Method
- The Setup
- Collect all required equipment.
- Clean the 10 copper electrodes thoroughly using sandpaper to remove impurities that could inhibit electrolysis.
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Assign half of the electrodes to be the anodes and mark them with A1 to A5 using a permanent marker, then the other half will be the cathodes, mark these with C1 to C5(there should be 5 C’s and 5 A’s). Be sure not to write on the side that will go into the solution as this could inhibit oxidation or reduction!
- Weigh the cathodes using the electronic balance and record this value as initial mass for each of the cathodes (C1 – C5).
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Collect 5, clean, 100 mL beakers and fill each with 100 mL of 1mol L-1 copper sulphate solution.
- Next, set up a circuit such that, the negative power pack terminal connects to the variable resistor using banana clips. From the variable resistor, connect a banana clip out to the ammeter’s positive terminal.
- Draw out a banana clip/alligator clip from the ammeters negative terminal and connect it to the first, C1, copper cathode. Connect the rest of the cathodes, respectively from C1 to C5 using alligator clips in series.
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Now, the positive terminal of the power pack connects to the anodes such that, the first anode connects to the second anode and so on just the like cathodes. Do this systematically from A1 to A5 using alligator clips until 5 anodes are connected in series.
- Place the each electrode pair in a beaker filled with sulphuric acid prepared previously about 4 cm into the beaker and fix both electrode pairs to polar sides of the beaker by bending. This controls the surface area and the distance between the electrodes. The circuit should now look something like the schematic to the right (figure 1), check this, and continue if correctly assembled.
Five electrolytic cells have now been connected in parallel and ready for electrolysis.
Figure 1.
- The Electrolysis
- Adjust the current to required starting amperage using variable resistor to 0.2 amps. Do this by connecting the multimeter straight to the variable resistor excluding the electrolytic cells.
(Figure 2)
- Now, connect the electrolytic cells back onto the circuit and turn on the power. Time for 2 minutes using the stopwatch, making sure current remains constant on the ammeter. (figure 2).After 2 minutes has elapsed, turn off power and take out the cathodes carefully.
- Wash the cathodes with distilled water and dry by lightly padding with a paper towel.
- Weigh each cathode using electronic balance. Record this result in a table as final mass.
- Repeat this experiment for the currents of 0.4, 0.6, 0.8 And 1 Amps keeping time constant to two minutes. Be sure to sand each cathode for each of the trials.
- The Calculation
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Form the results collected, find the mass deposited on the cathodes for each of the 6 trials by using,
final mass - Initial mass = Net mass. Record this calculation as mass deposited.
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For each of the 6 trails find the average mass deposited by adding the mass deposited on each of the five cathodes in each of the 6 trial and dividing by 5. E.g. C1+C2+C3+C4+C55= average mass deposited
- Graph the data and compare results to theoretical calculated mass values according to faraday’s law
The ∆ mass or mass deposited on the cathode was calculated for each of the trials by the following equation:
∆ mass (mass deposited) = final mass - Initial mass
The results are shown in the table.
Average mass gained for each of the trials was calculated by the following equation:
C1+C2+C3+C4+C55= Average ∆ mass (mass deposited)
The results are shown in the table.
From observing the data, it can be seen that, on average, there is a gain in mass as current is increased. However, there are some anomalous results
From the experimental results previously, average mass deposited is graphed against current. As time is kept constant at two minutes, time can be excluded from the quantity of current on the graph. The raw ∆ mass results for each current setting are also shown on the graph for scatter reference.
From the graph, it can already be seen that there is a gain in mass as the current increases. To some extent, the average mass deposited seems to show relationship of direct proportionality between current and mass deposited on the cathode as hypothesized. An explanation will be given for the large scatter of results present in the raw data after the results of the experiment have been analysed.
∆Mass Theoretical calculation
The theoretical values for mass deposited on the cathode will now be calculated for each reading of current quantity in order to determine the accuracy of the results for mass deposited obtained in the experiment. The technique that will be used to calculate these theoretical mass values involves using stoichiometric methods and faradays law, which states the quantity of current is directly proportional to mass formed on the cathode. The theoretical calculations for each current quantity tested in the experiment will be demonstrated below and explained in detail.
Time is constant at 2 minutes
=120 seconds.
The first current reading is 0.2 amps. The total quantity of charge transferred will first be calculated for this value of current.
Total Quantity of charge = Current Ax TimeS ➔ 0.2 x 120 = 24 Coulombs
Charge of 1 electron: 1.60217733x 10-19C
To find number of electrons transferred.Number of e- =Total Quantity of ChargeCharge of 1e- ➔241.60217733x10-19= 1.5 x1020electrons
In one mole there exists
6.02 x 10ˆ23 electrons
Avogadro’s number
To find moles of electrons transferred.
Moles of electrons =number of electronsAvogadro’s number ➔1.5 x10206.02 x 1023= 0.0002487046632 moles e-
Cu2+ + 2e- ➔ Cu
- Reaction at the cathode:
From the equation it can be seen that 1 mole of Cu requires 2 moles of electrons to form.
∴moles of Cu= moles electrons/2 ➔ 0.00024870466322= 0.0001243523316 moles
Molar Mass Cu : 63.5
To find Mass of Cu
mass = moles x Molar Mass = 0.001243523316 x 63.5 = 0.008 g of Cu deposited (1.s.f)
- Second current reading is 0.4 amps.
Total Quantity of charge = Current Ax TimeS ➔ 0.4 x 120 = 240 Coulombs
- To find number of electrons transferred.
Number of e- =Total Quantity of ChargeCharge of 1e- ➔481.60217733x10-19= 3 x1020electrons
-
To find moles of electrons transferred.
Moles of electrons =number of electronsAvogadro’s number ➔1.5 x10206.02 x 1023= 0.00049766163 moles e-
- Reaction at the cathode:
Cu2+ + 2e- ➔ Cu
∴moles of Cu= moleselectrons2➔ 0.00024870466322= 0.0002487046632 moles
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To find Mass of Cu
mass = moles x Molar Mass = 0.001243523316 x 63.5 = 0.016 g of Cu deposited (2.s.f)
- Third current reading is 0.6 amps.
Total Quantity of charge = Current Ax TimeS ➔ 0.6 x 120 = 72 Coulombs
- To find number of electrons transferred.
Number of e- =Total Quantity of ChargeCharge of 1e- ➔721.60217733x10-19= 4.5 x1020electrons
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To find moles of electrons transferred.
Moles of electrons =number of electronsAvogadro’s number ➔4.5 x10206.02 x 1023= 0.0007474083 moles e-
- Reaction at the cathode:
Cu2+ + 2e- ➔ Cu
∴moles of Cu= moleselectrons2➔ 0.00024870466322= 0.00037375415 moles
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To find Mass of Cu
mass = moles x Molar Mass = 0.001243523316 x 63.5 = 0.024 g of Cu deposited (2.s.f)
- Fourth current reading is 0.8 amps.
Total Quantity of charge = Current Ax TimeS ➔ 0.8 x 120 = 96 Coulombs
- To find number of electrons transferred.
Number of e- =Total Quantity of ChargeCharge of 1e- ➔961.60217733x10-19= 6 x1020electrons
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To find moles of electrons transferred.
Moles of electrons =number of electronsAvogadro’s number ➔1.5 x10206.02 x 1023= 0.00099667774 moles e-
- Reaction at the cathode:
Cu2+ + 2e- ➔ Cu
∴moles of Cu= moleselectrons2➔ 0.00024870466322= 0.00049833887 moles
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To find Mass of Cu
mass = moles x Molar Mass = 0.001243523316 x 63.5 = 0.032 g of Cu deposited (2.s.f)
- Fifth current reading is 1 amp.
Total Quantity of charge = Current Ax TimeS ➔ 1x 120 = 120 Coulombs
- To find number of electrons transferred.
Number of e- =Total Quantity of ChargeCharge of 1e- ➔481.60217733x10-19= 7.5 x1020electrons
-
To find moles of electrons transferred.
Moles of electrons =number of electronsAvogadro’s number ➔1.5 x10206.02 x 1023= 0.00124584717 moles e-
- Reaction at the cathode:
Cu2+ + 2e- ➔ Cu
∴moles of Cu= moleselectrons2➔ 0.00024870466322= 0.00062292358 moles
-
To find Mass of Cu
mass = moles x Molar Mass = 0.001243523316 x 63.5 = 0.04 g of Cu deposited (2.s.f)
- Sixth current reading is 1.2 amps.
Total Quantity of charge = Current Ax TimeS ➔ 1.2 x 120 =144 Coulombs
- To find number of electrons transferred.
Number of e- =Total Quantity of ChargeCharge of 1e- ➔1441.60217733x10-19= 9 x1020electrons
-
To find moles of electrons transferred.
Moles of electrons =number of electronsAvogadro’s number ➔1.5 x10206.02 x 1023= 0.00149501661 moles e-
- Reaction at the cathode:
Cu2+ + 2e- ➔ Cu
∴moles of Cu= moleselectrons2➔ 0.00024870466322= 0.0007475083 moles
mass = moles x Molar Mass = 0.001243523316 x 63.5 = 0.047 g of Cu deposited (2.s.f)
The theoretical mass calculation results are shown in table format and graphed bellow.
From the graph’s constant slope, it can be seen that the relationship between mass and current is one that exhibits direct proportionality. For example, if current where to double in magnitude so would the mass deposited on the cathode. This makes sense since, as current increases, more electrons are passed into the cell and so should increase the amount of copper deposited at the cathode proportionally to the amount of electrons.
Comparison of results
The graph above displays both the theoretically calculated mass and the experimentally obtained mass results in respect to current. From this comparison, it can be seen that the results obtained in the experiment are poorly correlated to, and show neither accuracy nor precision to the theoretical mass data.
Interpretation and Evaluation:
This electrolysis experiment is was performed using 5 electrolytic cells attached in parallel. This was intended to increase the sample size and thus, minimize random errors by taking repeated measurements of the same action and calculating the average. It is important to repeat experiments when performing any experiment. However, repetition does not improve accuracy of results when systematic errors are present. Systematic errors do not change with repetition and will produce the same inaccurate result with good precision if present.
From observing the results and graphs, it seems that the results obtained in the experiment show no real pattern or strong enough correlation to indicate a relationship of direct proportionality between current quantities applied to the electrolytic cell and mass deposited on the cathode and comparing the theoretical results of mass and the experimental results graph shows that the two are very different. The first experimental result is higher than the theoretical value by about a factor of 10 and can be excluded as an outlier right away. The rest of the experimental results were a lot smaller than what was expected from the theoretical values. Mass of Cu gained in each trial barely increased over changing the current from 0.4 to 1.2 amps and remained somewhat constant at 0.008 grams as the current was increased, opposed to the theoretical results, which showed a constant increase in mass as current increased. This is odd as 0.008 grams was the expected theoretical mass value for a current of 0.2 amps. Either this could indicate a loss of mass, either into the solution or when drying with paper towels, or an error in delivering the correct current quantity or an error in measuring the mass deposited on the cathode.
The experiment was not successful in confirming the hypothesis. The anomalous results obtained are likely to have been caused due to both, systematic and random errors that occurred and were noted as follows.
Random errors
Random errors occur as inconsistent and unpredictable changes in the experiment (i.e. extraneous variables) and often due to human mistakes resulting from the experimenter's inability to perform a measurement in exactly the setting to obtain the exact the same result.
A major error that occurred in this experiment is due to the way the cathodes were dried after being removed from the copper sulphate solution. Paper towels were used to absorb the solution on the cathodes which could have removed a lot of the copper that was deposited on the cathodes post-electrolysis and thus reduced the reading for the change in mass. This could explain why the results obtained in the experiment are low compared to the theoretical results. To improve the method of drying the electrodes, a blow dryer could be used. This would ensure distilled water is evaporated and copper is left behind.
Circuit break.
During the experiment, the current fluctuated franticly and was very hard to control. The reading on the ammeter would often casually drop down to zero. Attempts were made to adjust the current back as soon as possible but often took a significant amount of time to set the current back to the required setting. Its likely that the variable resistor used in the experiment was defective. This would have caused the mass of the cathodes after electrolysis to be either higher or lower than the result we should have obtained due to having an inconstant current but mostly lower due to the longer time spent at 0 reading. This would have affected the results significantly, as it happened in a majority of the trials. This error could be avoided by testing for functionality the resistor before performing the experiment.
Distance between the electrodes was not preserved throughout the experiment. In theory,
the closer the distance between the electrode pairs, the quicker the
rate of electrolysis as the current would flows through the electrodes faster, therefore increasing the rate of electrolysis, consequently, increasing the mass of Cu deposited and vice versa. This could have been avoided by placing electrode pairs at polar ends of the beaker.
In this experiment, not all the copper ions are guaranteed stick to the cathode, especially when using high currents as the reaction to form Cu is happening too fast to allow adhesion. Instead of bonds being formed on the cathode, bonds are being formed with other copper being formed so adhesion to the cathode becomes poor. Therefore, when the copper strips were removed between tests, some of the Cu was seen at the end of the beaker, this happened mostly at higher currents. This would have also contributed to reducing our measurement for the mass gained.
Copper sulphate solution was not replaced after each test. This could have led to the solution having excess amounts of copper in the solution, which, may have inhibited electrolysis and may well have had an effect on the amount of copper deposited on the cathode.
Systematic errors
Systematic errors are errors in calibration of instruments/chemicals or distortions in the system that occur throughout the whole experiment. Low accuracy but high precision could indicate the presence systematic error.
One systematic error that may have been present includes the washing of the cathodes with distilled water before drying and weighing. Washing the cathodes may have actually removed the copper that was deposited on the cathodes especially those formed on high current which were very weakly attached to the cathode and could have fallen off easily. This could account for the loss in mass over the trials. This was done for all cathodes and this is why this is a systematic error.
Systematic error that could have also occurred in this experiment in things such as the electronic balance not being correctly configured, giving false readings throughout all measurements and thus reducing the precision of the results by systematically gaining or decreasing mass.
The multimeter may also have a systematic error present such as an error in the configuration of settings on the meter or that the meter was actually giving inaccurate readings. This could have affected the results greatly as this experiment depends on the quantity of current used during electrolysis. It is important to test all measuring equipment for correct calibration in a trial run before performing the experiment and this should have been done
Extension and improvement
To reduce the scatter and improve accuracy of the results it would be possible to increase the number of sample electrolytic cells in each trial. Also, new electrodes and a new lot of Copper (II) sulphate solution could be used for each trial.
A Possible Extension to this experiment would be to investigate the relationship between the mass of copper removed from the anode and the current quantity. From observing the half reactions it can be shown that at the cathode, Cu is formed, and on the anode, Cu2+ is formed proportionally and falls into the solution. According to faradays law, quantity of current is directly proportional to mass deposited at the cathode. As the mass lost from the anode should be equal to the mass deposited on the cathode, therefore, the magnitude of mass lost at the anode should also be proportional to the current quantity. This experiment design has an advantage in that mass of copper gained at the cathode could be misplaced and fall to the bottom of the solution or wash off when cleaning and drying, while measuring the mass lost at the anode is much more effective as nothing will fall off from the anode that was not already converted into Cu2+. Also, one could test time instead of current
Conclusion
The group was not satisfied with the results obtained in the experiment. The experimentally determined mass was not accurate to the theoretical mass nor did it prove the qualitative prediction of a directly proportional relationship between the two variables. the determined mass was larger than the theoretical mass by a factor of 10 for the first trial at 0.2 amps, this was ruled out as an outlier while, all the other results after than remained at an almost constant rate of gain in mass of approximately 0.008 grams throughout changing the current from 0.4 amps to 1.2 amps. Strangely enough, this value of mass gain was the expected theoretical mass value for a current of of 0.02 amps. Either the variable resistor did not change current from 0.2 amps, or mass was lost from washing and drying the cathode or some other reason. The experiment contained too many errors and will need to be performed again to gain a reliable result
Bibliography
Electrolysis of copper(ii) sulfate solution | nuffield foundation. 2012. [online] available at: . [accessed 10 june 2012].
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Electrolysis of copper sulfate and its use to refine blister copper by electrolytic refining. - youtube .[online] available at:. [accessed 10 june 2012].
