The Reaction.

GCSE Science

The Reaction

The word equation for this reaction is:

Copper Sulphate + Iron → Iron Sulphate + Copper

The balanced chemical equation for this reaction is:

CuSO4(aq) + Fe(s) → FeSO4(aq) + Cu(s)

This reaction falls under several categories of chemical reaction, namely:

Displacement reaction
Exothermic reaction
Redox reaction

I will now explain each of these reactions in detail.

Displacement Reaction

In a displacement reaction, the more reactive substance in the reaction takes the place of the less reactive substance. As can be seen from the equations written above, the iron takes the place of the copper in the copper sulphate solution. The reaction can therefore be classed as a displacement reaction.

Redox Reaction

‘Redox’ stands for REDuction / OXidation, meaning that in the reaction, there is one substance which gains oxygen and another which loses it. Reduction and oxidation always happen together – if one thing is reduced, another is oxidised. By looking at the equations, it is easy to see that iron has been oxidised – gained oxygen, while copper has been reduced – lost oxygen, by losing the sulphate, SO4, to the iron.

However, it is also possible to tell which substance has been reduced and which has been oxidised by looking at how the charges of the substances change in the equation. This can be done by producing an ionic equation:

1) Highlight the spectator ions:

Cu2+(aq) + SO42-(aq) + Fe0 (s)→ Fe2+(aq) + SO42-(aq) + Cu0 (s)

2) Remove the spectator ions to leave an ionic equation:

Cu2+(aq) + Fe0 (s)→ Fe2+(aq) + Cu0 (s)

Since copper goes from having a positive charge, i.e, too few electrons, to having no charge, it must obviously have gained electrons in the process. Gaining electrons always means reduction. Iron, on the other hand, has gone from having no charge, i.e, the full amount of electrons, to having a positive charge – too few electrons. It has evidently lost some of its electrons, and a loss of electrons always means oxidation.

Exothermic Reaction

A reaction is always both exothermic and endothermic. There are two parts to every reaction – the breaking of the bonds of the reactants (in this case, iron and copper sulphate solution) and the making of the bonds of the new products (iron sulphate and copper). The breaking of the reactant bonds requires energy, and is therefore endothermic. In the process of making the product bonds, energy is given out - exothermic. In an endothermic reaction, more energy is taken in to break the bonds than is given out later on. In an exothermic reaction, more energy is given out while making the new bonds than was taken in during the breaking of the old bonds.

In my investigation, I recorded the increase in temperature. This is an indication of the amount of energy given out overall during the reaction.

Background Information

Before I began my investigation into the factors affecting temperature change in the reaction between iron and copper sulphate, I had to make a copper sulphate solution. Since my experiment would involve changing the concentration of solution, while the amount of iron remained constant throughout, I decided to begin with a molar solution and then reduce the concentration of the solution by adding water to it. To make a molar solution of CuSO4, ...

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In my investigation, I recorded the increase in temperature. This is an indication of the amount of energy given out overall during the reaction.

Background Information

Ar Copper (Cu) 63.6

Ar CuSO4 Ar Sulphur (S) 32

+ Ar Oxygen (O) x 4 64

159.6

I would use 159.6g of copper sulp hate if I wanted to make a litre of the solution. However, so much solution was unnecessary and wasteful. I therefore divided this figure by 10:

159.6 = 15.96.

I decided to mix 15.96g of copper sulphate with 100ml of water – one tenth of a litre. This way, I would still have a molar solution, but I would have less of it. To actually make the solution, I had to heat the water over a Bunsen burner and gradually add the copper sulphate powder while stirring the solution. Otherwise, the powder would simply crystallise as it reached the bottom of the beaker. I therefore had to allow the solution to cool down completely before using it in my experiment – otherwise, the results would certainly have been affected, and would thus be inaccurate.

My plan was to carry out six tests, with a different concentration of solution for each test. There would be altogether 20 ml of solution in each test, with the proportion of solution to extra water decreasing in each test, until, in the final test, I would have 20 ml of only water:

1 Molar solution - 20 ml copper sulphate : 0 ml water

0.8 M solution - 16 ml copper sulphate : 4 ml water

0.6 M solution - 12 ml copper sulphate : 8 ml water

0.4 M solution - 8 ml copper sulphate : 12 ml water

0.2 M solution - 4 ml copper sulphate : 16 ml water

0 M solution - 0 ml copper sulphate : 20 ml water.

Now that I had determined the quantities needed to make 100ml of copper sulphate solution, I had to work out how much iron would be needed – the amount used in each test would not vary.

Ar Fe = 55.85, i.e, in grams,1 mole of iron is 55.85g

If I was going to use 100 ml of water for each test, I would use a tenth of this amount (5.585g) in each test. However, as I explained before, there would be 20 ml of solution in each test. I therefore multiplied 5.585g x 0.2. The result was 1.117g. Therefore, this was the amount of iron I would be using in each test. Because the scales I would be using only measured up to one decimal place, I rounded this figure down to 1.1g of iron per test. I could now proceed with my investigation.

Apparatus

Anhydrous copper (II) sulphate - with which to make the solution for the experiment

Distilled water - to dissolve the copper sulphate to make solution

Fine iron filings - to put into the solution to carry out investigation

Electronic weighing scales - in order to weigh out iron and copper sulphate quantities fairly accurately

Boiling Tubes - in which to place solution and iron

Measuring cylinders - to measure out quantities of distilled water and copper sulphate solution

Thermometer - to record highest temperatures reached because of the exothermic reaction

Stirring rod - to stir the copper sulphate while making solution

Bunsen burner - to help the copper sulphate dissolve into the distilled water

Hypothesis

I predict that as the concentration of the copper sulphate solution gets weaker, the amount of energy released will decrease proportionately, i.e., if there is a change of 10 degrees Celsius when the iron is in a 0.8 M solution, there will be a change of 5 degrees Celsius in a 0.4 solution. However, if the concentration of solution was more than 1 mole, there should be no additional temperature rise; the ratio of 1M of solution to 1M of iron would produce the maximum temperature rise.

Therefore, if I drew a graph taking into account all these points, it would look something like the following sketch:

To ensure a fair and safe test:

I will use the same scales each time I measured out quantities of copper sulphate powder or iron filings, in case the scales were slightly biased.
I will make sure that the bulb of the thermometer was always covered by the copper sulphate solution
Tests will be repeated in order to make sure of results obtained.
The copper sulphate solution will be measured very carefully in order that the volume remains constant in all the tests.
The thermometer will be left inside the test-tube until the maximum temperature has been reached.
I will wear safety goggles while making the copper sulphate solution.
Because the solution will be made over a Bunsen burner, I will not use it in the reaction until it has cooled down.

Diagram

Below is a diagram of my experiment, for the 1 molar solution of copper sulphate.

Method

Make copper sulphate solution, with quantities given above.
Pour 20 ml of the solution into a boiling tube
Weigh out 1.1g of fine iron filings. Add to the solution.
Immediately insert a thermometer into the boiling tube. Record the highest temperature that is reached.
Repeat steps one to four 5 times, but each time decreasing the concentration of the solution, as shown above.
Repeat all the above steps once, in order make sure of results.

Results

I drew up several tables in which to record my results. In the first, I recorded the highest temperature compared to the strength of the solution for both experiments, and then took an average of each result. I then made a second table in which I recorded the increase in temperature from room temperature rather than the actual temperature reached and also included an average. I then drew a graph using the data from my second table, comparing the strength of the solution to the temperature increase.

Table 1

Table 2

N.B. Room temperature was 18°C

I noticed that the iron turned a reddish colour very soon after it entered the solution. This indicates that iron is more reactive than copper and had displaced the copper in the solution in my experiment.

Conclusion

Overall, I conclude that my prediction was correct – the temperature did increase more or less proportionately:

0.8 solution increase = 5

0.4 solution increase = 2.5

5/2.5 = 2

Not all my results were so exact; however this example shows that my experiment was fairly accurate. The line of best fit in my graph went through (0,0) and also showed that temperature increased in direct proportion to increase in the molar strength of the solution. None of the points hit the line of best fit exactly; however, there was quite an even correlation.

I did not have enough time to test whether the temperature increase would reach a peak with a ratio of 1M : 1M of solution and iron –i.e., if the solution was, for example, a 1.5 molar solution, this would not have any effect on the temperature rise. I therefore took this assumption to be correct.

Evaluation

1) Fortunately, I obtained no anomalous results. However, the temperature increase in general was not as large as I expected. One reason for this is probably that much heat was lost through radiation from the test tube, as there was no insulation to prevent such an occurrence. As I did not cover the tops of the test tubes, it is very likely that heat would be lost through convection currents. Had I covered the test tubes, the temperature increase would have doubtless been larger. The best choice would have been to cover them with something with a silvery surface, such as aluminium foil, which would have reflected any heat radiated out back into the test tube.

2) Another reason for my unexpectedly low results is the size of the iron filings. I used very fine iron filings, which meant that there was a large surface area for the copper sulphate to react with. Therefore, it is likely that most of the iron did react. However, during this reaction, a layer of copper will have formed over each iron filing. This would mean that the inside of each grain of iron would not have been able to react with the copper sulphate. Therefore, in effect, less than one mole of iron was added to the copper sulphate because not all of it was able to react. This resulted in lower temperature increase than expected. It would have been very hard to solve this problem as I would have had to work out how much a mole of iron was, taking into account the size of each iron filing as this would determine the size of the copper layer surrounding it. The best I could have done was to use even finer iron filings; however, I would still have had the same problem although to a lesser extent.

3) As explained above, the iron filings had a very large surface area because of their minute size. Since iron is a fairly reactive metal, it would have reacted with oxygen while in storage, resulting in a layer of iron oxide covering it. Therefore, when I weighed out a mole of iron for my experiment, I was weighing out both iron and iron oxide. This problem could have been reduced by storing the iron very carefully so that it should come into contact with the least amount of oxygen possible.