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In this experiment students have the opportunity to observe and measure the relative reduction potentials of various metals in order to have a better sense of comprehension of the topics associated with reaction spontaneity as well as reduction potentials. The objectives of this experiment were accomplished through the use of electrolysis and reduction potentials. Electrolysis is essentially when chemicals decompose through the addition of an electric current through some liquid or some solution that contains ions. The general method associated with part one is that metals and their corresponding metal solutions are placed on filter paper and then measured with a voltammeter in reference to either the copper or silver electrode. Part two just consists of placing aluminum in a CuCl2 solution and observing it dissolve. Part three just consists of placing copper in a AgNo3 solution and watching it dissolve. Part four is where the process of electrolysis comes in. In this experiment, the redox reaction was carried out through a galvanic cell where electroplating metals and water electrolysis produces hydrogen and oxygen, which require electrical energy to produce chemical change in transferring electrons from one substance to another. The galvanic cell was divided into an anode and cathode where the oxidation half reaction occurs at the anode and the reduction half reaction occurs at the cathode, and the electrons flow from anode to cathode.
We redid part one twice and for the most part we obtained accurate data, however there were several pieces of data that we acquired that were off. There is a 67.9 percent error for Zinc in reference to the silver electrode. A source of error for this could be that the solution on the metal had dried up or a student who had previously done this part had returned the metal to the wrong container, thus our recorded data would be off. Another source of error associated with part one can be how the electrodes could have been not fully cleaned before moving on to the next measurement. This remaining residue as well as some electron transfer could have possibly made Zinc’s measured value to be more positive than it should actually be. The general trend associated with measured reduction potentials is that if the voltage in reference to the Hydrogen electrode is positive for a specific half reaction, then it will be positive for the Copper and Silver electrode as well. If the voltage is negative in reference to the Hydrogen electrode, the voltage will also be negative for the Copper and Silver electrode as well. This trend is also similar for the actual reduction potential. When comparing the actual and measured standard reduction potential values the percent errors for the half reactions in reference to the copper electrode are as follows: Ag 8.69%, Cu 0%, Fe 8%, and Zn 19.14%. In reference to the silver electrode they are as follows: Ag 0%, Cu 6.5%, Fe 18.4%, and Zn 67.9%.
Standard Reduction Potentials at 25°C
In part two of this experiment we observed the aluminum foil turn from a silverfish color to a brown color. Bubbles were also forming within it, and the test tube felt hot to the touch, thus we deduced that an exothermic reaction was happening. When heat is released, that means the reaction is exothermic and not endothermic. The solution also appeared to change color as well. The solution originally was clear, however as time passed, the solution had a bluish tint to it. It is evident that a chemical reaction occurred primarily because of how there was a color change, the aluminum foil changed in appearance, and because of how heat seemed to be released from the test tube. The overall balanced equation and corresponding cell potential is as follows:
Al(s) à Al3+ + 3e- E°=+1.70 V
Cu2+ + 2e- à Cu(s) E°= +0.34 V
2Al(s) + 3Cu2+ à 2Al3+ + 3Cu(s) E°cell= 2.04 V
In Part three of this experiment we observed the copper wire with an original orange-red color obtain a gray/silverfish color. It appeared that no heat was emitted from the reaction and seeing as how the test tube was initially cold, however at the end of the reaction it felt cooler, it led us to believe that an endothermic reaction could possibly be taking place. The solution that the copper wire was submerged in appeared blue in color as several minutes passed from when the copper wire was initially placed in the AgNO3 solution. It is evident that a chemical reaction took place primarily because of how there was a color change, a temperature change, and how the copper wire changed in appearance. The overall balanced equation and corresponding cell potential is as follows: Ag+ + e- à Ag(s) E°= 0.80 V
Cu(s) à Cu2+ + 2e- E°= -0.34 V
2Ag+ + Cu(s) à 2Ag(s) + Cu2+ E°cell= 0.46 V
In part four of the experiment, the half reactions that took place at the anode and cathode are as follows: Anode half reaction: Cu(s) → Cu2+ + 2 e-
Cathode half reaction: 2H+ + 2e- → H2 (s)
Regarding method one of determining Faraday’s constant and Avogadro’s number found via collecting H2, the values we obtained for Faraday’s constant is 115711.2C/mol of electrons. The value we obtained for Avogadro’s number via the first method is 7.22×10
23 molecules/mol. In comparison to the actual value of Avogadro’s number (6.02214× 1023) and Faraday’s constant(96845C/mol), the experimental values both have a 19.9 percent error. This percent errors could be associated with how long part four took. The current measured by the multimeter appeared to be lower than what previous classes had recorded, and this may have had an impact on the numerical values we obtained, thus impacting the final Avogadro’s number and faraday’s constant that we acquired. Regarding the second method of determining Faraday’s constant and Avogadro’s number found from Cu reacted, the values we obtained for Faraday’s constant is 111925.75C/mol of electrons. The value we obtained for Avogadro’s number via the second method is 6.98×10
23 molecules/mol. In comparison to the actual value of Avogadro’s number (6.02214× 1023) and Faraday’s constant(96845C/mol), the experimental values both have a 16.0 percent error. The percent error for the first method appears to be higher than for the second method. A source of error associated with the first method’s percent error value being larger than the second methods could deal with the water displacement that took place. The water displacement values are integral for obtaining the proper numerical values to use for the calculations, if even a small quantity of H2 leaked out or if the copper wire was moved even a tiny amount, then that could have impacted the amount of H2 gas that was being collected within the burette.
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