Electrochemical Cells |
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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 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:
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.
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.
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.
Figure 2.1
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.
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
Table 1. Salt Bridge Solutions
Solution |
Redox Reaction |
Voltage |
1M K2SO4 |
Zn(s) + Cu2+” src=”https://s3-eu-west-1.amazonaws.com/aaimagestore/essays/1048788.002.png”>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+” src=”https://s3-eu-west-1.amazonaws.com/aaimagestore/essays/1048788.002.png”>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+” src=”https://s3-eu-west-1.amazonaws.com/aaimagestore/essays/1048788.002.png”>Mg2+ + Fe(s) |
2.10V |
.5M |
2.08V |
|
.25M |
2.06V |
|
0.5M |
2Al + 3Cu2+” src=”https://s3-eu-west-1.amazonaws.com/aaimagestore/essays/1048788.002.png”>Zn2+ + Cu(s) |
|
24°C |
||
60°C |
||
16°C |
Mg(s) + Fe2+” src=”https://s3-eu-west-1.amazonaws.com/aaimagestore/essays/1048788.002.png”>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 |
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+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 isFaraday’s constant96485.337 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:
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:
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.
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)
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.
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.
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
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+” src=”https://s3-eu-west-1.amazonaws.com/aaimagestore/essays/1048788.002.png”>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+
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