The purpose of this lab was to measure the reduction potentials of metals and to understand how a reduction-oxidation relationship works to produce spontaneity. Lastly, these results were used to determine the Faraday’s constant and Avogadro’s number by electrolysis, which is chemical decomposition produced by passing an electrical current through a solution with ions. The primary objective of part one of this experiment was to discover how the properties of an electrochemical cell works.
An electrochemical cell is based on an oxidation-reduction reaction and is composed of two half reactions: an anode and a cathode half reaction. Oxidation typically occur at the anode and reduction at the cathode. The electrochemical cell produces an electrical current which is driven by the potential differences in the two half reactions. The reduction potential of hydrogen was given and the reduction potential of copper and silver was able to be measured in reference to the hydrogen electron. It was predicted that everything below the referenced should be decreasing and anything above should increase.
The reasoning behind this is that if the meter reading is negative, that means the positive lead is connected to a metal that has a more negative potential than the metal connected to the negative lead wire. The predictions seemed to correspond correctly with the measurements. For copper, everything below the referenced 0. 0 V was decreasing below it [-0. 63,-0. 78] and increasing above it [0. 42]. As was the case with the silver electrode. During part two, the half reactions were provided in order to determine if Aluminum would dissolve in CuCl2.
The original solution Copper (II) chloride is blue in color, and when the aluminum sulfide was placed in, the solution began to burn and the CuCl2 began to work at the aluminum, turning it into a dark brown color of porous material. The reaction was exothermic, which is when a reaction releases heat. The way to determine this if the test tube felt hot after the heat was released, since it elevates the temperature of the mixture. The reactants were Copper (II) chloride and Aluminum, which produced Copper metal and Aluminum Chloride. The color changed from brown to a murky green and eventually clear by the end of the lab.
In the aluminum half-reaction, it has a charge of three and when it is reduced, it gains three electrons. It’s potential was -1. 70. For the half reaction for Copper, the copper ion has a positive charge of two and when reduced gains an electron of two. It’s potential is 0. 34 V. In order for a reduction reaction to occur, Aluminum must be reduced and Copper oxidized. Therefore, the overall cell potential is 2. 04. equation: Al3+ + 3e- > Al(s) E? cell= -1. 70 V Cu2+ +2e- >Cu(s) E? cell= 0. 34 V 2Al(s) + 3Cu2+ > 2Al3+ + 3Cu(s) E? cell= 2. 04 V In part three, copper was placed in a silver nitrate solution to see if it would dissolve.
The solution wa initially clear and after the Copper was placed in, changed to a light shade of blue. The copper seems to crystalize and form a fuzzy texture on the outside, which was very hard to mix in the solution. The color eventually changes from white to a light brown. The products produced are Copper nitrate and Silver metal and that’s because the reactivity of copper is stronger than silver and displaces the silver to form copper nitrate. This is a redox reaction, in which the copper gives up two of its electrons to the silver ion, which is being reduced, to form solid silver.
Equation: Ag+ + e > Ag(s) E? ell= 0. 80 V Cu2+ + 2e > Cu(s) E? cell= 0. 34 V 2Ag+ (aq) + Cu(s) > 2Ag(s) + Cu2+(aq) E? cell= 0. 46 V 2AgNO3 (aq) + Cu(s) > Cu(NO3)2(aq) + 2Ag(s) Lastly, for part four, the Faraday’s constant and Avogadro’s number was found using an electrolytic cell. The electrolytic cell with two copper electrodes were immersed in a diluted hydrogen sulfide solution. The power supply provided the electrical current, which was required to oxidize copper to Cu2+ at the anode and reduce H+ to H2 at the cathode. In order to determine Avogadro’s number, the number of copper atoms per mole of copper must be determined.
During electrolysis, the positive pin of the power supply is losing mass as the copper atom is converted to copper ions. However, the hydrogen gas, the cathode, is being reduced and the hydrogen gas produced from this is collected up through the burette. To calculate Faraday’s constant, the total amount of electrical current flowing through must be measured, which can be done by changing the number of charges of moles of electrons. For the first part, the calculated Faraday’s constant and Avogadro’s number were: [86,900c/mol e-] and [5. 368×1023]. That was for how much copper ions were produced.
Whereas for the measurement of Hydrogen gas produced, the Faraday’s constant and Avogadro’s number were: [80,627 c/mol e-] and [5. 033×1023]. Equation: Cu(s) > Cu2+(aq)+ 2e- 2H+(aq) +2e- > H2(g) 2H+(aq) + Cu(s) > Cu2+ + H2(g) There were multiple errors that occurred during this experiment, which was predicted since so many things had to be recorded and kept track of. For part one, it was hard to keep the filter paper down on the Petri dish, which may have accounted for some of the solutions ending up mixed together on certain parts of the paper or touching the metal and causing an error in the results.
Overwhelmingly, part one was accurate enough with what was predicted; therefore, the error was not an impacting part. For parts two and three, the error resided in how much copper was placed in the solution or how diluted the solution was, preventing the full scope of the reaction from occurring. Lastly, for part four, the biggest error was setting up the inverted burette and making sure the system was tight to not let any air through. Secondly, determining the solution in the first gradation, which could’ve been a lot more or less.
Determining the total volume of gas inside the burette was difficult considering that the beaker was a bigger and water had to be added to match the height of the burette when it reached zero. Some of the water added slipped inside the burette and changed the final volume. Therefore, the weight of the final water was significantly different than what it was supposed to be. Overall, these errors impacted the final Avogadro’s number and Faraday’s constant, since the percent error for both were around 12%, which indicated that many errors led to such a large difference in the actual number and the experimental numbers.
