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Electrochemistry lab report answers

08/01/2021 Client: saad24vbs Deadline: 2 Day

Electrochemical Cells and Cell Potentials


Objective:


The purpose of this experiment is to create and experiment galvanic cell and collect/interpret data by using a multimeter to describe the flow of electrons. The we g=had to determine how it is calculated by using the formulas given.


Procedure:


Exercise 1: Construction of a Galvanic Cell


1. Gather all of the supplies listed in the materials list. 2. Use the scissors to cut a strip of the filter paper approximately 1.5 inches in width (1/4 the size of the sheet of filter paper). 3. Cut a strip of filter paper and fold the strip of filter paper in half (widthwise) and then in half again. 4. Folding the filter paper in half and then in half again and put on the safety gloves and goggles. 5. Create the salt bridge by carefully winding the folded filter paper into a circle so that it fits into the bottom of the 9 oz. plastic cup. Add the potassium chloride to the cup with the filter paper until the paper is completely covered with the potassium chloride. 6. Folded filter paper in cup. The potassium chloride is added to the cup to cover the filter paper. 7. Allow the paper to soak up the potassium chloride for a minimum of 10 minutes or until you are ready to add it to the galvanic cell, as described later in the experiment. 8. Place the 2 glass beakers on a table. Add approximately 45 ml of zinc sulfate (approximately ½ of the bottle) to one of the beakers. To the second beaker, add approximately 45 ml of copper sulfate. 9. Pick up a fresh strip of zinc and insert 1 end of it into the copper sulfate solution. After approximately 5 seconds, remove the zinc from the copper sulfate and place it on a piece of paper towel.10. Pick up a fresh strip of copper and insert 1 end of it into the zinc sulfate solution. After approximately 5 seconds, remove the copper from the zinc sulfate and place it on the piece of paper towel.11. Metal in solutions. A. Zinc being inserted into copper sulfate. B. Copper being inserted into zinc sulfate and observe the 2 metal strips and record observations in Data Table 1 in your Lab Report Assistant. 12. From the observations, determine which of the 2 reactions is spontaneous. Record this in the observations section of Data Table 1. 13. Set up the voltmeter as follows; a. Make sure the on/off switch of the voltmeter is in the "off" position. b. Place the end of the black probe into the bottom right hole of the voltmeter. c. Place the end of the red probe into the hole directly above the location of the black probe. d. Turn the voltage dial so that the arrow end of the dial is pointing to 20 DCV. e. Add 1 jumper cable clip to each end of the probes. It does not matter what color jumper cable clips are provided in your kit, or which color is attached to either probe. 14. Put the salt bridge into place by submerging 1 end on the copper sulfate and the other end in the zinc sulfate. Adjust the beakers as necessary so that the salt bridge does not sink between the beakers. 15. Clip a fresh piece of zinc onto 1 of the jumper cable clips and clip a fresh piece of copper onto the other jumper cable clip. 16. Place the zinc into the zinc sulfate solution, so that the metal is submerged in the solution, but the jumper cable clip is above, and not touching, the solution or salt bridge. 17. Place the copper into the copper sulfate solution, so that the metal is submerged in the solution, but the jumper cable clip is above, and not touching the solution or salt bridge. Note: It may take a few minutes to find the correct placement of the copper and zinc into the solutions to keep the jumper cable clip above the solution. Adjust the jumper cable clips as necessary to find the correct placement. 18. Turn the voltmeter on, and observe whether the total voltage is positive or negative. If the voltage reads positive, the galvanic cell was prepared correctly and can be allowed to progress. If the voltage is negative, quickly turn off the voltmeter and swap the jumper cable clips from one metal to the other. 19. When the metals and jumper cable clips are arranged so that the voltmeter has a positive reading, allow approximately 5 minutes for the voltmeter reading to stabilize. When the voltmeter reading has stabilized record the voltmeter reading in Data Table 2 in your Lab Report Assistant, under 0 minutes. 20. look at a clock or watch and record the voltmeter reading for the galvanic cell every 15 minutes for 2.5 hours. 21. While the reaction in the galvanic cell is progressing, use Table 1 in the Background section to determine the 2 half-reactions and standard reduction potentials for the redox reaction occurring in your galvanic cell. Record the half reactions, identifying which is the oxidation and which is the reduction half-reaction. Also record the corresponding reduction potentials in Data Table 3 in your Lab Report Assistant. 22. Record the equation for the complete redox reaction occurring in the galvanic cell in Data Table 3. 23. Calculate the standard cell potential for the redox reaction occurring in the galvanic cell, and record in Data Table 3. 24. When all voltmeter readings have been taken and recorded in Data Table 2, take a photograph of your galvanic cell. In the photograph, include a small piece of paper that displays your name and the date. Resize and insert the photograph in Data Table 4 in your lab Report Assistant. Refer to the appendix entitled, "Resizing an Image" for guidance with resizing an image. 25. Turn the voltmeter off and carefully take apart the galvanic cell. Pour the solutions down the drain with ample amounts of running water. Place the salt bridge and metal pieces into the trash. 26. Wash all glassware with soap and water and return equipment to the kit for future use. 27. When you are finished uploading photos and data into your Lab Report Assistant


Exercise 1: Construction of a Galvanic Cell


Data Table 1. Spontaneous Reaction Observations.


Metal in Solution


Observations


Zinc in Copper Sulfate


The solution turns colorless and black precipitate forms


Copper in Zinc Sulfate


No change


Data Table 2. Voltmeter Readings.


Time (minutes)


Voltmeter Reading (Volts)


0


1.1


15


1.057


30


0.984


45


0.975


60


0.968


75


0.957


90


0.953


105


0.950


120


0.949


135


0.947


Data Table 3. Standard Cell Potential.


Equation


E°(Volts)


Oxidation Half-Reaction


+0.76


Reduction Half-Reaction


+0.34


Redox Reaction


+1.10


Data Table 4. Galvanic Cell Setup.


Photograph of galvanic cell




Questions


A. What were the concentrations of the solutions (zinc solution, copper solution, and salt bridge)? Were the concentrations consistent with those of standard state conditions? Explain your answer.


Concentration of copper solution and zinc solution was taken as 1 M each. And the salt bridge is not an issue to be considered.


As the reaction proceeds, although movement of anion allows the overall charges in solution to remain neutral, the net movement of anion produces a concentration gradient across the two solutions. In other words, after a while the net concentration of ions in the zinc oxidation side will be greater than in the coper reduction side. This concentration gradient will oppose movement of anion. In consideration of keeping the overall concentration of ions in balance between the two sides, cations will also be moving to the right.


B. Was the amount of electric energy produced in your galvanic cell consistent with the standard cell potential of the reaction (as calculated in Data Table 3)? Hypothesize why it was or was not consistent.


There is a little drop of resulting electrical energy from the standard cell potential of reaction.


The cause of voltage drop is: The actual electrical resistances in the "circuits", the cells themselves". The cell potentials that you calculate are the "ideal" situation and you would get those if there was not some electrical resistance. But like every machine has some friction, every circuit has some resistance, and the effect of the resistance is to lower the potential difference, the voltage of the cell. The internal resistance is not something you can get rid of. You can minimize it by using "clean" electrodes and keeping the distances between electrodes short.


You may also have an error in your voltmeter. You want to use a voltmeter with a high input impedance (resistance), ideally around 10 megohms. The voltmeter itself becomes part of the circuit and gives false readings. Avoid cheap voltmeters with lower input impedances


C. Was there evidence of electron transfer from the anode to the cathode? Use your data in Data Table 2 to explain your answer.


The cell potential is a measure of the tendency of the anode metal to be oxidized (lose electrons) and the tendency of the metal ions in the cathode compartment to be reduced (gain electrons). Thus, cell potentials vary with the composition of the substances being used as electrodes. Cell potentials also vary with the concentrations of ions and pressures of gases, as well as the temperature at which the reaction occurs. As the redox reaction proceeds, and the electrons travel from anode to the cathode, the total cell potential will decrease. Another evidence is that the direction of electron transfer is tested by dipping the copper electrode directly in zinc solution and zinc electrode directly in copper solution to see which electrode becomes plated with the ion of the solution.


D. For the following redox reaction in a galvanic cell, write the oxidation half-reaction and the reduction-half reaction, and calculate the standard cell potential of the reaction. Use Table 1 in the Background as needed. Explain how you identified which half-reaction is the oxidizer and which is the reducer. Show all of your work.


Equation


Eo(volts)


-0.34


Fe+3 (aq) + e- Fe+2(s)

+0.77


+0.78V


()*1 Eo = -0.34

(Fe+3 (aq) + e- Fe+2(s))*2 Eo= 0.77*2=1.54

Eo= Ereduction + Eoxidation = 1.54- (-0.34) = 1.88

Reducers lose electrons , so is reducer

And oxidizer gains electrons so Fe+3 (aq) + e- Fe+2(s) is oxidizer

3+2+2+


(s)(aq)(aq)(s)


Redox reaction: Cu + Fe Cu + Fe


®

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