Introduction
Electrochemistry is the science that studies this union of chemistry and electricity. Batteries and fuel cells utilize spontaneous redox processes to convert chemical energy into electrical energy.
Oxidation reactions (or redox reactions) are an important part of chemical reactions. They involve the transfer of electrons from one source to the other via oxidation and reduction. When this process occurs spontaneously, forming products that are in a lower energy state than the reactants, the excess energy is released to the surroundings, frequently in the form of heat. Combustion in your car’s engine, “burning” calories in the gym and the rusting of iron are some examples of exothermic reactions. When the oxidation and reduction processes are physically separated in an electrochemical cell, the electrons are transferred through a wire connecting the cells and an electrical current can either be generated or used to drive the reaction. Add conclusion?
Electrochemistry is the science that studies this union of chemistry and electricity. Batteries and fuel cells utilize spontaneous redox processes to convert chemical energy into electrical energy. On the other hand, electrical energy can be used to drive non-spontaneous processes, converting the electrical energy into chemical energy that is stored in the reaction products.
Electrochemical Cells
Electrochemical cells fall into two broad categories. Voltaic (or galvanic) cells produce electricity from spontaneous redox processes. Batteries are a common example of this type of cell. Cells that use electricity to drive non-spontaneous reactions are called electrolytic cells. The basic components of an electrochemical cell are:
Two compartments separated by a “salt bridge” through which ions can flow. Oxidation occurs in the anode compartment and reduction occurs in the cathode compartment.
Two solid electrodes that are connected by a wire. The electrodes themselves don’t necessarily participate in the reaction.
Two solutions of electrolytes into which the electrodes are immersed. The ions of the electrolytes may participate in the reaction or they may be inert electrolytes that are present to carry charge.
Different metals have different tendencies to undergo oxidation, or lose electrons. Likewise, their cations have different tendencies to undergo reduction. This tendency is measured in terms of the metal cation’s reduction potential. The cell potential for a given electrochemical cell is the difference between the tendencies of the metal cations in their respective half-cells to undergo reduction. In a voltaic cell, the substance with the highest (most positive, or least negative) reduction potential will undergo reduction and the metal in the other compartment will be oxidized. The cell potential, Ecell, represents the difference between the tendencies of the metal ions to undergo reduction. For the reaction to be spontaneous, the overall cell potential must be positive.
Aim
The aim of the experiment was to see which of the three factors affects electrochemical cells. The three factors, Surface area, Concentration and Temperature. Each of these factors will be explored to see how they affect the current generated by the cell.
Hypothesis
Electrochemical cells are different from normal reactions; however the amount of volts produced by the cell will be effected by the factors that affect chemical reaction rates. The main factors that affect chemical reactions are Temperature, Surface area and Concentration. By increasing the amount of collisions per second the amount of volts should also be affected by the increased collision rates too.
Planning and Preliminary trials
A panel volt meter was given, however it was quite hard to read so it was swapped with an electronic volt meter that would give a much more accurate reading.
Since the beakers lips where curved the electrode would be at an angle, to fix this, the electrodes where dangled from a utility stand.
When the concentrations were tested it was noticed strange results, like how 0.5M KNO3 would be better than 1M, and it was found that the surface area of the salt bridge was different, therefore surface area tests had to be conducted and then be cut out the same sizes for every experiment after.
The surface area of the anodes and cathodes did not make any difference in the test, however they should make a difference if the battery is of a larger capacity (See figure 2.1 for example)
Figure 2.1
The heat capacity of the solutions where different, so only a rough estimate could be taken (eg. CuSO4 is at 59.4°C while ZnSO4 is at 62.4°C)
Theory: CuSO4 98.53 J/(mol–1 K)
ZnSO4 116.0 J/(mol–1 K)
The copper sulphate will heat up faster than the zinc sulphate, which means that the copper sulphate needs to be taken off the hot plate before the Zinc or otherwise the copper has a much higher temperature compared to zinc.
Materials
100mL beakers
Strips of copper, zinc, magnesium and iron (different sizes)
200mL x 0.5 mol zinc sulfate solution
400mL x 1 mol zinc sulfate solution
200mL x 2 mol zinc sulfate solution
200mL x 0.5 mol copper sulfate solution
400mL x 1 mol copper sulfate solution
200mL x 2 mol copper sulfate solution
200mL x 0.5 mol magnesium sulfate solution
400mL x 1 mol magnesium sulfate solution
200mL x 2 mol magnesium sulfate solution
200mL x 0.5 mol iron sulfate solution
400mL x 1 mol iron sulfate solution
200mL x 2 mol iron sulfate solution
100mL x 1 mol potassium nitrate solution
100mL x 2 mol potassium nitrate solution
100mL x 1 mol potassium sulfate solution
100mL x 2 mol potassium sulfate solution
200mL x 1 mol aluminum sulfate solution
Aluminum
Filter paper
Measuring cylinders
Volt Meters
Alligator clips
Conductivity Meter
Electric Thermometer
Large container (Enough to hold 2 x 250ml beakers)
Goggles
Ice
Steel wool
Method
Salt bridge solution test
Use steel wool or sandpaper to polish the metal strips, wash with distilled water after and wipe dry with a towel.
Place 60 mL of the 1 M ZnSO4solution in a 50-mL beaker. Place a strip of polished zinc in the beaker.
Place 60 mL of the 1 M CuSO4solution in a 50-mL beaker. Place a strip of polished copper in the beaker.
Connect alligator clip probes to a DC voltmeter. Connect the clips to the metal strips.
Cut a strip of filter paper and soak it in 1 M KNO3 solution and slowly lower it so both sides of the filter paper touches the contents of both beakers. Measure the reading on the volt meter.
Repeat the step 5 but use different concentrations of KNO3 and K2SO4
Repeat steps 1-6 except use Fe(II)SO4 and MgSO4 instead of ZnSO4 and CuSO4
The following experiments will be done with the salt bridge that gave the best result during the above experiment.
Concentration of solution test
Place 60 mL of the 1 M ZnSO4solution in a 50-mL beaker. Place a strip of polished zinc in the beaker (take down the temperature of both the ZnSO4 and CuSO4 solution for the next part of the experiment).
Place 60 mL of the 1 M CuSO4solution in a 50-mL beaker. Place a strip of polished copper in the beaker.
Connect alligator clip probes to a DC voltmeter. Connect the clips to the metal strips.
Cut a strip of filter paper and soak it in KNO3 solution and slowly lower it so both sides of the filter paper touches the contents of both beakers. Measure the reading on the volt meter.
Repeat the following steps with 0.25 and 0.5 M concentrations of CuSO4and ZnSO4.
Polish the iron and magnesium strips with steel wool.
Place 60 mL of the 1 M Fe(II)SO4solution in a 50-mL beaker. Place a strip of polished iron in the beaker.
Place 60 mL of the 1 M MgSO4solution in a 50-mL beaker. Place a strip of polished magnesium in the beaker.
Connect alligator clip probes to a DC voltmeter. Connect the clips to the metal strips.
Cut a strip of filter paper and soak it in KNO3 solution and slowly lower it so both sides of the filter paper touches the contents of both beakers. Measure the reading on the volt meter.
Repeat the following steps with 0.25 and 0.5 M concentrations.
Temperature of solution test
Place two 50mL beakers into an ice-cream container
Pour 60 mL of the 1 M ZnSO4solution into one 50-mL beaker. Place a strip of polished zinc in the beaker.
Pour 60 mL of the 1 M CuSO4solution into the other 50-mL beaker. Place a strip of polished copper in the beaker.
Fill the ice-cream container with ice and then fill it up with water, wait till the temperature of both solutions becomes steady then continue with the following steps.
Connect alligator clip probes to a DC voltmeter. Connect the clips to the metal strips.
Cut a strip of filter paper and soak it in KNO3 solution and slowly lower it so both sides of the filter paper touch the contents of both beakers equally. Measure the reading on the volt meter.
Repeat except place the beakers on a hot plate instead of an ice cream container.
Surface area of salt bridge test
Cut strips of filter paper at different sizes (1cm x 6.25cm, 2cm x 6.25cm and 3cmx6.25cm)
Pour 60 mL of the 1 M ZnSO4solution into one 50-mL beaker. Place a strip of polished zinc in the beaker.
Pour 60 mL of the 1 M CuSO4solution into the other 50-mL beaker. Place a strip of polished copper in the beaker.
Connect alligator clip probes to a DC voltmeter. Connect the clips to the metal strips.
Carefully lower the salt bridge between the beakers making sure it is evenly placed in the middle.
Measure the voltage produced and repeat with the different salt bridges.
Experimental Results
Table 1. Salt Bridge Solutions
Solution
Redox Reaction
Voltage
1M K2SO4
Zn(s) + Cu2+Zn2+ + Cu(s)
1.06V
0.5M K2SO4
1.03V
1M KNO3
1.03V
0.5M KNO3
1.01V
1M K2SO4
Mg(s) + Fe2+Mg2+ + Fe(s)
2.06V
0.5M K2SO4
2.04V
1M KNO3
2.08V
0.5M KNO3
2.05V
Table 2. Surface Area of Salt Bridge
Surface Area
Redox Reaction
Average Voltage
6.25cm2
Zn(s) + Cu2+Zn2+ + Cu(s)
1.06V
12.5cm2
1.11V
18.75cm2
1.12V
6.25cm2
Mg(s) + Fe2+Mg2+ + Fe(s)
2.07V
12.5cm2
2.13V
18.75cm2
2.16V
Table 3. Concentration
Concentration
Redox Reaction
Average Voltage
1M
Zn(s) + Cu2+Zn2+ + Cu(s)
1.09V
.5M
1.07V
.25M
1.06V
1M
Mg(s) + Fe2+Mg2+ + Fe(s)
2.10V
.5M
2.08V
.25M
2.06V
0.5M
2Al + 3Cu2+2Al3+ + 3Cu(s)
1.95V
0.25M
1.93V
0.125M
1.92V
Table 4. Temperature
Temperature
Redox Reaction
Average Voltage
16°C
Zn(s) + Cu2+Zn2+ + Cu(s)
24°C
60°C
16°C
Mg(s) + Fe2+Mg2+ + Fe(s)
24°C
60°C
16.7°C and 18.3°C
2Al + 3Cu2+2Al3+ + 3Cu(s)
0.39V
24°C both
0.5V
51.3°C and 48.6°C
0.57V
Table 5. Conductivity of Salt Bridge
Solution
Seimens (µs/m)
Higher is better
1M K2SO4
3839
0.5M K2SO4
3831
1M KNO3
3865
0.5M KNO3
3861
Discussion
To theoretically calculate the amount of voltage produced by each cell, the theoretical standard potential of the half cells need to be found.
The standard potential for the chemicals used in this experiment are:
Oxidants ⇌ Reductants
E°(V)
Cu2+ + 2e– ⇌Cu(s)
0.34
Fe2+ + 2e– ⇌Fe(s)
-0.41
Zn2+ + 2e– ⇌Zn(s)
-0.76
Al3+ + 3e– ⇌Al(s)
-1.71
Mg2+ + 2e– ⇌Mg(s)
-2.38
These values are when the cell is at STP. Source: Text book
To get the cell potential at STP:
Zn(s) + Cu2+Zn2+ + Cu(s)
Cu2+ + 2e– ⇌Cu(s) = 0.34V
Zn2+ + 2e– ⇌Zn(s) = -0.76
E°oxidation of Zn= – (-0.76 V) = + 0.76 V
E°cell= E°reduction+ E°oxidation
E°cell= 0.34 + 0.76
E°cell= +1.10 V
For all standard cell potential calculations refer to appendix 1.1
When the cell isn’t at STP the Nernst Equation has to be used.
Ecellis the cell potential E°cellrefers to standard cell potential R is thegas constant (8.3145 J/mol·K) T is theabsolute temperature n is the number of moles of electrons transferred by the cell’s reaction F is Faraday’s constant (96485.337 C/mol)
The Nernst equation was designed to be used when the values or environment was not at STP. The Nernst equation however does not support what happens if temperature changes but the concentration values are both equal. See example below.
Even though the temperature is at 350K the voltage will not change since log of 1 will be 0 and it cancels out RT/nF.
0.5M Aluminum Sulfate + 1M Copper Sulphate cell at 24.8°C
For all calculations containing the Nernst equation refer to appendix 1.2
The evidence backs up the hypothesis that states “the factors that affect reaction rate will also affect electrochemical cells” The voltages does increase and decrease in all 4 test:
As surface area increases so does the voltage for both cells.
As concentration increases so does the voltage for all 3 types of cells.
As temperature increases so does the voltage for all 3 types of cells.
The only downside is how the change in tests 2 and 3 are very small and almost negligible.
However the experimental results obtained did not match up to the theoretical results; but they are very similar and there might be a few reasons on why:
The resistance of the copper wire
The alligator clips don’t provide a lot of contact (decrease in possible surface area)
The temperature of the substances could be different since only room temperature was taken not the temperature of the actual solution.
The salt bridge could’ve been too small and stemmed the flow of electrons that could flow through it.
In the “Salt Bridge Solutions” experiment the results proved inconclusive so a conductivity meter was used to test the conductivity of each substance.
KNO3 will make a better salt bridge due to its higher conductivity rate; this is because of its solubility level which is at 316 g/Lat 20°C compared to K2SO4 solubility which is only 111g/L at 20°C.
Surface area test:
The tests were successful and backed up the hypothesis, however further tests were not taken since the beakers were too small and it would be a waste of material when the desired voltage has already been achieved.
It can be seen how the change in voltage is decreasing as the surface area increases which means that it will probably flatten out soon and there will not be a change at all. (See graph 1 and 2 in results)
Concentration:
For all the experiments on the 3 electrochemical cells the result backs up the hypothesis however concentration does not provide as much impact on voltage (See table 3 for results). The electrons are already there and making contact with the plate since it doesn’t really matter about which angle it hits the electrode nor how much power it hits with. It is hypothesized that by increasing concentration you increase the amount of electrons and that should be directly proportional to how long the cell would last; if you have double the amount of electrons but you are using them at the same rate that means that it should last twice as long since you have twice as many electrons.
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Temperature
There are no big changes in voltage when temperature is changed, like stated before “it doesn’t really matter about which angle it hits the electrode nor how much power it hits with”. Changing the temperature does two things, increases the rate at which the atom hits per second and changes how much force it hits with by providing it kinetic energy. However if a copper atom hits an electrode stronger, the KE doesn’t transfer over to the electron due to its size, also the electron will be transferred at the same rate since it has to go through a copper wire before it can get over to the cathode. Secondly if you hit an object more even though it can’t carry any more electrons won’t do anything, just like how a reaction is limited by its limiting reagent.
Percentage Error
P.E of 1M Copper and Zinc Cell
P.E =
P.E =
P.E = 0.00909%
See all Percentage Error calculations in Appendix 1.3
Conclusion
By conducting various experiments to see which of the three factors affects electrochemical cells. The three factors being Surface area, Concentration and Temperature. The results obtained were very close to the accepted value and the average percentage error was only XXX%.
The results obtained demonstrates that the hypothesis was indeed correct in stating that the factors that effect rate of reaction will also effect electrochemical cells, as stated by the hypothesis “by increasing the amount of collisions per second the amount of volts should also be affected by the increased collision rates too.” It should also be possible to gather further data to back up the hypothesis however the timeframe did not allow it.
Nevertheless even in ideal situation and also the simplicity of the experiment, the results should not have changed. In the 1M Copper and Zinc cell the voltage should’ve been 1.10V however 1.09V was the output which leads to further questions as to why it wasn’t exact. The reaction remains the same and energy was not lost since neither heat nor sound was produced during the experiment. It is assumed that unless at STP then the voltage can only be determined to ±0.02V (based on other experiments) For a more precise measurement a larger experiment needs to be conducted to a higher level where the experiment cannot be subjected to human error.
Appendix
Zn(s) + Cu2+Zn2+ + Cu(s)
Cu2+ + 2e– ⇌Cu(s) = 0.34V
Zn2+ + 2e– ⇌Zn(s) = -0.76
E°oxidation of Zn= – (-0.76 V) = + 0.76 V
E°cell= E°reduction+ E°oxidation
E°cell= 0.34 + 0.76
E°cell= +1.10 V
_____________________________________
2Al(s) + 3Cu2+2Al3+ + 3Cu(s)
Cu2+ + 2e– ⇌Cu(s) = 0.34V
Al3+ + 3e– ⇌Zn(s) = -1.71
E°oxidation of Al= – (-1.71 V) = + 1.71 V
E°cell= E°reduction+ E°oxidation
E°cell= 0.34 + 1.71
E°cell= +2.05 V
______________________________________
Mg(s) + Fe2+Mg2+ + Fe(s)
Fe2+ + 2e– ⇌Fe(s) = -0.41V
Mg2+ + 2e– ⇌Mg(s) = -2.38
E°oxidation of Zn= – (-0.76 V) = + 0.76 V
E°cell= E°reduction+ E°oxidation
E°cell= 0.34 + 0.76
E°cell= +1.10 V
References
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