How duration affects the rate of electrolysis in a Voltaic Cell Essay
How duration affects the rate of electrolysis in a Voltaic Cell
Design and Conduct an experiment to investigate the effect of ONE FACTOR on redox reactions.
Introduction:
The two main components of redox reactions are reduction and oxidation. Reduction is a gain in electrons and the decrease in oxidation number whereas oxidation is the loss of electrons and the increase in oxidation number. Voltaic cells, also known as galvanic cells generate their own electricity. The redox reaction in a Voltaic cell is a spontaneous reaction. For this reason, voltaic cells are commonly used as batteries. Voltaic cell reactions supply energy which is used to perform work.
The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus (known as the salt bridge which primarily completes a circuit and maintains electrical neutrality) that allows electrons to flow. The functions of a voltaic cell are quite simple. There happens to be an anode and a cathode. The positive ions go the negative electrode (anode) whereas the negative ions go to the positive electrode (cathode). Electrons always flow from the anode (where oxidation takes place) to the cathode (where reduction takes place). Electrons flow across wires whereas ions flow across the electrolyte and the salt bridge.
Aim:
The objective of this experiment is to see how the time affects the mass of the zinc electrode (anode) and the copper electrode (cathode) in a voltaic cell.
Variables:
Variable
Type of variable
How it will be controlled
Time (s)
Independent (The one you change)
Values from 5 to 35 minutes will be used
Mass of anode & cathode (g)
Dependent (The one you measure)
Electrodes will be measured after each time interval
Current (A)
Controlled
Measure the current with the help on an ammeter
Initial mass of cathode and anode (g)
Controlled
Weigh out the electrodes using top pan balance from the beginning of the experiment
Charge on ion
Controlled
Use the same solution for all the trials. The charge on the copper ion should be 2+ since the copper 2+ is being converted to copper metal. The charge on the zinc ion should be 0 because Zn is being converted to Zn 2+
Concentration of electrolyte
Controlled
Use the same solution for all the trials. The solution primarily should be 1 mol dm3 (just like standard conditions)
Area of electrodes (cm2)
Controlled
Measure the electrodes to ensure they have the same dimensions (9×2.5cm). Use the same electrodes for all the trials.
Volume of electrolyte (cm3)
Controlled
Use a measuring cylinder to measure out the electrolyte’s volume
Atmosphere which we are working under
Controlled
Primarily we are working under standard room temperature of 298 K
Apparatus:
* 1×22.5cm2 copper electrode
* 1×22.5cm2 zinc electrode
* 100cm3 1mol dm3 Zinc sulphate solution
* 100cm3 1mol dm3 copper (II) sulphate solution
* Filter paper (required to create a salt bridge)
* 100cm3 of potassium nitrate solution (the spectator ion which I will require for creating the salt bridge which will complete the circuit and maintain electrical neutrality)
* 2x200cm3 beakers
* Stopwatch (ï¿½0.01s)
* 1x100cm3 measuring cylinder (ï¿½1.0cm3)
* Voltmeter
* 2 connecting wires
* Top pan balance (ï¿½0.01g)
Method:
1) Set up the voltaic cell. Use a measuring cylinder to measure out 100cm3 of copper sulphate solution. Pour it into the 200 cmï¿½ beaker.
2) Next do the same for zinc sulphate. Use a measuring cylinder to help measure out 100cm3 of zinc sulphate solution. Pour it into a different 200 cmï¿½ beaker.
3) Weigh the masses of the electrodes separately using a top pan balance. Record the initial masses.
4) Connect the wires to the outlets in the zinc and copper electrode. Place them in the corresponding outlets of the voltmeter.
5) After that we cut out some filter paper and dip that into our spectator ion (potassium nitrate) in order to build a salt bridge. The salt bridge will primarily complete the circuit, allow flow of ions and maintain electrical neutrality. The salt bridge will be placed in such a way that the ends of the salt bridge will be touching separate solutions of zinc sulphate and copper sulphate. The overall circuit should resemble the diagram in Figure.1.
6) Place the zinc electrode into the beaker with the zinc sulphate solution and the copper electrode into the beaker with the copper sulphate solution and at the same time, start the stopwatch. Keep the stopwatch running until 200 seconds elapse. *Note we will be recording the time every 5 minutes because 1 or 2 minutes simply isn’t enough for the change to take place
7) Take the cathode out of the solution and measure its mass (remember, before doing so, shake it a couple of times in order to remove any moisture). Record the mass. Do the same for the zinc electrode
8) Place the electrodes into their respective solutions once again and start timing. Repeat steps 5 to 6
9) Repeat the same steps until we get mass readings for up to 60 minutes of experimenting.
Data Collection and Processing
Raw data:
– Initial mass of anode (zinc electrode): 31.29 ï¿½0.01g
– Initial mass of cathode (copper electrode): 32.05 ï¿½0.01g
Table 1 – Mass of anode and cathode obtained from different time intervals
Duration of electrolysis (ï¿½0.21s)
Mass of anode (zinc electrode) (ï¿½0.01g)
Mass of cathode (copper electrode) (ï¿½0.01g)
300.00 (5 minutes)
31.27
32.08
600.00 (10 minutes)
31.14
32.16
900.00 (15 minutes)
31.08
32.27
1200.00 (20 minutes)
31.00
32.42
1500.00 (25 minutes)
30.83
32.49
1800.00 (30 minutes)
30.61
32.80
2100.00 (35 minutes)
30.25
33.08
Qualitative observations:
– We can see that the copper is deposited at the cathode where the cathode begins to get more pink/ brownish colour.
– Blue colour of copper sulphate solution begins to get paler.
– Zinc electrode begins to corrode a bit. Most corrosion can be observed at 35 minutes time interval.
Note*
– Uncertainties:
The average reaction time was ï¿½0.5s even though it did alter from interval to interval. Note that there is also a ï¿½0.01s time uncertainty in the stopwatch itself. The uncertainty for mass is inscribed on the top pan balance as well.
Data Processing:
We must now calculate the mass changes which have taken place due to experimenting with different time intervals. (Different time intervals would result in a different mass change)
This can be calculated simply by doing the following:
Mass change = final mass – initial mass
Due note however that this formula can only be used for calculating the mass change taking place at the cathode (copper electrode where reduction takes place). This is because copper 2+ is being converted to copper metal and is being deposited at the cathode. Obviously this would result in a mass gain at the cathode. Therefore, it would be better for us to use the formula ‘Mass change = final mass – initial mass’ so that it gives us a positive value for the mass change taking place at the cathode.
Example 1
Mass change = final mass – initial mass
=> 32.08 – 32.05
=> 0.03g
Example 2
Now to calculate the mass change taking place at the anode (zinc electrode), we use the following formula, Mass change = initial mass final mass. In this case we use this formula because we know that the zinc is being oxidized to zinc 2+ leading the zinc electrode to corrode. This therefore results in a decrease in mass of the anode (zinc electrode). Thus, it would be better for us to use the formula ‘Mass change = initial mass – final mass’ so that it gives us a positive value for the mass change taking place at the anode.
Mass change = initial mass – final mass
= > 31.29 – 31.27
= > 0.02
Table 2 Mass changes of anode and cathode for each time interval
Time (ï¿½0.21s)
Mass change of Anode (Zinc electrode)(ï¿½0.01g)
Mass change of cathode (copper electrode) (ï¿½0.01g)
300.00 (5 minutes)
0.02
0.03
600.00 (10 minutes)
0.15
0.11
900.00 (15 minutes)
0.21
0.22
1200.00 (20 minutes)
0.29
0.37
1500.00 (25 minutes)
0.46
0.44
1800.00 (30 minutes)
0.68
0.75
2100.00 (35 minutes)
1.04
1.03
Graph 1:
Graph 2:
To derive the equation for the two separate reactions, the number of electrons gained or lost during the process has to be deduced.
The mass change per minute can be deduced from the gradient. Therefore we first calculate the gradient of graph 1 (mass changes for zinc electrode). For calculating the gradient, find two points which perfectly fits in the grid. In this case, the points (0.04. 100) and (0.08, 200)
Gradient= (Y2 – Y1) ï¿½ (X2 – X1)
= (0.08 0.04) ï¿½ (200 – 100)
= (0.04) ï¿½ (100)
= 0.0004
Therefore, the gradient of the first graph is 0.0002. So the mass change per minute for the anode is 0.0004.
Next, we calculate the gradient of graph 2 (mass changes for copper electrode). To find the gradient, we work with the points (0.20. 500) and (0.24, 700)
Gradient= (Y2 – Y1) ï¿½ (X2 – X1)
= (700 – 500) ï¿½ (0.24 0.20)
= (200) ï¿½ (0.04)
= 0.0002
Therefore, the gradient of the first graph is 0.0002. So the mass change per minute for the cathode is 0.0002.
The uncertainties also need to be propagated through the summation of the fractional uncertainties.
Uncertainties regarding zinc electrode:
Fractional uncertainty of mass = absolute uncertainty ï¿½ actual value
= 0.01 ï¿½ 0.02
= 0.500
Fractional uncertainty of time = absolute uncertainty ï¿½ actual value
= 0.21 ï¿½ 300
= > 0.0007 = 0.001
Total uncertainty = 0.001 + 0.500 = 0.501 to 3 decimal places
Therefore the rate of change is 0.004 ï¿½ 0.501 g/s
Table 3 – Rate of change for each time interval for anode (zinc electrode)
Time (ï¿½0.21s)
Rate of change of anode (zinc electrode) (g/s)
60.00
0.004ï¿½0.501
120.00
0.004ï¿½0.067
180.00
0.004ï¿½0.048
240.00
0.004ï¿½0.035
300.00
0.004ï¿½0.022
360.00
0.004ï¿½0.015
420.00
0.004ï¿½0.001
To calculate the number of electrons in zinc electrode, the following equation may be used:
Number of electrons = molar mass ï¿½ mass of electrode (mass of one of the samples)
= 65.37 ï¿½ 31.27
= 2.09
Therefore, this would be the halfequation which would occur at the cathode:
Zn–> Zn2.09+ + 2.09e
Due to the loss in a bit more electrons compared to the theoretical formula, it would be a stronger reducing agent therefore the electrode potential would be lower (more negative) than that of the original value. Nevertheless, the electrode potential cannot be determined.
Uncertainties regarding copper electrode:
Fractional uncertainty of mass = absolute uncertainty ï¿½ actual value
= 0.01 ï¿½ 0.03
= 0.333
Fractional uncertainty of time = absolute uncertainty ï¿½ actual value
= 0.21 ï¿½ 300
= > 0.0007 = 0.001
Total uncertainty = 0.001 + 0.333= 0.334 to 3 decimal places
Therefore the rate of change is 0.002 ï¿½ 0.334 g/s
Table 3 – Rate of change for each time interval for cathode (copper electrode)
Time (ï¿½0.21s)
Rate of change of cathode (copper electrode) (g/s)
60.00
0.002ï¿½0.334
120.00
0.002ï¿½0.091
180.00
0.002ï¿½0.046
240.00
0.002ï¿½0.027
300.00
0.002ï¿½0.023
360.00
0.002ï¿½0.013
420.00
0.002ï¿½0.010
To calculate the number of electrons in copper electrode, the following equation may be used:
Number of electrons = molar mass ï¿½ mass of electrode (mass of one of the samples)
= 65.50 ï¿½ 32.08
= 2.04
Therefore, this would be the halfequation which would occur at the cathode:
Cu2.04+ + 2.04e –> Cu
Due to the gain of a bit more electrons compared to the theoretical formula, it would be a slightly weaker oxidizing agent therefore the electrode potential would be slightly lower than that of the original value. Nevertheless, the electrode potential cannot be determined.
Conclusion
My results show that as the duration/ time intervals increase, the mass of the anode (zinc electrode) decreases and the mass of the cathode (copper electrode) increases. We can see that there is a strong positive correlation between the time it takes for both electrodes to change in masses. If the duration is longer, then more electrons flow from the zinc electrode to the copper electrode (anode to cathode) through the electrical wires, while ions flow through the salt bridge to complete.
As we know, in a voltaic cell/ galvanic cell, oxidation occurs at the anode (negative electrode) where as reduction occurs at the cathode (positive electrode). Primarily, zinc is oxidized at the anode and converted to zinc 2+. This causes corrosion at the zinc electrode due to the metal being converted to ions thus the mass of the zinc electrode (anode) decreases. On the other hand, copper undergoes reduction at the cathode and the copper 2+ ions get converted to copper metal. This causes the copper metal to be deposited at the cathode thus leading to the copper electrode (cathode) to increase in mass as the duration is increased. The following anodic reaction takes place at the zinc electrode (this is the theoretical equation):
Zn (s) –> Zn2+ (aq) + 2e
However the equation we found experimentally is:
Zn–> Zn2.09+ + 2.09e
Hence, this suggests that since the former zinc sample has more electrons to lose, it is an even stronger oxidizing agent compared to the theoretical equation and is slightly higher in the electrochemical series than the latter zinc samples.
According to the results that have been gathered, there is a positive correlation between the time it takes to electrolyse an aqueous solution and the rate of electrolysis. The rate of electrolysis was measured using the mass of cathode. If the duration of electrolysis is longer, then more electrons will flow through the circuit and more ions will flow from the anode to the cathode. Oxidation occurs at the anode whereas reduction occurs at the cathode. The cathode gains electrons therefore the mass decreases. The following reaction has taken place (although this is the theoretical equation):
Cu2+ (aq) + 2e –> Cu (s)
However, the experimental equation is:
Cu1.75+ + 1.75e –> Cu
Therefore this implies that since the former copper sample has more electrons to gain, it is a stronger oxidizing agent and it is lower in the electrochemical series than the latter copper sample.
The value of the electrode potential hasn’t been calculated, however, the number of electrons is 25% off there that shows that there is a great difference between the literature value and the experimental value. According to the graph in the previous page, there is a very strong positive correlation between the mass change and duration of electrolysis as can be deduced from the high R squared value. The change in mass over a certain period of time is very gradual because of the size of the electrons. Although a lot of electrons are able to flow through the electrolyte, there is not such a drastic change. By looking at the graph, almost all the error bars for the points touch the line of best fit which means the data is fairly accurate.
The theoretical mass of a copper electrode would be 31.75g. From the results that have been tabulated, the mass of a copper electrode is 36.21g.
The percentage error can be calculated using the following formula:
Percentage error = difference x 100
theoretical value
= 4.46 x 100
31.75
= 14.04%
This shows that although there is not such a big difference between the theoretical value and the experimental value.
Evaluation
Limitation
Type of error
Improvement
The mass of the anode was not measured therefore the rate of electron transfer between the two electrodes could not be determined. This could have increased or decreased the mass of the cathode.
Random
Measure the mass of the anode
The power pack has internal resistance therefore not all the current was emitted. This could have decreased the current, thus decreasing the number of electrons produced.
Random
Use a resistor to accurately measure the current
The top pan balance had a zero offset error. This could have increased the mass of the cathode.
Systematic
Use the top pan balance with the 0.001 uncertainty to obtain more accurate values.
a
A

Subject: Chemistry,

University/College: University of California

Type of paper: Thesis/Dissertation Chapter

Date: 17 November 2017

Words:

Pages:
Let us write you a custom essay sample on How duration affects the rate of electrolysis in a Voltaic Cell
for only $16.38 $13.9/page