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12 Dec 2019

1.Use the activity series from your text to determine whether tin will dissolve in a 1.0M acid solution (a solution of H+ ions) according to the following equation. Explain your answer. Sn(s) + 2H+ (aq) J Sn+2(aq) + H2(g)

2. Would a battery constructed of tin (Sn/Sn+2) and lead (Pb/Pb+2) be of much value? Why or why not? Hint: calculate the reduction potential for the battery.

3. Martin Smythe constructed a potato battery by inserting strips of copper and zinc into a potato. The metal strips were attached to the electrical leads of a clock. What is the voltage produced by the battery?

4. What is the purpose of the potato in question 3?

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Jean KeelingLv2

13 Dec 2019

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How do you do the math for this? I know cost of Zinc is $5.30/100 g, Silver is $0.56/g, and Lead is $2.45/100g..

The construction of a cell with maximum voltage

Use metal strips and solutions of your choice to determine which combination will give the largest potential. Several different metals and metal salt solutions will be available. In general, the solid metal should be the same metal as that in the salt solution it is in contact with. Utilize a table of standard reduction potentials to help decide which cells to construct. Keep in mind that you can connect more than one cell together. You can also vary the concentration of the metal salt solutions. A minimum of 3 different batteries should be made and tested. Your grade will not be based solely on the number of batteries you make, but also on the strategy used for choosing battery designs and the thoroughness of your investigation. In addition to finding the highest voltage you can achieve, determine which battery gives you the best ratio of volts to cost (V/$).

Battery 1.) HALF CELL ONE: 1.0 Mg²⁺ (aq) and 1.0 Mg (s) HALF CELL TWO: 1.0 M Zn²⁺ (aq) and Zn (s) observed potential = + 1.12 V

Battery 2.) HALF CELL ONE: 1.0 Mg²⁺ (aq) and 1.0 Mg (s) HALF CELL TWO: 1.0 M Ag⁺ (aq) and Ag (s) observed potential = + 2.39 V

Battery 3.) HALF CELL ONE: 1.0 Mg²⁺ (aq) and 1.0 Mg (s) HALF CELL TWO: 1.0 M Pb²⁺ (aq) and Pb (s) observed potential = + 1.42 V

You are asked to construct a voltaic cell that would simulate an alkaline battery by providing an electrical output of 1.50 VV at the beginning of its discharge. After you complete it, your voltaic cell will be used to power an external device that draws a constant current of 0.50 AA (amperes) for 2.0 hh. You are given the following supplies: electrodes of each transition metal from manganese to zinc; the chloride salts of the +2 transition metal ions from Mn2+Mn2+ to Zn2+Zn2+ (MnCl2MnCl2, FeCl2FeCl2, CoCl2CoCl2, NiCl2NiCl2, CuCl2CuCl2, and ZnCl2ZnCl2); two 100 mLmL beakers; a salt bridge; a voltmeter; and wires to make electrical connections between the electrodes and the voltmeter. The standard reduction potentials for the divalent metal ions are as follows: Half-reactionMn2+(aq)Fe2+(aq)Co2+(aq)Ni2+(aq)Cu2+(aq)Zn2+(aq)++++++2e−2e−2e−2e−2e−2e−→→→→→→Mn(s)Fe(s)Co(s)Ni(s)Cu(s)Zn(s)E∘(V)−1.18−0.44−0.28−0.250.34−0.76 Since you were tasked with constructing a cell that produces a potential of 1.50 VV, you decide to fine tune your concentrations to minimize the error. If you used the same concentration for each electrolyte, your voltaic cell would produce a potential of 1.52 VV. Therefore, you should adjust the concentration ratio that would lower your cell potential by 0.02 VV. You and a fellow researcher decide that you should prepare two 100 mLmL solutions for the voltaic cell. Your fellow researcher already prepared one of the electrolyte solutions for you: 0.228 MM CuCl2 CuCl2. At what concentration should you prepare the MnCl2MnCl2 solution such that your initial voltaic cell potential is 1.50 VV? Assume that the temperature is consistently 25 ∘C∘C. If the battery is permitted to discharge for a period of two hours at a constant current of 0.50 AA, what will be the concentration of the manganese(II) ion in the anode electrolyte solution at the end of the discharge? Similarly, what will be the concentration of the copper(II) ion in the cathode electrolyte solution at the end of the discharge? Given the nature of how a cell is constructed, you would usually deplete the ions at the cathode before completely oxidizing the anode (since the latter would break the circuit). This is because the complete oxidation of the anode would likely only occur if you were using an extremely thin wire of the metal. In this case, let us assume you constructed the voltaic cell using sufficiently sized rods for each metal. This means that the total time you can run your cell will depend on the initial copper ion concentration. Calculate the maximum run time for your voltaic cell if it runs at 0.5 AA and your initial concentration of CuCl2(aq)CuCl2(aq) was 0.228 M'

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OBJECTIVE:

In this lab, you will produce several different voltaic cells, measure the total voltage, and determine the corresponding half-cells and total potentials. A copper electrode and a copper gluconate electrolyte solution will be used in every system, and four additional half-cells with different metals and salts will be run against it to compare the half-cells and total overall voltage. The salt bridge, consisting of filter paper and potassium chloride, connects the two cells.

PROCEDURE

Put on your goggles and lab apron.

Label five clean, dry 100 mL beakers as copper gluconate, tin sulfate, aluminum sulfate, iron sulfate, and zinc

sulfate.

Using a clean, dry 100 mL graduated cylinder (and pipette if needed), measure and pour 25.0 mL of each

solution into the corresponding beaker. Be sure to thoroughly rinse the graduated cylinder with dH2O before measuring out a new solution. Some of your bottles of solution may need shaken up before using to re- dissolve anything that has settled out. If shaking doesnât completely work, use a clean, dry glass stirring rod to help break up and dissolve. Mix thoroughly. Just be sure to clean the stir rod before moving it to a different solution bottle.

Lightly sand the provided metal strips of copper, tin, aluminum, iron, and zinc down with sand paper to remove the oxidized layer. When finished, wipe off any debris. Set the strips aside. Be sure you have sanded these well to ensure reaction.

Use scissors to cut the filter paper into five strips. Each strip should measure about 20 mm (2 cm) x 150 mm (15 cm).

Note: When your electrochemical cell apparatus is set-up, each strip should be long enough to connect the electrolyte solutions in the two adjacent beakers (half-cells). The strips may look short now, but after they soak, you will see that are the proper length. See pic below.

Using a clean, dry 100 mL graduated cylinder, measure and pour 25.0 mL of potassium chloride solution into a 100 mL beaker. Place the filter paper strips into the beaker. Allow the strips to soak for approximately 5 minutes. This process creates the salt bridge. Be sure all papers are in contact with KCl and get soaked.

Place the metal copper strip in the copper gluconate beaker, allowing it to lean against the side of the beaker.

Repeat Step 7 for each of the metal strips and the corresponding beakers (see pic below).

Place one of the potassium chloride soaked filter papers between the copper gluconate solution and the tin

sulfate solution so that the filter paper is immersed in both solutions simultaneously. The KCl strip must be touching both solutions (itâs OK for the strip to cling to the wall of each beaker, so long as it is touching both solutions)! Pick up the strip by the middle, and drain excess solution by touching the ends to a paper towel. Position the paper strip salt bridge so that it makes contact with both solutions, but is not touching the electrodes.

Assemble the 7-Function Multi-Meter by firmly pressing the positive (red) Test Lead into the V?mA (center) jack. Press the negative (black) Test Lead into the COM (bottom) jack. Rotate the red, center dial until it is pointing towards the DCV 2000 m position (Hint: this position is similar to the northwest position on a compassâsee pic below). 2000 m is equal to 2 volts. That means your read out is now in mV, so you will have to divide the readout by 1,000 when recording your data in V. For example, in the picture below, my readout of 980 mV would be 0.980 V. Turn the multi-meter on.

Touch the red multi-meter probe to the top of the copper strip, and the black multi-meter probe to the top of the tin strip. This works best when you press lightly down on the top edge of the metal strip with the side of the probe, as this creates more constant contact and less fluctuation movement. See pic below. If you happen to have a set of alligator clips to connect the probes to the metal, that would work even better.

Note: The red probe should be placed on copper for all of the half-cell apparatuses. The black probe should be placed on the remaining metals.

Note: Alternatively, you can bend the metal strips to rest over the edge of the beaker, and to give a flat surface that makes it easier to rest the probes:

Record the voltage differential created between the two half-cell in Table 1 (see below). Watch the voltage for a few seconds and then take a reading. Some readings stabilize quickly, others will drift for as long as you want to watch. (Swirling the solution with the electrode can also change the reading.) Typically, waiting a few seconds after connecting the cell yields the most reliable reading.

Take a picture of your voltaic cell. This will be a little tricky since you have to hold the probes, but try to get someone to help. Otherwise, do the best you can. The digital readout on the multimeter must be visible for each sample, and both half cells must be visible and clearly labeled. Your face must be visible in the picture also, as well as a label with your name/student ID, the lab title, and the date. Unless all fourâyour face, name/ID, the lab title, and the date are clear, you will not receive credit. Save this picture, and combine all four voltaic cells (Sn, Al, Fe, Zn) into a single PDF document in order to submit it in the Lab 14 Assignment.

Discard the salt bridge, and repeat Steps 9 -13 three more times using the aluminum sulfate solution, iron sulfate solution, and the zinc sulfate solution. Each solution should be paired with the copper gluconate solution, placing the red probe of the volt meter on the copper each time. A new salt-bridge should be used for each electrochemical cell. Record your data in Table 1 (see below).

Wash your glassware (and plasticware) with warm water and soap. To minimize spotting and mineral contamination, it is best to do a final rinse with dH2O. Dry your glassware (or leave set to dry) and put it away. All chemical solutions used in this lab can now be safely washed down the sink, and solid metals can be thrown away or recycled. Clean up your lab and put away your materials.

RESULTS

Make a data table in either Word or Excel that shows the multi-meter voltage measurements (V), experimental half-cell potential using copper electrode as the standard, the theoretical half-cell voltage, percent error, the half-cell and net cell reactions, and free energy (?G). Your table should look something like the one on page 6 of this document. Also see the description below the sample table for help in determining all of these values. Make sure that the calculation sections shown below are included within your Table 1 document as well.

Save your Table 1 as a PDF file to be submitted with the Lab 14 Assignment. Be sure your document includes your name/student ID, lab title, and date. Unless all threeâyour name/student ID, the lab title, and the date are on the document and legible, you will not receive credit. All columns and values must be clearly labeled.

You should have = 1 PDF of a pic of you and your 4 electrochemical cells being measured (Step 13) + (Steps 16-17). All pictures, tables, and labels must be clear and properly labeled or they will not receive credit.

Question:

You probably observed the cell voltage change slightly while you were observing it. Explain what is occurring at the electrode that causes the observed voltage to change?

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