The amount of copper deposited on the cathode and lost from the anode depends on the number of electrons passing through the circuit, i.e. upon the charge passed through the cell. Now the charge passed, q (in Coulombs), is related to the current. I) in amps) and time, t (in seconds), by Faraday's law:
Therefore I will predict that the mass change of the copper electrodes is directly proportional to the current and the time.
Positive ions (cations) are attracted to the negative electrode (cathode).
Negative ions (anions) are attracted to the positive electrode (anode).
When the ions present get to the electrodes the are oxidised or reduced.
Oxidation and Reduction:
Oxidation is loss of electrons. Reduction is gain of electrons.
At the Anode (positive electrode), the Anions (negative ions) lose electrons, and are oxidised. At the Cathode (negative electrode), the cations (positive ions) gain electrons, and are reduced.
Aim:
In this experiment there will be two factors that I will be testing for:
1) Changing time
2) Changing current
Each factor will require a different experiment and thus allowing two experiments to occur.
Hypothesis: It is possible to predict that the relationship will be directly proportional between the time the current flows and the mass of Copper deposited on the Cathode (negative electrode). I can therefore predict that if I double the time of the experiment, I will therefore be doubling the charge. This statement can be supported by both of Faraday’s Laws.
Faraday’s First Law of electrolysis states that:
“The mass of any element deposited during electrolysis is directly proportional to the number of coulombs of electricity passed”
Faraday’s Second Law of electrolysis states that:
“The mass of an element deposited by one 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.”
Prediction:
As in all cases using electrolysis, I predict that the positive copper ions will move towards the negative cathode and likewise the negative ions will move towards the positive anode. As for the factors that I am testing with, I predict that they will affect the results slightly as for time. If I were to speed up the time then I would need a catalyst and thus will affect the actual amount of ions moved from the solution to its opposites force. (i.e. Positive – Negative). For the current, If I were to increase it, in turn it would make the ions move at a faster rate and possibly again speed up the reaction. However if I were to slow it down then It would make the particles move at their own steady speed.
Apparatus:
-
Copper sulphate solution aq
- Copper electrodes (Copper Strips)
- Power pack
- Crocodile clips
- Beaker
- Amp meter
Method:
(1) Current
- We set a certain current and connected the circuit as shown in the diagram.
- We connected the circuit to a power supply and timed the reaction at intervals of 2 minutes. We went up to 20 minutes
- We then noted our results at every interval and drew up graphs on our results.
- We continued the experiment by increasing the current for each experiment, however in each experiment we changed the electrodes to make it a fair test.
- To make this a fair test we timed each current experiment twice to obtain more accurate results.
- We tested from 0.5 amps to 2amps.
(2) Time
- For the experiment relating to time, we selected a certain current to test the time intervals. We connected the circuit as shown below.
- We then connected the circuit to a power supply and timed the reaction at intervals of 2 minutes as well.
- Overall we kept the current the same as this was not an experiment to test the current, however to test the time intervals.
- We continued up to 20 minutes each time and wrote down our results.
- We then changed the time intervals from 2 minutes to 5 minutes in between and tested it all the way up to 30 minutes.
Diagrams
Results:
(1) Current
0.5amps
1.0amps
2.0amps
(2)
4amps
Analysis:
Overall my results came out similar to what I had expected them to be.
Cu(s) (r) Cu2 +(aq)+2e- (oxidation)
During the electrolysis of a copper salt is the reverse of the cathode reaction:
Cu2 +(aq) + 2e- (r) Cu(s) (reduction)
So for every two electrons passing through the external circuit, one copper ion should be formed at the anode and one copper ion discharged at the cathode. One would expect the mass loss of the anode to equal the mass gain at the cathode, as explained earlier, for every two electrons, at the cathode one copper ion is discharged, whilst at the anode, one copper ion is formed This can be explained with the ionic theory, which basically states that the electrons flow away from the cathode, to the anode where the Cu2+ ions take 2 electrons from the negative electrode and become Cu atoms, thus mass loss at cathode = mass gain at the anode. This does support the prediction, as the two lines are at most only 0.018 grams apart, or 10% inaccurate, using the formula difference ¸ theoretical X 100. The other pattern is that the mass change µ current, This is shown by the construction lines on the graph, which show that when the current is 0.2A, the mass lost at the anode is 0.035g, and the mass gained at the cathode is 0.04g, and when the current doubles to 0.4A, the mass change also doubles as the mass lost at the anode is 0.07g, and the mass gained at the cathode is 0.078g. This is because, as explained in the planning section, The amount of copper deposited on the cathode and lost from the anode depends on the number of electrons passing through the circuit, i.e. upon the charge passed through the cell. Now the charge passed, q (in Coulombs) is related to the current. I) in amps) and time, t (in seconds), by Faraday's law: CjbCKT1I Visit
Evaluation:
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 the ions "stick" to the anode, and so end up at the bottom of the solution. This happens most at higher levels of current, and causes the mass lost at the cathode to be greater than the mass gained at the anode. Also the temperature of the solution raised at higher currents by 5° C this would cause fewer ions to turn to copper at the anode, and make the current more, as there is less resistance. The size of the electrodes was also never exactly the same, as they were reused, so the amount of electrolysis differed from experiment to experiment. The separation of the electrodes was a small source of error, as they were not always exactly the same distance apart. The current, which was controlled with the rheostat, was not always the same, as the amount of copper decreases, so does the resistance, and so the current increases. 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, could have caused other errors. The rest are cut of with out rounding. Therefore this experiment could have been made more accurate by using lower current values, with the same size and separation of electrodes, controlling the current so that the temperature is constant, and the current more accurately controlled, and using a more accurate ammeter and a balance which rounds the other decimal places. My results showed many inaccuracies, shown by the accuracy bars on the graph. Which show the highest value and the lowest, with the average in the middle. This shows that for the 0.20A reading, the anode difference is 0.01A, and the cathode difference is 0.02A, both very small variations. For the 0.40A reading, the anode difference is 0.07A, a much greater difference, and the cathode variation was smaller, at 0.02A. The 0.60A anode difference was only 0.01A, and the cathode was the same. The 0.80A anode and cathode variation were also 0.01A. The final reading, 1.00A anode difference was 0.03, and the cathode variation was 0A. This nearly fits the pattern of the greatest variation being at the top, except for the 1.00A cathode variation of 0A. This increasing variation is caused primarily by two things, firstly the temperature of the solution increases more at higher current values, so the ions travel faster, and so do not stay on to the anode as well, and secondly the increased current itself has the effect of making less ions sticking to the cathode. The anomalous result for in the 0.40A value for the anode was probably caused by one or both of the crocodile clips touching the solution, so less electrons flow through the copper, and so less are transferred to the cathode.
The range of my results were from 0.5A to 2.00A, with an average discrepancy of 0.02A from the average reading, which although there was one large anomalous result is quite small, is quite a small variation, therefore The evidence is strong enough to say that the mass lost at the cathode equals the mass gained at the anode, and that q µ i, as the greatest error was only 0.01g, or 12.5%.
If this experiment were to be done more accurately, I would have to use more accurate apparatus, such as a newer ammeter, a balance with more digits, a more accurate way of controlling the current, maybe with a computer, and likewise with the temperature. I also could have kept the size and separation of the electrodes the same. I also could have made sure that the crocodile clips were completely out of the electrolyte. Also I could have taken a much wider range of readings, from 0.01A to 10A at smaller intervals, and I could have timed for different times, and I could have investigated the other variables, such as the temperature of the electrolyte, the concentration of the electrolyte, the separation of he electrodes, and the size of the electrodes. The Electrolysis Of Copper Sulphate Solution Using Copper Electrodes
