Method
I poured 200ml of Copper Sulphate solution into a beaker. I got two copper strips to use them as electrodes and cleaned them thoroughly using sandpaper. I then weighed the two electrodes and recorded them so I could compare their weights after the experiment. These electrodes were then put into the beaker of copper sulphate, and then they were connected up to an ammeter and a power pack. A consistent current of 0.54 amps (the variable resistor would not allow a lower current) was passed through the solution. If I had used a higher current then too much copper may have collected at the cathode and then the adhesion would not be as good and parts may have fallen off and caused innacuracies. This was done for various different times. Each time I did the experiment I turned the power pack off when the time was up, and then I removed the copper electrodes from the solution. I washed these in water and then dried them by dipping them into propanone. When both of the electrodes were clean and dry I weighed them both and recorded their weights. I will alterate the electrodes so that one doesn’t get too built up. From the recordings from all the different times I will be able to plot the results onto a graph and it will be easy to see how the amount of time that the current is allowed to flow affects the deposition of copper. From these results I should be able to suppost Faraday's first and second law.
Safety
I must be aware of the dangers of the experiment. I will use goggles at all times because the solution is irritant to the eyes. The glass beaker must be handled with care and water must be kept away from the electrics at all times and the electrics must be turned off when not in use.
Electrolysis Theory
Electrolysis is the process of splitting an electrolyte into its ions by passing an electric current through it. Obviously for this to work, the ions have to be able to move with relative freedom, and so the electrolyte has to be either molten or part of an aqueous solution.
The electrolysis of an aqueous solution of copper sulphate using copper electrodes results in transfer of copper metal from the anode to the cathode during electrolysis. The copper sulphate is ionised in aqueous solution.
The positively charged copper ions migrate to the cathode, where each gains two electrons to become copper atoms that are deposited on the cathode.
Reaction at cathode: Cu + 2e ⇒ Cu (Reduction)
At the anode, each copper atom loses two electrons to become copper ions, which go into solution.
Reaction at Anode: Cu ⇒ Cu + 2e (Oxidation)
The sulphate ion does not take part in the reaction and the concentration of the copper sulphate in solution does not change. The reaction is completed when the anode is completely eaten away.
The result of this is that the anode loses mass, the cathode gains mass and the solution remainsthe same concentration. Therefore the loss from the anode should exactly equal the gain at the cathode.
Faraday’s laws say that ‘the mass of any element deposited during electrolysis is directly proportional to the number of coulombs of electricity passed’ and that ‘the mass of an element deposited by 1 Faraday of electricity is equal to the atomic mass in grams of the element, divided by the number of electrons required to discharge one ion of the element’. Therefore by changing the amount of time the current is applied for (and thus the number of coulombs of electricity passed), I should be able to vary the amount of copper deposited proportionately. Also, it means I am able to make accurate predictions before the experiment is actually carried out and use them to test how accurate my results are.
Predictions
The following formulae will be used to predict my results:
Quantity of charge (coulombs) = current (amps) x time (seconds)
Number of moles of electrons = charge/96500
Cu2+ + 2e(-) → Cu (for one mole of copper, 2 moles of electrons are needed)
Mass = moles x RFM (to convert to mass of solid)
0 MINS
No solid will be formed
5MINS
0.54 x 300 = 162
162/96500 = 0.00167
0.00167/2 = 0.000835
0.000835 x 64 = 0.05344
0.053g
10 MINS
0.54 x 600 = 324
324/96500 = 0.00335
0.00335/2 = 0.00167
0.00167 x 64 = 0.10688
0.107g
15 MINS
0.54 x 900 = 486
486/96500 = 0.00503
0.00503/2 = 0.002515
0.002515 x 64 = 0.16096
0.161g
20 MINS
0.54 x 1200 = 648
648/96500 = 0.00671
0.00671/2 = 0.003355
0.003355 x 64 = 0.21472
0.215g
25 MINS
0.54 x 1500 = 810
810/96500 = 0.00839
0.00839/2 = 0.004195
0.004195 x 64 = 0.26848
0.267g
Chemistry Coursework - Obtaining
No alterations to the proposed method were needed because the obtaining aspect of the experiment was very successful. I think the the results obtained are very accurate when the equipment used is put into consideration. They were accurate in that they all followed a pattern and they were reasonably close to the predicted results the largest difference from the predicted results being +0.015 grams. The only result that could be considered anomalous has been highlighted yellow and that is not particularly extreme.
In order to make the experiment more precise and reliable I did 3 repetitions for each time, anymore would've taken far too long. I then calculated the mean of the three. This helped to balance out any figures that were too high and too low and thus gave me a more reliable figure.
Observations
There werent really any observations to be made except for the obvious. The obvious being that there the anode became slightly thinner for each reading and small amounts of copper collected on the cathode. Also, a sludge was depositied in the bottom of the beaker.
Chemistry Coursework - Analysis
Problems
- There were several sources of error in this experiment as none of the results were 100% accurate. These errors could have been caused by the fact that not all th ions "stick" to the cathode, and so end up at the bottom of the solution. This happens most when there is more charge, this would explain why, for the longer time readings, the lost at the anode may have been greater than the mass gained at the cathode. The fact that not all the ion "stuck" onto the cathode can be explained by the fact that the adhesion of the copper added onto the cathode by electrolysis would be nowhere near as good as that between the copper atoms already part of the electrode. This would have reduced the ability of the copper as a conductor and the adhesion would be poorer. The surface area of the electrodes in the solution was never exactly the same because I couldn’t clip them in in exactly the same place each time, we also alternated between to stop one getting built up and so this wouls cause further error because they are not exactly the same size. The size of the electrodes may have differed awell because they reused. Because of the errors concerning the electodes the amount of electrolysis that took place would've varied eacht time. Other errors could have been caused by the apparatus, such as the ammeter, which is quite old and may not be perfectly calibrated, and the scales which only show the mass to 2 decimal places. The rest are cut off without rounding. Even with great care, small amounts of copper may well have been removed during the washing of the electrodes each time, this would also cause innacuracies.
Conclusions
My results show that the longer the electrolysis experiment is carried out for, the more copper will be lost from the anode and collected at the cathode, and that this is directly proportional to the current used and the mass of the substance.
In the results all the actual differences were within 0.015grams. Although the numbers were small this still suggests that all results were reasonably accurate and close to the predicted results. This tells me that the results are very good evidence to support Faraday’s 1st and 2nd laws of chemistry and electricity – ‘the mass of any element deposited during electrolysis is directly proportional to the number of coulombs of electricity passed’ and ‘the mass of an element deposited by 1 Faraday of electricity is equal to the atomic mass in grams of the element, divided by the number of electrons required to discharge one ion of the element’. This is exactly what I predicted in my hypothesis, and as I expected a sludge was deposited on the bottom of the beaker.
Chemistry Coursework - Evaluation
I think the the experiment was a success. I developed a simple method, yet it proved very effective. It has provided me with good enough results to support Faraday's 1st and 2nd laws and proven my hypothesis correct. I think that the results were quite good and reliable because there was only one abnormal result and any other highs and low would've been removed by calculating a mean. I am sure that all inaccuracies in the results came from experimental errors because there no results that particualrly extreme. I am sure of this because the results were accurate enough to show that my science was right which only leaves mistakes to be cause during the experiment.
I this experiment were to be done more accurately I would have to use more accurate apparatus, such as a new ammeter, a new blance with more digits in the display, a more accuarate way of controlling the current (e.g. with a computer). I would use new electrodes for each reading and they must be exactly the same size and weight.This would also ensure that the same surface area is submerged into the solution. I would also use a different method of holding them (not crocodile clips) which would ensure that the same surface area is submerged into the solution each time. I would also take a much wider range of readings at many more, smaller, intervals.This would produce more precise and reliable results and a better graph. I would take more time and care could have been taken over washing the electrodes after each reading to try to ensure as much of the solution but as little of the copper was removed as possible. I would use entirely accurate electronic timing device to shut off the current after exactly the required amount of time to remove the possibility of human error. I could investigate the other variables, such as the temperature of the electrolyte, the concentration of the electrolight, the distance between the electrodes and the size of the electrodes. However investigating these would not allow me to make predictions or refer to Faraday's laws.
I have used my method successfully to perform an accurate experiment and obtain accurate results form which I have proven my hypothesis correct and acquired only one abnormal result. This abnormality was almost certainly caused by one or more of the following:
∙ Washing off some of the copper when cleaning the elctrode with propanone.
∙ The copper not sticking to the electrode and fallen into the solution when removing the electrode.
∙ Apparatus problems; there may have been something adding weight on the scales thus not giving the true weight of the electrode.
∙ Human error; while leaving the experiment someone may have stopped the stop watch and then starting it again giving us a longer time for electrolysis or someone may have done something to the circuit by mistake,the probability of this is very small though.
If I wanted to prove Faraday's 1st and 2nd laws and my hypothesis even further I would do the electrolysis of sulphuric acid using graphite electrodes. This woul produce hydrogen gas at the cathode and oxygen gas at the anode. Because they are gases I would have to measure the volume rather than the gas. I woiuld do this by collecting the gas in a measuring cylinder filled with water.
Diagram