Investigation- To investigate the factors involved in the following reaction:Na2S2O3(aq) + 2HCl(aq) 2NaCl(aq) + H2O(l) + SO2(aq) + S(s)

How I will make my investigation a fair test:

To make my investigation a fair test I must only change one variable at a time and keep everything else constant. The following will be kept constant in all my experiments.

Temperature: The whole investigation will be done at lab temperature. I will have to look over the fact that the temperature will be different on different days, although the effect on my results should only be very small.

Pressure: The investigation will be done at normal air pressure

Amount Stirred: I will stir each run once just to mix the reactants totally and then leave it. I will try to mix each run by the same amount although any inaccuracies which occur here are unavoidable at the level I'm carrying out the experiment.

Experiment 1

I will vary the concentration of the Na2S2O3, by diluting it with deionised water. I will keep the total volume constant at 11ml.

Na2S2O3

HCl

H2O

0

0

9

8

2

7

3

6

4

5

5

4

6

3

7

2

8

9

Experiment 2

I will vary the concentration of the HCl again by diluting it with deionised water. The total volume will be 11ml.

Na2S2O3

HCl

H2O

6

5

0

6

4

6

3

2

6

2

3

6

4

What I will measure:

I am going to measure the time it takes for a set amount of sulphur to be produced. This amount, x, is the amount it takes for the black cross to be obliterated. I will assume that this is the same for each run. With these results I can say that x grams of Sulphur are produced in, t, seconds. Therefore the rate of Sulphur produced per second will be x/t. Since l just want to compare all the data I collect the value of x doesn't really matter. To male things simple I will say x=1, so the rate of the reaction is 1/t.

Ideally I want to measure the initial rate of the reaction. As the reaction takes place the reactant particles are being turned into product particles. So the concentration of the reactants decreases with time. I don't want to have to wait for too much Sulphur to be produced though as long as I keep x constant it shouldn't matter that much. The more Sulphur that I wait for to be produced the more under the actual value my rate times will be.

Apparatus:

Two 10cc measuring cylinders accurate to (0.1ml

2M HCl

0.25M Na2S2O3

distilled water

paper with a cross drawn on it

stopwatch accurate to (1 second

50cc glass beaker

Thermometer accurate to (0.5(C

Method:

Measure out the desired quantities of HCl and Na2S2O3 at the correct concentrations. Put the Na2S2O3 into the 50cc glass beaker which is placed over a black cross. Add the HCl and start the stopwatch as soon as all the HCl is added. Stir the solution once to mix the reactants. Stop the stopwatch as soon as you can no longer see any of the cross. I will say this is the time taken for x grams of sulphur to be produced. Record the time and repeat each run twice and three times if necessary.

Safty Precautions:

I will be very carful when using acid as it is corrosive. I will wear a lab coat and safty specs at all times.

DOING STAGE

Changes to original plan:

In the planning stage I said I would use 0.25M Na2S2O3(aq), in fact what we were given was a 0.16M solution. When I was carrying out the planned experiment for the concentration of HCl I noticed that there wasn't any real change in the times which I was getting. In view of this I tried using smaller concentration to see what effect that would have. The problem now was that even though I did notice a change in the times I recorded the values I was using were so small (0.2ml) that I couldn't be sure of much accuracy in my measurements. I scaled the volumes up by 5 (the concentration would stay the same) to compare the rate times which I got. By doing this I had to use a different beaker in which to do the experiment which meant that the depth of the solutions was now 2.6cm, rather than 1.4cm. Therefore the value for x ( the amount of sulphur needed to blot out the cross) was different. I will call this new value y grams of Sulphur. I then decided to try the smaller volumes again but this time doing my measurements with a buiret to see if that made any difference to the level of accuracy.

Results, Observations and Measurements:

See next page for results tables and graphs

Since I did the experiment over several days the temperature was likely to change. Just to see I measured the temperature of all the solutions on each day. I noticed that the temperature of the reactants stayed constant all the time the reaction was taking place.

Solution

Temp 1(ºC)

Temp 2(ºC)

Temp 3(ºC)

Temp 4(ºC)

Temp 5(ºC)

H2O

21.5

20

21.5

23

23.5

Na2S2O3

22

20

21.5

23.5

23.5

HCl

22

20

21.5

23.5

23.5

Reactants (mixed)

22.5

21

22

24

23.5

On the graph showing rate of reaction varying with concentration of Na2S2O3, I have joined all the points and then drawn a best fit straight line through the origin.

On the other graph showing rate of reaction varying with concentration of HCl, I have plotted 3 lines. Graph 1 shows the rate of reaction when the volumes were measured with a measuring cylinder. Graph 2 shows the rate of reaction when the volumes were measured with a buiret. Graph 3 shows the rate of reaction when I scaled the volumes up. I couldn't really measure the mass of Sulphur produced therefore I said that I would just say x=1. This is fine until I come to saying what y is. Even though y will be greater than x I have just used y=1 as well.

REPORTING AND EVALUATING STAGE

Patterns in my results:

Looking at my first variable, concentration of Na2S2O3, the graph shows the relationship to be directly proportional since the points fit around a best fit straight line through the origin. Looking at the results table when I double the concentration from 0.016M of Na2S2O3 to 0.032, the rate of reaction doubles from 0.002 to 0.004. In the planning stage I predicted that the line on the graph would level off at a certain point. In fact on my graph this hasn't happened, this is because with the volumes I used the HCl was always in excess. At the maximum concentration of Na2S2O3 the ratio of Na2S2O3 :HCl was 1:12.5. Had I used higher concentrations of Na2S2O3 then I might have seen this levelling off effect. Other than this the relationship was as I predicted.

The results for the second variable, concentration of HCl, are slightly more complicated. If you look at all three graphs then you can see the same pattern in each, which is the one I predicted. They start off as a straight line and then level off at a fairly constant rate. The levelling off effect is much more visible on graphs 1 and 2, when the lower volumes were used, but it seems to occur at the concentration just below 0.4M of HCl. On graph 3 the rate of reaction at the greatest concentration is much higher than that of graphs 1 and 2. The curves are not really smooth curves, most of which will be down to experimental error. You can see though that the measurements done with a buiret seem more accurate.

Scientific explanation for my results:

As I predicted in the planning stage you can see the effect of the rate doubling when the concentration of the variable doubles (shown by the straight line through the origin). This (as I have said in the planning) is because when I double the concentration of X (either the Na2S2O3 or the HCl, which ever is being varied.)then I am doubling the number of X molecules in a given volume. If you keep the number of Y molecules (either the Na2S2O3 or the HCl depending on what X is) constant then what was originally a ratio of X:Y now becomes 2X:Y. Therefore the chance of the Y molecules colliding with the X molecules has doubled. I will assume that the number of successful collisions doubles proportionally. This will mean that the rate will double.

In order for both reactants to react fully the ratios of Na2S2O3:HCl should be 1:2. When we look at the concentration of HCl when it no longer seems to have any effect on the rate of reaction (when concentration of HCl is approximately 0.40M), the ratio is 1:2.5. This is just above 1:2, when both reactants will react fully.

Comparing both variables, you can see that whichever variable isn't in excess is the one on which the rate of reaction depends. The breakdown of the thiosulphate ion is the slowest step in the reaction therefore the one on which the rate depends whilst there is enough HCl for it to react with. When you decrease the concentration to a point where there isn't enough HCl for the thiosulphate ion to react with the reaction becomes dependent on the concentration of the HCl. Therefore now the slowest step of the reaction involves the HCl.

When I measured the time taken for y grams of Sulphur to be produced, I assumed x=y. In fact the times (in seconds) for the cross to be obliterated will be quicker when the depth is greater. Therefore this will mean that the rate values will be larger, which explains why graph 3 has much higher values then graphs 1 and 2.

How I could have improved the accuracy:

There were several sources of inaccuracy in my experiment. One which I have looked at was the temperature. I can see that it did vary and although it didn't seem to have too greater effect on my investigation, I could improve on this by using a water bath preheated to 20ºC. The method of measuring rate was probably the most effective although different equipment would improve the set-up a lot. Instead of just using a cross and a stopwatch you really need something more accurate. This set-up leads to quite a lot of human error. How do you know when exactly the same amount of sulphur has been produced? How can you be sure exactly when to stop the stopwatch? To improve on this you could use a colorimeter. Here a narrow beam of light is transmitted through the solution to a photocell. A current is generated in the photocell which is proportional to the amount of light transmitted through the solution. The meter is actually calibrated to show the fraction of light absorbed by the solution and therefore is proportional to the concentration of the coloured substance( in our case sulphur). Using this method you could measure exactly the time it took for a known amount of sulphur to be produced. This would have cut out the inaccuracies when I changed my total volume, and began to measure the time taken for y grams of Sulphur to be produced.

You could improve even further on how you measured the rate using this set-up by taking values for the concentration of sulphur over a period of time. You could then plot a graph of time against concentration and to find the rate at a particular instant. You would then end up with a different line for each concentration of reactant. From these lines you could find the initial rate of reaction by taking a tangent of the curve where it starts. This would give you very accurate values for the initial rate of the reaction. Even though I tried to measure the initial rate of the reaction what I actually measured was an average rate of reaction. I measured the rate of reaction over a period of time. The further the reaction is to completion the more error will creep in.

An alternative method of measuring the rate would be to monitor the pH. This probably would work although I doubt if it would be as effective as measuring the concentration of Sulphur. Monitoring the pH would be measuring the concentration of H+ ions. The results which you would get would be affected by the fact the SO2 (aq) will be slightly acidic, but this shouldn't be too much of a problem. The success of this method will depend on what part the H+ ion plays in the breakdown of the S2O32- ion.