased, the reaction rate increases. This means that the reaction, with respect to potassium bromide, is also first order. This in turn means that the rate equation so far looks like this:
Rate = k[BrO3][Br -]
Results With Respect to Sulphuric Acid
The table below shows my results when varying the concentration of sulphuric acid:
As can be seen from the table above, at concentrations of 0.3 and 0.4M, the solution took much too long to turn colourless. This was impractical, as it meant that I was less likely to be able to conduct a suitable number of repeats to get an accurate average time. Because of this I was forced to make some minor modifications to my original method. As I only had a range of four concentrations recorded for sulphuric acid, I knew that this would not be enough to draw a very useful graph of concentration against reaction rate. For this reason, I decided to conduct experiments testing two more concentrations within the range of the other four, thus making up for the two concentrations I had not been able to investigate. The results for these extra concentrations are shown below:
The result in my table highlighted both in bold and by an asterisk is clearly an anomaly. For this reason I decided to leave it out when calculating the reaction rate at 0.6M. Although this meant that I only used four repeats for this concentration, the average results I obtained still coincides with other results and is therefore still relatively accurate.
Graph 4 on the following page shows the concentration of sulphuric acid plotted against the reaction rate. From my line of best fit it is clear that there is a positive correlation between the two. My line of best fit is in a curved shape. Relating back to my theory section, this shows that the reaction, with respect to sulphuric acid, is second order. This in turn means that I have proved the final rate equation to be:
Rate = k[BrO3][Br -][H+]2
This is exactly what I predicted at the beginning of my investigation.
Now that I have proved the rate equation experimentally, I want to find a value for the constant, ‘k’. This is called the rate constant (1, page 229) and its value changes with temperature. Using my results at a certain temperature I should be able to calculate a value for ‘k’. I will use the following values:
[BrO3] = 0.008
[Br -] = 0.01
[H+]2 = (0.1)2 = 0.01
At room temperature, I calculated the reaction rate for these values to be 0.0221seconds-1. I will now substitute the above values into the rate equation in order to calculate the constant ‘k’, although I first need to rearrange the equation to make ‘k’ the subject:
Reaction Rate = k[BrO3][Br -][H+]2
k = Reaction Rate
[BrO3][Br -][H+]2
k = 0.0221
(0.008)(0.01)(0.01)
k = 0.0221
(8 x 10-7)
k = 27625
The units for ‘k’ vary depending on the reaction investigated. In this case they are worked out as follows:
k = (moldm-3)(seconds-1)
(moldm-3)(moldm-3)(moldm-3)2
k = 27625s-1mol-3dm9
As I mentioned earlier, a change in temperature alters the value for ‘k’. The value I calculated above was taken at room temperature. On the following page is a table showing my results when I investigated how temperature affects the reaction rate.
Results When Varying Temperature
From Graph 5 on the next page, it is clear that there is a negative correlation between temperature and time taken for the solution to turn colourless – as the temperature of the reactants is increased, the time taken for the solution to turn colourless decreases. Graph 6 shows that there is a positive correlation between temperature and reaction rate – as the temperature of the reactants is increased, the reaction rate increases. Both graphs have a curve of best fit, rather than a straight diagonal line.
Using the rearranged Arrhenius equation discussed in my theory, I can draw up a table of values needed to work out the activation enthalpy for this reaction:
These results are plotted in Graph 7 on separate graph paper. This graph shows 1/temperature in Kelvin against log to the base e of reaction rate. I chose to use a completely hand-drawn set of axes so that I could accurately work out the gradient of the line, thus enabling me to work out an accurate value for the activation enthalpy. I could not easily draw the triangle on the line of best fit, which allows me to calculate the gradient, using a computer. Graph 7 shows that there is a negative correlation between the two sets of data – as 1/temperature is increased, log to the base e of reaction rate decreases.
To find the gradient of a straight line, you use the following equation: ()
Gradient = Change in y value
Change in x value
For my graph, I worked out the gradient as follows:
Gradient = -1.75
0.00024
= -7291.7
Gradient = -Ea = -7291.7
R
R (the gas constant) = 8.31 J K-1 mol-1
The gradient equation can be rearranged as:
-Ea = gradient x R
= -7291.7 x 8.31
= -60594
Ea = 60594Jmoles-1
This can be converted to kJmoles-1 by dividing by 1000 (the number of joules in a kilojoule):
60594
1000 = 60.594kJmoles-1
This is the activation enthalpy for my reaction.
Conclusion
I have already drawn out all of my results tables and have drawn graphs for each one. From these I have been able to work out the rate equation, rate constant and activation enthalpy for this reaction. Therefore I have fulfilled all of the aims that I set out to complete.
I collected four different sets of results throughout my experiment. The first three sets of results concerned the concentration of different substances in the reaction. I varied these to try and find the rate equation. In general, I found that the solution turned colourless at a much slower speed at lower concentrations. At the same time this means that the reaction occurred quicker at higher concentrations. This can be simply explained by referring back to the collision theory that I mentioned earlier in my investigation. This suggests that when there is less reactant molecules it is less likely that there will be a collision between two particles. If molecules are not colliding as often then reactions between molecules will also occur less often. This means that the reaction occurs at a much slower rate. Graph 1 shows this relationship in my investigation - as the concentration of potassium bromate(V) is decreased, the amount of time taken for the solution to turn colourless, and thus the reaction to finish, increases. This is a strong negative correlation.
The reaction rate of a reaction can be worked out using the following equation:
Reaction Rate = 1
Reaction Time
I calculated the reaction rate for all of my sets of results and then plotted graphs showing concentration of solution against reaction rate. I could then determine the order of reaction for each substance used in the reaction. I already knew from my research that the reaction should be first order, with respect to potassium bromate(V) and potassium bromide, and second order with respect to sulphuric acid. I was able to clearly prove the correct order of reaction for each of these chemicals by using my graphs of concentration against reaction rate. To work out the overall order of the reaction all of the individual orders of reaction must be added together. As I have correctly calculated the individual orders of reaction, I can accurately state that the overall order of reaction is as follows:
Overall order = 1 (KBrO3) + 1 (KBr -) + 2 (H2SO4)
= 4
Every reaction has a reaction mechanism. This looks at the individual ‘steps’ in the reaction which may not obviously be occurring. Using the rate equation I can work out how many of each ion is involved in the rate determining step. The rate determining step is described as the slowest step in the reaction mechanism. To find how many ions are used in this step, you simply look at the power to which each ion is raised. For example, the reaction, with respect to H+ ions, is second order. This means that there are two H+ ions involved in the rate determining step. The general equation for this reaction is as follows:
BrO3-(aq) + 5Br -(aq) + 6H+(aq) 3Br2(aq) + 3H2O(l)
The rate equation, as previously proved, is:
Rate = k[BrO3][Br -][H+]2
I can now see that the rate determining step involves one BrO3 - ion, one Br - ion and two H+ ions. It is likely that this slow rate determining step is preceded by faster reactions. The mechanism which I suggest occurs is as follows:
Step 1: (fast)
H+ + Br - HBr
Step 2: (fast)
H+ + BrO3 - HBrO3
The HBr and HBrO3 formed in the previous two reactions now react in Step 3.
Step 3: (rate determining step)
HBr + HBrO3 HBrO + HBrO2
The products of Step 3 then react further to give water and bromine in Step 4 and 5.
Step 4: (fast)
HBrO2 + HBr 2HBrO
Step 5: (fast)
HBrO + HBr H2O + Br2
Reactions mechanisms cannot be proven – they are just suggestions to what happens in reactions, based on a good understanding of chemistry. The reaction mechanism above is simply a suggestion as to what may happen in this reaction.
Concentration was not the only thing that I varied in my investigation. I also chose to vary the temperature of the reactants. I already knew from my research that changing temperature has an affect on rate of reaction, as it changes the value of the rate constant, ‘k’. My results and graphs show clear correlation that if temperature is increased, reaction rate increases. This means that the time taken for the solution to turn colourless decreases as temperature is increased. I also found that even a small increase in temperature could dramatically affect the reaction rate. For example, when I changed the temperature of the reactants from 10 to 20°C, the average time taken for the reaction to occur fell from 102.98 to 42.76 seconds. This means that, with a temperature change of just 10°C, the time take for the solution to turn colourless fell by 60.22 seconds.
The reason for this dramatic change in reaction rate can again be explained by the collision theory and kinetics. As temperature increases, molecules gain more kinetic energy and can therefore move quicker. The fact that they are moving more quickly means that molecules are statistically more likely to collide. This causes an increase in reaction rate. Molecules will not only collide more frequently with higher temperatures, but they will collide more often with the required energy for a reaction to occur. This required energy is called the activation enthalpy. As temperature is raised, more molecules reach the activation enthalpy and therefore the reaction is more likely to take place sooner.
I plotted a graph of log to the base e of reaction rate against the reciprocal of the temperature in Kelvin (1/temperature). This gave me a gradient of –Ea/R as is mentioned in the section earlier in my investigation about the Arrhenius equation. Using this graph I could then work out the activation enthalpy. I found it to be 60.594kJmoles-1. The Maxwell-Boltzmann diagram shown in my theory section explains this.
Earlier in my investigation I predicted that with an decrease of just 10°C, between 20 and 10°C, the reaction time would at least double. The percentage increase for my results can be worked out as follows:
Reaction time at 10°C = 102.98 seconds
Reaction time at 20°C = 42.76 seconds
42.76 x ? = 102.98
? = 102.98
42.76
? = 2.41
This means that, in my experiment, a decrease of 10°C brought about a 241% increase in reaction time. This shows that the effects of temperature are very important on reaction rate.
At room temperature, I calculated the value for the rate constant, ‘k’. I worked out a value of 27625s-1mol-3dm9. Although the value for ‘k’ does not change with concentration, it is affected by a change in temperature. This is why the value for rate of reaction changes as temperature is changed.
In my reaction an orange solution turns colourless after time. This is because the bromine that is being formed, as shown in my suggested reaction mechanism, reacts with phenol. Once the bromine has filled all of the binding sites on the phenol present in the solution, is reacts with the methyl orange indicator. This causes the orange colour to turn colourless.
When I was varying temperature, I obtained an anomalous result. This result clearly did not fit in with the trend that was developing with the rest of the results. I chose to ignore this result as it would obscure my average reaction time and therefore my calculations for rate of reaction would also be affected. I will elaborate on this and other problems in my evaluation section.
Evaluation
In general I think that my investigation went well. I succeeded in all of my initial aims and produced accurate, reliable results. I only encountered one real anomaly throughout the investigation process (highlighted in my conclusion).
I conducted a pre-test in my planning section of my investigation. This allowed me to highlight any problems with my initial method and I therefore had a chance to rectify them, rather than encountering problems later on. I investigated suitable concentrations of each solution to use. I noticed that at much lower concentrations, the reaction went much too slowly. If the reaction goes too slowly I would not have time to do five repeats for each concentration. This would reduce both accuracy and reliability of my results. For this reason I decided to exclude the two lowest concentrations of each solution. Although this meant that I had lost some of the range of my results, I still investigated six different concentrations for each substance. This means that I was still able to obtain accurate, reliable results.
Another thing I wanted to find out in my pre-test was how to stir the solution whilst the reaction was taking place. I conducted an experiment to investigate this. I found that the most useful method of stirring would be to use a magnetic stirring device. This ensured that I only had to concentrate on one thing – the point at which the solution turned completely colourless. If I had to stir the solution by hand I may not notice the colour change straight away. Using a magnetic stirrer I could ensure that my results were more accurate. I also put my conical flask, with reacting solution in it, on a white tile. This meant that any colour change was more obvious, as it was against a white background. The surface of the table I was working at was a darker colour, meaning that if the conical flask was placed straight onto the table it would be much harder to see when the solution had turned colourless.
I also conducted a pre-test to trial my method for investigating how temperature change affects the reaction rate. I had originally decided to monitor only the temperature of the water bath heating my solutions in boiling tubes. However, I realised that this would most likely not share the same value as the temperature of the solutions in the boiling tubes. To overcome this problem I placed a thermometer in each boiling tube – one in the tube containing potassium bromide and one in the solution of the other reactants. Once both had reached the correct temperature I was able to mix the two solutions, knowing that both were at the required temperature. I also discovered that when investigating temperatures above 40°C the reaction became too fast to record accurate timings. I would have lowered the concentration of potassium bromide, although I wanted to keep this constant to my previous experiments. This would mean that I could calculate a more accurate rate equation. Instead, I chose the following range of temperatures: 10, 15, 20, 25, 30, 35 and 40°C. This range made it easier for me to keep the temperature of the reactants stable, as they do not fluctuate to far above or below room temperature.
I took into account how precise and reliable all of my separate measurements throughout the investigation were. This is why I decided to use burettes instead of measuring cylinder. As explained in my pre-test section, the percentage error associated with measuring 5cm3 of solution in a burette is just 2%. This is much more accurate than a measuring cylinder. The percentage error for a measuring cylinder is as follows:
Percentage error = error
measured value x 100
So for measuring, for example, 5cm3 from a measuring cylinder (with error +/- 0.5cm3) the error would be:
Percentage error = 0.5
5 x 100
= 10%
This is a much larger percentage error than for a burette, so I made the right decision to ensure accuracy by using a burette rather than a measuring cylinder.
The volumes of solution I used were as follows:
Potassium bromate(V) = 5cm3
Potassium bromide = 5cm3
Phenol = 5cm3
Sulphuric acid = 5cm3
The error for all of these measurements is just 2%. This small percentage error is almost negligible in terms of how much it will have affected my results, due to the fact that the error is so small. I could also use all of these volumes when varying temperature. Although I was now using boiling tubes rather than a conical flask to house the reactions, the fact that all volumes were just 5cm3 meant that none of the volumes had to be changed in order to fit all the solutions into a single boiling tube.
In my experiment I chose to study rate of reaction as a ‘clock reaction’. This means that rate is measured by timing how long it takes to produce a certain amount of one product. In my experiment this product was bromine. If the rate of reaction is high then the reaction time must be lower, as rate is the inverse of reaction time. This is shown as:
Rate = 1
Reaction Time
This method works in my reaction due to the presence of phenol. At first any bromine formed reacts with phenol. However, when all of the binding sites on all of the phenol molecules present are filled, bromine begins to react with the methyl orange indicator. This in turn causes the solution to turn colourless by decolourising the methyl orange indicator. Because I kept the concentration and amount of phenol, and the amount of methyl orange added, the same throughout my investigation I know that the amount of bromine needed in each test is the same. This ensures that the test is fair. This is an accurate method as it is simple and therefore it is easy to avoid mistakes. I did not need to take a series of measurements throughout the reaction, which increases the chances of human error. By plotting a graph of reaction rate against concentration I could easily find the order of reaction. I also did not need to draw tangents to any of my curves, which further reduces any chance of human error when drawing and analysing my data in graph form.
I tested four variables in my investigation. The first variable I tested was the concentration of potassium bromate(V). I drew up results tables and drew a graph of reaction rate against concentration of potassium bromate(V). I found that the reaction was first order with respect to potassium bromate(V). This was exactly as I had predicted in my aim. I did not encounter any anomalous results whilst investigating this variable and so can be sure that these results are accurate. I carried out five repeats for each concentration for each variable, ensuring that I had accurate and reliable results.
Next I investigated how changing the concentration of potassium bromide affected the rate of reaction. Again I drew up tables and a graph showing reaction rate against concentration of potassium bromide. I found that the reaction was also first order with respect to potassium bromide. This again was as I had predicted earlier on in my investigation. There were no anomalous results for potassium bromide, so I am sure that my results are accurate.
I conducted an experiment to test how the concentration of sulphuric acid affects the rate of reaction. I drew up a results table and drew a graph of reaction rate against concentration of sulphuric acid. I found that the reaction was second order with respect to sulphuric acid. This was what I predicted in my aim. However, I did have an anomalous result whilst recording my results for sulphuric acid. This came at a concentration of 0.06M and the value was 101.7seconds. When compared with the other results for the same concentration (65.4, 64.9, 64.8 and 65.0seconds) this is clearly an anomaly. Although this could have potentially affected my values for rate of reaction, I came up with a solution. Because I had conducted five repeats for each concentration, I knew that the exclusion of one of these repeats would not have a huge affect on my results. For this reason I was able to exclude the value of 101.7seconds when calculating an average reaction time for the concentration of 0.06M. This also meant that it was not included when further investigating the order of reaction for sulphuric acid. The fact that I still found that the order of reaction, respect to sulphuric acid, was two means that my results were still accurate even when excluding this anomaly.
The aforementioned anomaly could be caused by several different things. My own human error may have meant that I produced a solution of sulphuric acid with a lower concentration than I had intended to. This would make the reaction go much slower than in the other repeats. One other possibility is that the solution had somehow become contaminated. If another chemical was included in the sulphuric acid solution, it would change how quickly the reaction occurred. This could be due to not washing out the apparatus thoroughly enough. Another possible reason is simply that I did not notice the colour change at the right time. This is human error but I don’t think that this was very likely, as the colour change is very easy to see. I think that the most likely reason was making up the solution incorrectly, as this is most likely to affect the reaction time as greatly as the anomaly suggests. For this reason I disposed of the sulphuric acid solution and made up a fresh batch. This appears to have worked, as in my results table for sulphuric acid the values were much closer together and continued the trend, after the first repeat had been ignored. This human error can easily be avoided by ensuring that each solution is measured out very carefully.
Another potential problem that I could have encountered would have been if I had added too much phenol, or had added phenol of a stronger concentration. This would mean that there were more phenol molecules. If there are more phenol molecules then the bromine made in the reaction may not exceed the amount of phenol. This would in turn mean that all of the phenol binding sites are not filled, and bromine does not get a chance to react with the methyl orange indicator. This would mean that the solution would not turn colourless.
Overall, I think that my investigation was a success. I managed to prove that the rate equation for the bromine clock reaction was as follows:
Rate = k[BrO3-][Br -][H+]2
This is exactly as I had predicted in my aim.
I investigated the affect that temperature has on the rate of reaction, and found that the rate constant, ‘k’, changes with temperature. I then used this knowledge, along with the Arrhenius equation to find the activation enthalpy of this reaction. I found that the activation enthalpy in my investigation was 60.594kJmoles-1. This seems to be a reasonable value, and I can therefore conclude that I achieved all of my aims.
References
Salters Advanced Chemistry - Chemical Ideas (Heinemann) - page numbers included where referenced.
http://www.science.uwaterloo.ca/~cchieh/cact/c123/coneffec.html
http://www.chemguide.co.uk/physical/basicrates/temperature.html
http://www.chemguide.co.uk/physical/basicrates/arrhenius.html
http://www.hometrainingtools.com/articles/making-chemical-solutions-science-teaching-tip.html
http://www.mathsteacher.com.au/year10/ch03_linear_graphs/02_gradient/line.htm