Electrolysis Of Copper Sulphate Chemistry Essay
Aim: to find out how the amount of current affect the amount of copper deposited at the cathode during the electrolysis of copper sulphate with inert graphite electrodes
Hypothesis: Michael Faraday’s first law of electrolysis states:
“The mass of any element deposited during electrolysis is directly proportional to the number of coulombs of electricity passed”
From this statement we can say that the amount of electricity that we pass in the circuit will directly affect the rate of electrolysis. Ions are moving towards the electrodes when electricity is passed through the circuit. If higher amounts of electricity is passed, the flow of ions will be faster therefore travelling faster to their respective electrodes. Ions moving faster will increase the rate of electrolysis.
My hypothesis based on faraday’s first law of electrolysis is that as current increases and more amps are passed through the circuit the rate of electrolysis will increase. The rate of electrolysis can be seen from the increase in mass of cathode.
Variables:
Independent
Dependent
Controlled
Current- this will be the only variable that will be changed because we are looking at the how current affects the mass of copper deposited at the cathode
Mass of copper- we believe that the amount of copper at the cathode will depend on the current. From this we can say that the mass will change according to the amount of current
Concentration of electrolyte (1 mole)- Increased concentration of ions increases the current and consequently the rate of electrolysis. Therefore because we only want to look at how current affects the mass deposited we should keep concentration constant
Time (5minutes)- time just like concentration affects the amount of copper deposited at the cathode. The longer we leave it for more copper will form.
Temperature- the temperature of the room should be monitored and recorded
Surface area of electrodes- Increased surface area of the electrodes used increases the current and the rate of electrolysis
Separation of electrodes – Decreased distance between the electrodes increases the rate of electrolysis
Quantity of copper sulphate solution- the amount of copper sulphate should remain constant to make it a fair test.
Equipment and materials
(100ml ±5) Beaker
Inert electrodes(graphite) [x2]
Electronic balance
Stop watch
Tile
Crocodile clips and Wires [x3]
Distilled water
Power supply
Ammeter( ± 0.05 A)
Variable resistor
Measuring cylinder
80ml of 1mol of copper sulphate
Card board
Sand paper
Thermometer (± 0.05 ËšC)
heater
Control of variables:
All of these variables will have an effect on the mass of copper deposition therefore these are the methods use to ensure they are keep constant:
variables
How will they be kept constant
concentration
The same concentration of the solution will be used
Quantity of copper sulphate
(80ml ±5) of copper sulphate will be measure out each time using the same measuring cylinder
current
experiment will be carried out with the same amount of current using the same ammeter when repeated
temperature
Temperature will be monitor throughout the experiment using the same thermometer, as we cannot keep the temperature in the laboratory throughout.
Diagram:
Crocodile clip
Method:
Set up the apparatus as illustrated
Clean one of the graphite electrodes, the cathode, with sandpaper.
Weight the electrode (cathode) on an electronic balance and record its mass in grams with its uncertainty. Make sure the electronic balance is accurate thus being able to detect small change. ( a suitable balance would have uncertainty ±0.01)
Cut out a piece of cardboard in a rectangular shape and make two holes into the cardboard. These holes should allow the electrodes to fit into them.. (Note that the distance between the electrodes can affect the rate of electrolysis)
After placing the electrodes through the holes, connect the clean electrode(cathode) to the negative terminal of the power supply using the wire that is attached to crocodile clip
Connect the other electrode (anode) to the positive terminal of the power supply using the other crocodile clip. (This electrode will act as the anode. It will not be necessary to clean the anode as it will not be weighed)
Measure out (80ml ± 5) of 1 mol copper sulphate solution using a measuring cylinder and pour it into a (100ml ±5) beaker
immerse the two electrodes 3cm into the copper sulphate solution (note that the depth of the electrodes affect the rate of electrolysis)
Before starting the experiment makes sure the circuit is working fine and use the variable resistor to adjust the current to 2.0A.
Turn on the power and at the same time start the timer
After 5 minutes(300 s) switch of the power supply and stop the timer
Carefully remove the electrodes from the copper sulphate solution
Gently dip the cathode in distilled water and dry it using a heater.
Remove the cathode from the cardboard and re-weight the cathode
Record its new mass and after, place it back in the cardboard.
Repeat the experiment again to increase accuracy.
Once you are satisfied repeat step 1-17 with a different current. (0.4, 0.6,0.8)
Theoretical data:
The quantity of electricity or charge in a current for a specified amount of time is:
Q = I x t
Q = quantity of electricity or charge in coulombs (C)
I = current in amps (A)
t = time (s)
The Faraday constant, F, is the quantity of electricity carried by one mole of electrons
F = Avogadro’s Number x charge on electron in coulombs
Avogadro’s number = 6.022 x 1023 mol-1
Charge on electron in coulombs = 1.602192 x 10-19 C
Therefore
F = 6.022 x 1023 mol-1 x 1.602192 x 10-19 C
F = 96,484 C mol-1
The quantity of electricity required to deposit an amount of metal can be calculated:
Q = n(e) x F
Q = quantity of electricity in coulombs (C)
n(e) = moles of electrons
F = Faraday constant = 96, 500 C mol-1
Mass of copper deposited from copper (II) sulphate solution using a current of (0.2 A ± 0.05) over 300 seconds:
quantity of electricity: Q = I x t
I = 0.20 A ± 0.05
t = 300 s ± 0.01
Q = (0.20) x (300) = (60.0) C
moles of electrons: n(e) = Q ÷ F
Q = (60.0) C
F = 96.5 x 10³ C mol-1
n(e) = (60.0) ÷ 96.5 x 10³ = (6.22 x 10-4) mol
mass of copper: mass = n x Molar mass
Cu2+ + 2e —–> Cu(s)
1 mole of copper is deposited from 2 moles electrons
n(Cu) = ½n(e) = ½ x (6.22 x 10-4)= (3.11 x 10-4) mol
Molar mass = 63.55 g mol-1
mass (Cu) = (3.11 x 10-4) x 63.55 = (1.98x 10-2) g
Mass of copper deposited from copper (II) sulphate solution using a current of (0.4 A ± 0.05) over 300 seconds
quantity of electricity: Q = I x t
I = 0.40 A ± 0.05
t = 300 s ± 0.01
Q = (0.40) x (300= (120) C
moles of electrons: n(e) = Q ÷ F
Q = (120 ± 15.00) C
F = 96.5 x 10³ C mol-1
n(e) =(120) ÷ 96.5 x 10³ = (1.24 x 10-3) mol
mass of copper: mass = n x Molar mass
Cu2+ + 2e —–> Cu(s)
1 mole of copper is deposited from 2 moles electrons
n(Cu) = ½n(e) = ½ x (1.24 x 10-3) = (6.22 x 10-4) mol
Molar mass = 63.55 g mol-1
mass (Cu) = (6.22 x 10-4) x 63.55 = (3.95 x 10-2) g
Mass of copper deposited from a copper (II) sulphate solution using a current of ( 0.6 A ± 0.01) over 300 seconds
quantity of electricity: Q = I x t
I = 0.60 A ± 0.05
t = 300 s ± 0.01
Q = (0.60) x (300) = (180) C
moles of electrons: n(e) = Q ÷ F
Q = (180) C
F = 96.5 x 10³ C mol-1
n(e) =(180) ÷ 96.5 x 10³ = (1.87 x 10-3) mol
mass of copper: mass = n x Molar mass
Cu2+ + 2e —–> Cu(s)
1 mole of copper is deposited from 2 moles electrons
n(Cu) = ½n(e) = ½ x (1.87 x 10-3) = (9.33 x 10-4) mol
Molar mass = 63.55 g mol-1
mass (Cu) = (9.33 x 10-4) x 63.55 = (5.93 x 10-2) g
Mass of copper deposited from a copper (II) sulphate solution using a current of (0.8 A ± 0.01) over 300 seconds
quantity of electricity: Q = I x t
I = 0.80 A ± 0.05
t = 300 s ± 0.01
Q = (0.80) x (300) = (240) C
moles of electrons: n(e) = Q ÷ F
Q = (240) C
F = 96.5 x 10³ C mol-1
n(e) = (240) ÷ 96.5 x 10³ = (2.49 x 10-3) mol
mass of copper: mass = n x Molar mass
Cu2+ + 2e —–> Cu(s)
1 mole of copper is deposited from 2 moles electrons
n(Cu) = ½n(e) = ½ x (2.49 x 10-3) = (1.24 x 10-3) mol
Molar mass = 63.55 g mol-1
mass (Cu) = (1.24 x 10-3) x 63.55 = (7.90 x 10-2) g
Data Collection
RAW DATA:
Trial 1
current (A)
±0.05
Initial mass of cathode (g) ±0.01
Final mass of cathode (g) ±0.01
0.2
4.62
4.65
0.4
4.65
4.75
0.6
4.75
4.84
0.8
4.84
4.96
Trial 2
current (A)
±0.05
Initial mass of cathode (g) ±0.01
Final mass of cathode (g) ±0.01
0.2
4.65
4.68
0.4
4.68
4.79
0.6
4.79
4.87
0.8
4.87
4.98
Trial 3
current (A)
±0.05
Initial mass of cathode (g) ±0.01
Final mass of cathode (g) ±0.01
Change in mass (g)
±0.01
0.2
4.64
4.69
0.00
0.4
4.69
4.71
0.00
0.6
4.71
4.79
0.20
0.8
4.85
4.97
Time (s)
600
1200
1800
2400
Temperature (ËšC) ±0.05
22
23
24
23
Observations: –
Bubbles slowly forming at the anode
Copper forming at cathode
Side reaction black particles settling at the bottom
Rapid Bubbling as current increase
Solution gets darker, eventually almost completely black
PROCESSED DATA:
Trial 1
current (A)
±0.05
Initial mass of cathode (g) ±0.01
Final mass of cathode (g) ±0.01
Change in mass (g)
±0.01
0.2
4.62
4.65
0.03
0.4
4.65
4.75
0.10
0.6
4.75
4.84
0.09
0.8
4.84
4.96
0.12
current (A)
±0.05
Initial mass of cathode (g) ±0.01
Final mass of cathode (g) ±0.01
Change in mass (g)
±0.01
0.2
4.65
4.68
0.03
0.4
4.68
4.79
0.11
0.6
4.79
4.87
0.08
0.8
4.87
4.98
0.11
current (A)
±0.05
Initial mass of cathode (g) ±0.01
Final mass of cathode (g) ±0.01
Change in mass (g)
±0.01
0.2
4.64
4.68
0.04
0.4
4.68
4.74
0.06
0.6
4.74
4.85
0.11
0.8
4.85
4.97
0.12
current (A)
±0.05
Average mass of cathode (g) ±0.01
0.2
0.03
0.4
0.09
0.6
0.09
0.8
0.12
Conclusion:
The result shows us that the rate of electrolysis increased as the current increased. This is shown by a faster increase in mass of the cathode. From the graph we can see that almost all our results pass through the line of origin indicating that current is directly proportional to the rate of electrolysis. This proves our hypothesis to be true. As the current is doubled so is the mass.
From observations we see bubbles of gas being formed at the cathode. This gas is oxygen.
Cathode reaction: Cu2+(aq) + 2e- → Cu(s)
Anode reaction: 2H2O(l) → O2(g) + 4H+(aq) + 4e-
Since we used graphite electrodes the oxygen at the anode reacted with the oxygen to form CO₃. We also find during the experiment as the current is increased particles of carbon settle at the bottom of the electrolyte. We also find that the electrolyte slowly starts to turn darker.
Evaluation:
The procedure that we used to collect the data was rather effective. However from the experiment that was carried out there was several factors that could have made the experiment much more successful and to some extent more accurate.
The result that we got from the experiment was somehow what we expected if we look at our theoretical graph. The change in mass was somehow greater than expected but it still did follow Faraday’s law. There were some results that did not follow Faraday’s law and did not pass through the line of origin. This was the average for the 0.4A. It did not fit along with the others.
The factors that could have reduced the accuracy of my experiment are:
The electrodes: we used graphite electrodes to perform this experiment. We had used a cardboard to place the two electrodes in, however at times the electrodes would come close together. The separation has a great effect on the rate of electrolysis. Decreasing the separation would increase the rate of electrolysis. We also find that when removing the electrodes to weigh some of the copper would stick to the cardboard, therefore the electrodes would have to be made completely dry before weighing. Rather than using a cardboard, to make the experiment more accurate we could use something that could easily remove the electrodes as well as keep the separation constant.
The variable resistor: Even with the use of a variable resistor the current constantly fluctuated. Not being able to keep the current constant can affect the rate of the reactions. Therefore a more accurate resistor could have been used; however this was not available at the time of the experiment.
Temperate: the temperature of the room had a small fluctuation. To make the experiment more accurate a water bath could have been used.
My results were fairly accurate; however the experiment could have been repeated several more times to increase its accuracy.
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