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Electrochemistry

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Chris Caron

on 2 June 2013

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Transcript of Electrochemistry

Electrochemistry Lab Compilation Chris Caron
Honors Chemistry
Ms. Farrell
D Block
April 24, 2013 Objective The objective of this lab was to determine the half reactions for Copper and Zinc Sulfate, Iron and Copper (II) Sulfate, Zinc and Copper (II) Sulfate,and Copper and Silver Nitrate.The anode, cathode, flow of electrons, and spectator ions were identified. The EMF of the cell was determined. The reactions were then observed and explained. Additionally, students must create a working battery, and then combine it with another groups battery. Materials (4) Petri Dishes
(2) 1x60 mm Copper wiring
(1) Iron nail
(1) 10x15 mm Zinc Strip
100 mL of Copper (II) Sulfate
100 mL of Zinc Sulfate
25 mL of Silver Nitrate
(1) 2 Alligator clips with connecting wire
(1) Pen
Paper
(2) Ceramic tubing (Porous)
(2) Iron Nail (2) 20x100 Copper Strip
(2) 20x100 mm Zinc strip
(1) Voltmeter
(2) Plastic container with clamps
(2) 100 mL beaker
20x20 mm filter paper
(1) 9 Volt Battery
(1) Penny
(1) Nickel
25x25 mm Aluminum Foil
(1) Micropipette
3 drops of Sodium Sulfate Apparatus One Procedure/Results Experiment 1 Copper + Zinc Sulfate 1. The half-reactions were written down and each reaction was determined to be spontaneous or non-spontaneous.

2. Petri Dish was filled with 25 mL of Zinc Sulfate.

3. 1 piece of 1x60 mm copper wiring was placed in the Petri Dish.

4. The reaction was observed. Experiment 2 Zinc + Copper (II) Sulfate 1. The half-reactions were written down and each reaction was determined to be spontaneous or non-spontaneous.

2.Petri Dish was filled with 25 mL of Copper (II) Sulfate.

3. The 10x15 mm Zinc strip was placed in the Petri Dish.

4. The reaction was observed. Experiment 3 Iron + Copper (II) Sulfate 1. The half-reactions were written down and each reaction was determined to be spontaneous or non-spontaneous.

2.Petri Dish was filled with 25 mL of Copper (II) Sulfate.

3. The iron nail was placed in the Petri Dish.

4. The reaction was observed. Experiment 4 Copper + Silver Nitrate 1. The half-reactions were written down and each reaction was determined to be spontaneous or non-spontaneous.

2. Petri Dish was filled with 25 mL of Silver Nitrate.

3. One piece of 1x60 mm copper wiring was placed in the Petri Dish.

4. The reaction was observed. Experiment: Making a Battery 1. The half- reactions were written down and the cathode/anode were identified

2. Apparatus was assembled. Labeled "Cup 1".

3. One end of the voltmeter was connected to the copper strip, the other was connected to the zinc strip. The voltage was then recorded.

4. A second apparatus was assembled. Labeled "Cup 2"

5. One of the alligator clips was connected to the copper strip in Cup 1, and the other clip was connected to the zinc strip in Cup 2.

6.One end of the voltmeter was connected to the Zinc strip in Cup 1and the other end was connected to the Copper strip in Cup 2. The voltage was then recorded. 20x100 mm Copper strip 20x100 mm Zinc strip Plastic container with 2 clamps Ceramic tubing filled with 25 mL of Copper Sulfate Alligator Clip The Copper wiring became coated in a grey/ brown substance (silver) and the silver nitrate turned a light blue color. No reaction occurred. The Zinc strip turned a black color and disintegrated into smaller pieces. Some steam formed on the color of the Petri Dish. The iron nail, became coated in a brown color ( copper) and eventually, the copper sulfate turned green and then turned brown. The reading of the voltage from step #3 was +1.08 V, and the voltage from step #6 was 2.14 V. 100 mL Beaker Experiment 6:Electrochemical Analysis of Metals 1. The items were stacked in the following order: 20x20 mm Aluminum Foil was placed on the bottom, them the filter paper, 3 drops of Sodium Sulfate, and a penny.

2. The negative terminal of the 9 volt battery was touched to the aluminum foil, while the positive terminal was touched to the penny for no longer than 3 seconds.

3. The penny was removed, and the reaction was observed.

4. Steps 1-3 were repeated using an nickel and an iron nail. The nickel left a green imprint. The penny left a blue imprint. The iron nail left a brown imprint. Analysis Experiment 1 Copper + Zinc Sulfate Experiment 2 Zinc + Copper (II) Sulfate Experiment 4 Copper + Silver Nitrate Experiment 5 Making a Battery Experiment 3 Iron + Copper (II) Sulfate Experiment 6 Electrochemical Analysis of Metals In this reaction, Copper was oxidized, and Silver was reduced. This makes Copper the reducing agent and Silver the oxidizing agent. The spectator ion in this case was nitrate. The EMF of the cell was +0.28 V because the reduction of Silver is +0.80 V and the oxidation of Copper is -0.52 V. if we add these together, we get +0.28 V which is the EMF of the cell. This reaction was spontaneous because it occurred on its own. The oxidation half reaction was Cu --> Cu^+1 + e-, and the reduction half reaction was Ag^+1 + e- --> Ag. In this voltaic cell, the zinc was oxidized while the copper was reduced. Therefore, the reducing agent was zinc metal while the oxidizing agent was copper in the form of copper sulfate (CuSO4). The electrolyte, or spectator ion, in this case was sulfate (SO4^-2) which moved from the copper(II) to the zinc(II). The EMF of this cell was +1.05 V. This is due to the fact that the reduction of copper(II) to copper has a voltage of +0.34 V and that the oxidation of zinc to zinc(II) has a voltage of +0.76 V. These two voltages were added together in order to find the cell potential. One can see that the predicted voltage should be +1.10 V, but this was not the voltage in the readings taken. +1.05 V is very close though making it an acceptable answer. As one can see in the images above, this reaction was spontaneous since it occurred on its own. As a clarification, a voltaic cell is an electrochemical cell used to convert chemical energy into electrical energy while a battery is a group of voltaic cells connected together. The anode in this lab was zinc while the cathode was copper. The flow of electrons was from zinc to copper. The electrolyte (spectator ion) lost its bond from copper(II) and went through the porous ceramic tubing to bond with the newly formed zinc(II) molecules. The oxidation half reaction for this lab was Zn -> Zn^+2 + 2e^- and the reduction half reaction was Cu^+2 + 2e^- -> Cu. All the LEDs except the blue one lit up due to the fact that blue and violet wavelengths require a higher amount of energy to be produced.When the two voltaic cells were connected, a charge of +2.11 V was produced. Hypothetically, this number should be +2.20 V since two batteries with a hypothetical charge of +1.10 V were connected. In reality though, this numbers are not always produced due to a variety of reasons. +2.11 V is very close to +2.20 V making it an acceptable answer. The reading for voltage double because the cathode(+) of one voltaic cell was connected to the anode(-) of another. The flow of electrons was from the zinc strip in cup 1 to the copper strip in cup 1 to the zinc strip in cup 2 and to the copper strip in cup 2. In the first part of this lab, the copper (penny) was oxidized while the water was reduced. Therefore, the reducing agent was copper (penny) while the oxidizing agent was water (H2O). The electrolyte, or spectator ion, in this case was sulfate (SO4^-2) which came from sodium sulfate which was necessary to keep the reaction from producing charged ions. The sulfate moved from the sodium ions to the copper(II) ions. The oxidation half reaction for this experiment would be Cu -> Cu^+2 + 2e^- and the reduction half reaction was 2H2O -> H2 + 2OH^- . The predicted EMF of this cell would be -1.35 V; this was not measured during the lab. Based off of this voltage, one can see that this reaction would be non-spontaneous. Therefore, an external power source, in this case a 9 volt battery, was used to form an electrolytic cell. The touching of the positive terminal to the penny forced the penny (copper) to become the cathode and oxidizing agent in the reaction. The touching of the negative terminal to the aluminum foil forced the foil to become the anode and reducing agent thereby reducing the water molecules. The sodium sulfate then came into play by balancing out the negative and positive ions. This was true for all of the reactions, except copper was replaced with whatever metal was touching the positive terminal. The blue markings on the paper were formed due copper(II) sulfate, which has a distinct blue color.
In the second part of the lab, the copper was oxidized while the water was reduced. As a side note, the reasons copper was oxidized and not nickel is that a nickel coin is made up of 75% copper and 25% nickel. Therefore, copper was oxidized rather than nickel. The reducing agent was copper while the oxidizing agent was water. The electrolyte, or spectator ion, in this case was sulfate (SO4^-2) which came from the sodium sulfate that was crucial in the balancing of the equation. The sulfate moved from the sodium ions to the copper(II) ions. The oxidation half reaction for this lab was Cu -> Cu^+2 + 2e^- and the reduction half reaction was 2H2O + 2e^- -> H2 + 2OH^-. The predicted EMF of this cell would be -1.35 V; this was not measured during the lab. Based off of this voltage, one can see that this reaction was non-spontaneous.
In the third part of the lab, iron was oxidized while the water was reduced. Therefore, the reducing agent was iron while the oxidizing agent was water. The electrolyte, or spectator ion, in this case was sulfate (SO4^-2) which came from the sodium sulfate that was crucial in the balancing of the equation. The sulfate moved from the sodium ions to the iron(II) ions. The oxidation half reaction for this lab was Fe -> Fe^+2 + 2e^- and the reduction half reaction was 2H2O + 2e^- -> H2 +2OH^-. The predicted EMF of this cell would be -0.79 V; this was not measured during the lab. Based off of this voltage, one can see that this reaction was non-spontaneous. The explanation for how this cell worked can be viewed at the end of paragraph 1; simply replace the word "penny/copper" with "iron." Unlike copper, the iron nail did not leave behind a stain. One explanation for this is that FeSO4 may not have a color or that it was very faint. In this reaction, iron was the cathode and oxidizing agent while sodium was the anode and reducing agent. Experiment 4 In this reaction, it was concluded that the half reactions would be Cu --> Cu^+1 + e-, and Ag^+1 + e- --> Ag. The EMF of the cell was +0.28 V because the reduction of Silver is +0.80 V and the oxidation of Copper is -0.52 V. This reaction was spontaneous because it occurred on its own. Experiment 2 In this experiment, is was concluded that the half reactions would be Zn --> Zn^+2 +2e-, and Cu^+2 +2e- --> Cu. The electron flow was from iron to copper and this reaction was spontaneous because it produced a positive voltage of +1.10 V. Zinc was the anode and copper was the cathode. Sulfate ions were the electrolytes, that moved from the copper(II) ions to the zinc(II) ions. The electrons flowed from zinc to copper. Experiment 5: Making a Battery Experiment 1 From this experiment, it was concluded that the half reactions would hypothetically be Cu --> Cu^+2 + 2e- and Zn^+2 + 2e- --> Zn. This did not happen though because the cell was non-spontaneous and produced a negative voltage of -1.10 V. Experiment 6: Electrochemical Analysis of Metals Experiment 3 In this reaction, it was concluded that the half reaction would be Fe --> Fe^+2 + 2e-, and Cu^+2 + 2e- --> Cu. The reaction was spontaneous because it produced a +0.78 V. In this reaction iron was the anode, copper was the cathode, and sulfate was the spectator ion. The electrons flowed from iron to copper. Conclusion In this experiment, no reaction occurred, making it non-spontaneous. The copper was oxidized and the zinc was reduced, making copper the reducing agent and zinc the oxidizing agent. The oxidation half reaction was Cu --> Cu^+2 + 2e-, while the reduction half reaction was Zn^+2 + 2e- --> Zn. The oxidation reaction was equal to -0.34 V and the reduction reaction was equal to -0.76 V, making the cell potential -1.10 V. The spectator ion was sulfate, the anode was copper, and the cathode was zinc. The flow of electrons in the reaction would be from copper to zinc, but only if energy was added. In this experiment, the EMF of the cell was positive, making it a spontaneous reaction. Iron was oxidized (reducing agent), and zinc was reduced (oxidizing agent). The oxidation half reaction was Fe --> Fe^+2 + 2e- , and the reduction half reaction was Cu^+2 + 2e- --> Cu. The oxidation reaction was equal to +0.44 V and the reduction reaction was equal to +0.34 V, making the cell potential equal to +0.78 V. In this reaction iron was the anode, copper was the cathode, and sulfate was the spectator ion. The electrons flowed from iron to copper. In this experiment, the EMF of the cell was positive, making it a spontaneous reaction. Zinc was oxidized (reducing agent), and copper was reduced (oxidizing agent). The oxidation half reaction was Zn --> Zn^+2 +2e-, and was equal to +0.76 V. The reduction half reaction was Cu^+2 +2e- --> Cu, and was equal to +0.34 V, making the cell potential +1.10 V. Zinc was the anode and copper was the cathode. Sulfate ions were the electrolytes, that moved from the copper(II) ions to the zinc(II) ions. The electrons flowed from zinc to copper. From this experiment it was concluded that the half reactions for a Cu | CuSO4 || ZnSO4 | Zn voltaic cell are Zn -> Zn^+2 + 2e^- (oxidation) and Cu^+2 + 2e^- -> Cu (reduction). In addition, it was discovered that these reactions were spontaneous and produced a cell potential of +1.05 V. Zinc was the anode and reducing agent in this reaction and copper was the cathode and oxidizing agent. The electrolyte in this reaction was sulfate which moved from the copper(II) ions to the zinc(II) ions. Furthermore, when two voltaic cells are successfully combined, they form a small battery that is double each of their charges which was +2.11 V. From the first part of this experiment it was concluded that the half reactions for a copper-water reaction are Cu -> Cu^+1 + 1e^- (oxidation) and H2O + 2e^- -> H2 + 2OH^- (reduction). The direction of the electron flow was from the copper to the water. In addition, it was discovered that these reactions were non-spontaneous and produced a cell potential of -1.35 V. Therefore, an external power source (9 volt battery) was used to force this reaction. Copper was the anode and reducing agent while aluminum foil (it directly caused the water to be reduced) was the cathode and the oxidizing agent. The electrolyte in this reaction was sulfate which moved from the sodium ions to the copper(II) ions.The second part of the experiment was identical to the first due to the fact that a nickel coin mainly comprised of copper.From the third part of this experiment we concluded that the half reactions for a iron-water reaction are Fe -> Fe^+2 + 2e^- (oxidation) and H2O + 2e^- -> H2 + 2OH^- (reduction). The direction of the electron flow was from the iron to the water. In addition, it was discovered that these reactions were non-spontaneous and produced a cell potential of -0.79 V. Therefore, an external power source (9 volt battery) was used to force the reaction. Iron was the anode and reducing agent in this reaction while water was the cathode and oxidizing agent. The electrolyte in this reaction was sulfate which moved from the sodium ions to the iron(II) ions.
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