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Electrochemical Cells and Cell PotentialsHands-On Labs, Inc. Version 42-0153-00-02

Review the safety materials and wear goggles when working with chemicals. Read the entire exercise before you begin. Take time to organize the materials you will need and set aside a safe work space in which to complete the exercise.

Experiment Summary:

You will learn about galvanic cells and how cell potential is calculated. You will prepare a copper/zinc galvanic cell and measure the cell potential of the reaction. You will monitor the potential of the cell as the reaction proceeds.

EXPERIMENT

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Learning ObjectivesUpon completion of this laboratory, you will be able to:

● Define electrochemistry and compare redox, oxidation, and reduction reactions.

● Describe electrochemical cells including the flow of electricity through a galvanic cell.

● Predict the anode and cathode of a redox reaction using the standard reduction potentials.

● Construct a galvanic cell.

● Operate a multimeter and interpret voltage data.

● Calculate the standard cell potential for a redox reaction.

Time Allocation: 4 hours

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Experiment Electrochemical Cells and Cell Potentials

MaterialsStudent Supplied Materials

Quantity Item Description1 Camera, digital or smartphone1 Pair of scissors1 Roll of paper towels

HOL Supplied Materials

Quantity Item Description1 Digital multimeter1 Filter paper, 20 cm x 20 cm2 Glass beakers, 100 mL2 Jumper cables1 Pair of gloves1 Pair of safety goggles1 Plastic cup, 9 oz1 Experiment Bag: Electrochemical Cells and Cell Potentials

1 - Copper sulfate (CuSO4), 1.0 M, 75 mL 1 - Potassium chloride (KCl), 1.0 M, 30 mL 3 - Strips of copper metal, 2 in. x ¼ in. 3 - Strips of zinc metal, 2 in. x ¼ in. 1 - Zinc sulfate (ZnSO4), 1.0 M, 75 mL

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.



BackgroundElectrochemistry

Electrochemistry is the study of the electrical aspects of chemical reactions, concerned with two processes: the generation of an electrical current resulting from a spontaneous chemical reaction, and the use of an electrical current to produce a chemical reaction. These two processes describe oxidation-reduction (redox) reactions. A redox reaction is a chemical reaction in which there is a transfer of electrons (change in oxidation state) from one substance to another. The reaction is termed “redox” because it is composed of two half-reactions: an oxidation reaction in which electrons are lost and a reduction reaction during which electrons are gained. In the oxidation reaction the loss of electrons causes an increase in the oxidation number. Likewise, in a reduction reaction the gain of electrons causes a decrease in the oxidation number. See Figure 1.

Figure 1. Redox reaction between zinc and copper. The full reaction is shown in the top line. In the middle line is the oxidation reaction; notice that zinc loses two electrons to form the zinc ion. In the bottom line is the reduction reaction; notice that copper ion gains two electrons to

form the copper atom. The electrons gained and lost in the half-reactions cancel each other out in the full redox reaction.

Electrochemical and Galvanic Cells

A device that uses redox reactions to either use or produce electricity is called an electrochemical cell. There are two types of electrochemical cells: electrolytic cells, which use electrical energy, and galvanic cells, which produce electrical energy from a spontaneous redox reaction. A spontaneous reaction occurs naturally and does not require external influence (such as electrical energy). The focus of this experiment will be on galvanic cells. See Figure 2.



Figure 2. A simple galvanic cell for the redox reaction between zinc and copper. Oxidation occurs at the anode end, as copper gains electrons. Reduction occurs at the cathode end, as

zinc donates electrons. The voltmeter measures the amount of electrical energy produced by the cell.

In a galvanic cell the oxidation and the reduction portions of the redox reaction occur in separate locations (such as glass beakers), with a wire to facilitate the transfer of electrons between the locations. As shown in Figure 2, the wire may be attached to a voltmeter that measures the potential difference of electrical charge between the two locations. If a light bulb were hooked up to the wire, the light would burn dimly when a small potential difference exists and brightly when a large potential difference exists. In each of the two locations, an electrode is placed in a solution containing the same ion as the electrode. For example, in Figure 2, a copper electrode is placed in the copper solution and a zinc electrode is placed in the zinc solution. The electrode where oxidation occurs is called the anode, and the electrode where reduction occurs is called the cathode. To complete the cell (electrical circuit), the two locations are connected with a medium that facilitates the transfer of the ions (zinc ions and copper ions) from one location to another. This connection between the two half-cells is called the salt bridge, and it contains an inert electrolyte solution. A solution is inert if it does not react with the ions of either the electrodes or the solutions holding the electrodes. When the galvanic cell is complete, the electrons flow through the cell, from the anode to the cathode.



Reduction Potentials

A galvanic cell produces electrical energy that can be measured by a voltmeter. The cell voltage is the difference in electric potential between the cathode and the anode. The total amount of electric energy that a cell is expected to produce is called the standard cell potential (E°cell). Standard cell potential is calculated based on the assumption that the cell is in standard state conditions: the concentration of anode solution and cathode solution is 1M, the pressure is 1 atmosphere, and the temperature is 25°C. The standard cell potential is the contribution of standard reduction potential from the reduction half-reaction (E°cathode) and the standard reduction potential from the oxidation half-reaction (E°anode), as shown in the equation below:

The standard reduction potentials of half-reactions are constants. See Table 1 for a list of standard reduction potentials for a number of half-cell reactions. All half-reactions are shown as reduction reactions, hence standard reduction potentials.

Table 1. Standard Reduction Potentials.

Half-Reaction E°(Volts)F2(g) + 2e- → 2F-(aq) +2.87

Cl2(g) + 2e- → 2Cl-(aq) +1.36Br2(l) + 2e- → 2Br-(aq) +1.07Ag+(aq) + e- → Ag(s) +0.80

Fe3+(aq) + e- → Fe2+(aq) +0.77Cu2+(aq) + 2e- → Cu(s) +0.34

One of the most common galvanic cells is the battery. A battery contains a positive electrode (the

cathode) and negative electrode (the anode). These are denoted by “+” and “-“ symbols on the side of the battery. The electrodes take up most of the internal space inside the battery and access areas where chemical reactions occur. The anode experiences an

oxidation reaction in which charged ions interact with the anode to produce and release electrons. The cathode experiences a reduction reaction, whereby electrons are absorbed. The reactions result in the production of electricity, energy that travels in a circuit to

power cell phones, flashlights, and cars. © Eric Strand, © ekler



Half-Reaction E°(Volts)2H+(aq) + 2e- → H2(g) 0.00Fe2+(aq) + 2e- → Fe(s) -0.44Zn2+(aq) + 2e- → Zn(s) -0.76Al3+(aq) + 3e- → Al(s) -1.66

Mg2+(aq) + 2e- → Mg(s) -2.37Ca2+(aq) + 2e- → Ca(s) -2.87

K+(aq) + e- → K(s) -2.93

The more positive the reduction potential, the larger the ability of the half-reaction to behave as the oxidizing agent. Likewise, the more negative the potential, the larger the ability of the half-reaction to behave as the reducing agent. Given two half reactions, the one with more negative potential value will be the oxidizer. For example, consider the role of zinc as a reducer in Equation 1 below and as an oxidizer in Equation 2 below:

In equation 1, the cell potential of the half-reaction of zinc is -0.76V and the cell potential of the half-reaction of copper is +0.34V. In this reaction, the cell potential of the zinc is much more negative than the copper, and thus the zinc acts as the reducing agent (anode) in the reaction.

In equation 2, the cell potential of the half-reaction of zinc is -0.76V and the cell potential of the half-reaction of calcium is -2.87V. The cell potential of the calcium is much more negative than the zinc, and thus the calcium acts as the reducing agent (anode) in the reaction. The driving force of a reaction, pulling electrons from the anode in one location to the cathode in the other location, is dependent on the difference between the cell potentials of the half-reactions. The larger the difference, the more electrical energy the redox reaction will create.



The standard cell potentials for Equation 1 and Equation 2 are calculated below:

From the calculations, more electrical energy will be produced from the reaction occurring in Equation 2 (2.11V) than the reaction occurring in Equation 1 (1.00V). As a redox reaction proceeds, and the electrons travel from the anode to the cathode, the total cell potential for the reaction will decrease.

In the experiment, a galvanic cell for the redox reaction between copper and zinc will be prepared. In a galvanic cell, the Ecell must be positive for a spontaneous reaction to occur. The zinc solution will be zinc sulfate (ZnSO4) and the copper solution will be copper sulfate (CuSO4). The direction of electron transfer in the redox reaction will be tested by dipping the copper electrode directly into the zinc solution and the zinc electrode directly in the copper solution to see which electrode becomes plated with the ion of the solution. The total potential of the cell will be calculated and compared to the total amount of electrical energy produced in the galvanic cell, as measured with a multimeter.



Exercise 1: Construction of a Galvanic CellIn this exercise, you will create and experiment with a galvanic cell.

Procedure

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). See Figure 3.

Figure 3. Cutting a strip of filter paper.

3. Fold the strip of filter paper in half (widthwise) and then in half again. See Figure 4.

Figure 4. Folding the filter paper in half and then in half again.

4. 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. See Figure 5.



Figure 5. Folding filter paper in cup. The potassium chloride is added to the cup to over the filter paper.

6. 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.

7. 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.

8. Pick up a fresh strip of zinc and insert one 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. See Figure 6.

9. Pick up a fresh strip of copper and insert one 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. See Figure 6.

Figure 6. Metal in solutions. A. Zinc being inserted into copper sulfate. B. Copper being inserted into zinc sulfate.



10. Observe the 2 metal strips and record observations in Data Table 1 in your Lab Report Assistant.

11. From the observations, determine which of the 2 reactions is spontaneous. Record this in the observations section of Data Table 1.

12. Set up the multimeter as follows and see Figure 7:

a. Make sure the on/off switch of the multimeter is in the “off” position.

b. Place the end of the black probe into the bottom right hole of the multimeter.

c. Place the end of the red probe into the hole directly above the location of the black probe. Ensure that the probes are pushed all the way into the multimeter.

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.

Figure 7. Multimeter setup.

13. 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. See Figure 8.



Figure 8. Salt bridge. Notice that either end of the salt bridge is fully submerged in solution.

14. Clip a fresh piece of zinc onto one of the jumper cable clips and clip a fresh piece of copper onto the other jumper cable clip.

15. 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. See Figure 9.

16. 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. See Figure 9.

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.

Figure 9. Metals placed into their solutions. Notice the placement of the metal and the jumper cable clips relative to the solution.



17. Turn the multimeter 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 multimeter and swap the jumper cable clips from one metal to the other. For example, if a negative voltage was measured with the setup in Figure 9, the black jumper cable clip would be switched to hold the zinc, and the yellow jumper cable clip would be switched to hold the copper.

18. When the metals and jumper cable clips are arranged so that the multimeter has a positive reading, allow approximately 5 minutes for the multimeter reading to stabilize. When the multimeter reading has stabilized record the voltmeter reading in Data Table 2 in your Lab Report Assistant, under 0 minutes.

19. Look at a clock or watch and record the multimeter reading for the galvanic cell every 15 minutes for 2.5 hours.

20. 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.

21. Record the equation for the complete redox reaction occurring in the galvanic cell in Data Table 3.

22. Calculate the standard cell potential for the redox reaction occurring in the galvanic cell, and record in Data Table 3.

23. When all multimeter 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.

24. When you are finished uploading photos and data into your Lab Report Assistant, save and zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the correct format.

Cleanup:

25. Turn the multimeter off and carefully take apart the galvanic cell.

26. Properly dispose of solutions, metal pieces, and the salt bridge.

27. Wash lab equipment with soap and water and thoroughly dry.

28. Return cleaned equipment to the lab kit for future use.



QuestionsA. 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.

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.

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

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.



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