electric current: the movement of charged...

Electrochemistry Unit Name: _________________________________ of 16

Electrolysis of Solutions (Handwritten)Objectives• Observe the electrolysis of aqueous solutions and identify the reaction products at the cathode and

anode.

• Describe electrolysis reactions by writing half-reactions and net ionic equations for electrolysis reactions.

Introduction -

You have already learned that aqueous solutions of ionic compounds are electrolytes. They conduct electricity because water causes the ions to dissociate or break apart. The positive and negative ions are free to move about in solution, thus enabling them to carry an electric current.

You may have noticed that some of the solutions reacted chemically when you tested them for conductivity in Small-Scale Experiment 23. The ions in solution not only conducted electrons, they also gained and lost them. In other words, the ions were oxidized or reduced.

Electrolysis is the process by which an electric current causes a chemical reaction to take place. An electrolytic cell is a vessel in which electrolysis reactions occur. An electrolytic cell consists of two electrodes, a positive electrode called an anode, and a negative electrode called a cathode. When a voltage is applied to the cell, an electrolysis reaction takes place as shown in Figure 38.1.

Electroplating is an industrial electrolytic process used to deposit a thin layer of metal “plate” onto another metal. For example, electroplating is used to plate automobile bumpers with chromium to improve their appearance and to protect them from corrosion. Similarly, silverware is plated with silver. The object to be plated is the cathode, and the plating metal is the anode of an electrolytic cell. The aqueous solution contains ions of the plating metal.PurposeIn this lab you will investigate electrolysis of various aqueous solutions and use indicators to identify the reaction products at the cathode and anode. You will electroplate metals onto other metals and identify them.Safety• Wear your safety glasses.• Use small-scale pipets only for the carefully controlled delivery of liquids.MaterialsSmall-scale pipets of the following solutions:

Electrochemistry Unit Name: _________________________________ of 16water (H2O) sodium sulfate (Na2SO4)bromthymol blue (BTB) potassium iodide (KI)starch sodium chloride (NaCl)potassium bromide (KBr) copper(II) sulfate (CuSO4)Equipmentsmall-scale reaction surface electrolysis apparatus (9volt Batter with cap.)Experimental PagePlace one drop of each solution in the indicated place, and apply the leads of the electrolysis apparatus. Be sure to clean the leads between each experiment. Look carefully at the cathode (negative lead) and the anode (positive lead), and record your observations for each electrode in Table 38.1. BTB turns blue when OH- is present, BTB turns yellow when H+ is present, and Starch turns Blue/black when starch is present. These color changes tell you what products are being produced.

Experimental Data


Record your results in Table 38.1 or a table like it in your notebook. Assign each given half-reaction as occurring at the anode or the cathode.Questions for AnalysisUse what you learned in this lab to answer the following questions.1. Explain how you know that pure water does not conduct electricity and why it does not undergo

electrolysis.

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2. What do you observe when you apply the electrolysis apparatus to the water + sodium sulfate, Na2SO4. Why do experiments a and b give different results?

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3. Which reaction of the half-reactions below is the oxidation? Which is the reduction? Which is the cathode reaction? Which is the anode reaction?

Electrochemistry Unit Name: _________________________________ of 162H2O + 2e- H2(g) + 2OH-

H2O O2(g) + 2H+ + 2e –

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4. What do you observe when you electrolyze H2O + Na2SO4 + BTB? Given the electrode half-reactions in Question 3, what must be one product at the cathode (negative electrode)? What must be one product at the anode (positive electrode)?

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5. Add the half-reactions in Question 3 to obtain the net ionic equation. Simplify the result by adding together the OH- and H to get HOH, and then cancel out anything that appears on both sides of the equation.

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1. What do you observe when you electrolyze H2O + KI? How do the given half- reactions explain your observations? Add the half-reactions below to obtain the net ionic equation.

2H2O + 2e- H2(g) + 2OH-

2I- I2(g) + 2e-

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2. What is the effect of the starch added to the H2O + KI? What chemical species does starch detect?

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3. Recalling that Cl and I are in the same family, predict the half-reactions that occur when H2O + NaCl is electrolyzed. Cite evidence from your experiments that supports your theory. Repeat this for KBr.

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4. What happened to the color of the cathode as you electrolyzed CuSO4? What did you observe at the anode? Write half-reactions to explain.

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The Effect of Concentration (email)If an ionic compound is dissolved in water, it dissociates into ions and the resulting solution will conduct electricity. Dissolving solid sodium chloride in water releases ions according to the equation:

NaCl(s) Na+(aq) + Cl-(aq)

In this experiment, you will study the effect of increasing the concentration of an ionic compound on conductivity. Conductivity will be measured as concentration of the solution is gradually increased by the addition of concentrated NaCl drops. The same procedure will be used to investigate the effect of adding solutions with the same concentration (1.0 M), but different numbers of ions in their formulas: aluminum chloride, AlCl3, and calcium chloride, CaCl2. A computer-interfaced conductivity probe will be used to measure conductivity of the solution. Conductivity is measured in microsiemens (µS).

Figure 1

MATERIALSMacintosh or IBM-compatible computer 100-mL beakerSerial Box Interface or ULI distilled waterLogger Pro 1.0 M NaCl solutionVernier Conductivity Probe 1.0 M CaCl2 solutionring stand 1.0 M AlCl3 solution utility clamp stirring rod

PROCEDURE1. Obtain and wear goggles.2. Prepare the computer for data collection by opening “Exp 14” from the Chemistry with Computers

experiment files of Logger Pro. The vertical axis will have conductivity scaled from 0 to 2000 µS. The horizontal axis will have volume scaled from 0 to 8 drops.

3. Your experiment setup should look like Figure 1. The Conductivity Probe is already attached to the interface box and computer. It should be set on the 0-2000 µS position. Conductivity is measured in microsiemens (µS).

4. Add 70 mL of distilled water to a clean 100-mL beaker. Obtain a dropper bottle that contains 1.0 M NaCl solution.

5. Before adding any drops of solution:• Click Collect .

Electrochemistry Unit Name: _________________________________ of 16• Carefully raise the beaker and its contents up around the conductivity probe until the hole near the

probe end is completely submerged in the solution being tested. Important: Since the two electrodes are positioned on either side of the hole, this part of the probe must be completely submerged as shown in Figure 1.

• Monitor the conductivity of the distilled water until the conductivity reading stabilizes.• Click Keep , and then lower the beaker away from the probe. Type “0” in the edit box (for 0 drops

added). Press the ENTER key to store this data pair. This gives the conductivity of the water before any salt solution is added.

6. You are now ready to begin adding salt solution.• Add 1 drop of NaCl solution to the distilled water. Stir to ensure thorough mixing.• Raise the beaker until the hole near the probe end is completely submerged in the solution. Swirl the

solution briefly.• Monitor the conductivity of the solution until the reading stabilizes.• Click Keep , and then lower the beaker away from the probe. Type “1” (the total drops added) in the

edit box and press ENTER.7. Repeat the Step 6 procedure, entering “2” this time.8. Continue this procedure, adding 1-drop portions of NaCl solution, measuring conductivity, and entering

the total number of drops added—until a total of 8 drops have been added. 9. Dispose of the beaker contents as directed by your teacher.

10. Prepare the computer for data collection. From the Data menu, choose Store Latest Run. This stores the data so it can be used later, but it will be still be displayed while you do your second and third trials.

11. Repeat Steps 4-10, this time using 1.0 M AlCl3 solution in place of 1.0 M NaCl solution. 12. Repeat Steps 4-9, this time using 1.0 M CaCl2 solution.13. Click on the Linear Regression button, . Be sure all three data runs are checked, then click OK . A

best-fit linear regression line will be shown for each of your three runs. In your data table, record the value of the slope, m, for each of the three solutions. (The linear regression statistics are displayed in a floating box for each of the data sets.)

14. To print a graph of concentration vs. volume showing all three data runs:• Label all three curves by choosing Make Annotation from the Analyze menu, and typing “sodium

chloride” (or “aluminum chloride”, or “calcium chloride”) in the edit box. Then drag each box to a position near its respective curve.

• Print a copy of the Graph window, with all three data sets and the regression lines displayed. Enter your name(s) and the number of copies of the graph you want.

Electrochemistry Unit Name: _________________________________ of 16DATA TABLE

Solution Slope, m

1.0 M NaCl _________

1.0 M AlCl3 _________

1.0 M CaCl2 _________

PROCESSING THE DATA1. Describe the appearance of each of the three curves on your graph.

2. Describe the change in conductivity as the concentration of the NaCl solution was increased by the addition of NaCl drops. What kind of mathematical relationship does there appear to be between conductivity and concentration?

3. Write a chemical equation for the dissociation of NaCl, AlCl3, and CaCl2 in water.

4. Which graph had the largest slope value? The smallest? Since all solutions had the same original concentration (1.0 M), what accounts for the difference in the slope of the three plots? Explain.

5. Write a conclusion paragraph.


Electrochemistry: Voltaic Cells (Record your answers on this file and email to me.)

In electrochemistry, a voltaic cell is a specially prepared system in which an oxidation-reduction reaction occurs spontaneously. This spontaneous reaction produces an easily measured electrical potential. Voltaic cells have a variety of uses.

In this experiment, you will prepare a variety of semi-microscale voltaic cells in a 24-well test plate. A voltaic cell is constructed by using two metal electrodes and solutions of their respective salts (the electrolyte component of the cell) with known molar concentrations. In Parts I and II of this experiment, you will use a Voltage Probe to measure the potential of a voltaic cell with copper and lead electrodes. You will then test two voltaic cells that have unknown metal electrodes and, through careful measurements of the cell potentials, identify the unknown metals. In Part III of the experiment, you will measure the potential of a special type of voltaic cell called a concentration cell. In the first concentration cell, you will observe how a voltaic cell can maintain a spontaneous redox reaction with identical copper metal electrodes, but different electrolyte concentrations. You will then measure the potential of a second concentration cell and use the Nernst equation to calculate the solubility product constant, Ksp, for lead iodide, PbI2.

Figure 1

OBJECTIVESIn this experiment, you will

Prepare a Cu-Pb voltaic cell and measure its potential. Test two voltaic cells that use unknown metal electrodes and identify the metals. Prepare a copper concentration cell and measure its potential. Prepare a lead concentration cell and measure its potential. Use the Nernst equation to calculate the Ksp of PbI2.

Electrochemistry Unit Name: _________________________________ of 16MATERIALSVernier computer interface 0.10 M copper (II) nitrate, Cu(NO3)2, solutioncomputer 0.10 M lead (II) nitrate, Pb(NO3)2, solutionVoltage Probe 1.0 M copper (II) sulfate, CuSO4, solutionthree 10 mL graduated cylinders 0.050 M potassium iodide, KI, solution24-well test plate 1 M potassium nitrate, KNO3, solutionstring 0.10 M X nitrate solutionCu and Pb electrodes 0.10 M Y nitrate solutiontwo unknown electrodes, labeled X and Y steel wool150 mL beaker plastic Beral pipets

PRE-LAB EXERCISEUse the table of standard reduction potentials in your text, or another approved reference, to complete the following table. An example is provided.

Electrodes Half-reactions E° E°cell

ZnCu

Zn(s) → Zn2+ + 2e–

Cu2+ + 2e– → Cu(s)+0.76 V+0.34 V

+1.10 V

Cu

Pb

Pb

Ag

Pb

Mg

Pb

Zn

PROCEDUREPart I Determine the Eo for a Cu-Pb Voltaic Cell

1. Obtain and wear goggles.

2. Use a 24-well test plate as your voltaic cell. Use Beral pipets to transfer small amounts of 0.10 M Cu(NO3)2 and 0.10 M Pb(NO3)2 solution to two neighboring wells in the test plate. CAUTION: Handle these solutions with care. If a spill occurs, ask your instructor how to clean up safely.

3. Obtain one Cu and one Pb metal strip to act as electrodes. Polish each strip with steel wool. Place the Cu strip in the well of Cu(NO3)2 solution and place the Pb strip in the well of Pb(NO)3 solution. These are the half cells of your Cu-Pb voltaic cell.

4. Make a salt bridge by soaking a short length of string in a beaker than contains a small amount of 1 M KNO3 solution. Connect the Cu and Pb half cells with the string.

5. Connect a Voltage Probe to Channel 1 of the Vernier computer interface. Connect the interface to the computer with the proper cable.

Electrochemistry Unit Name: _________________________________ of 166. Start the Logger Pro program on your computer. Open the file “20 Electrochemistry” from the Advanced

Chemistry with Vernier folder.

7. Measure the potential of the Cu-Pb voltaic cell. Complete the steps quickly to get the best data.a. Click to start data collection.b. Connect the leads from the Voltage Probe to the Cu and Pb electrodes to get a positive potential

reading. Click immediately after making the connection with the Voltage Probe.c. Remove both electrodes from the solutions. Clean and polish each electrode.d. Put the Cu and Pb electrodes back in place to set up the voltaic cell. Connect the Voltage Probe to the

electrodes, as before. Click immediately after making the connection with the Voltage Probe. e. Remove the electrodes. Clean and polish each electrode again.f. Set up the voltaic cell a third, and final, time. Click immediately after making the connection

with the Voltage Probe. Click to end the data collection.g. Click the Statistics button, . Record the mean in your data table as the average potential. Close the

statistics box on the graph screen by clicking the X in the corner of the box.

Part II Determine the Eo for Two Voltaic Cells Using Pb and Unknown Metals

8. Obtain a small amount of the unknown electrolyte solution labeled “0.10 M X” and the corresponding metal strip, X.

9. Use a Beral pipet to transfer a small amount of 0.10 M X solution to a well adjacent to the 0.10 M Pb(NO3)2 solution in the test plate.

10. Make a new salt bridge by soaking a short length of string in the beaker of 1 M KNO3 solution. Connect the X and Pb half cells with the string.

11. Measure the potential of the X-Pb voltaic cell. Complete this step quickly.a. Click to start data collection.b. Connect the leads from the Voltage Probe to the X and Pb electrodes to get a positive potential

reading. Click immediately after making the connection with the Voltage Probe.c. Remove both electrodes from the solutions. Clean and polish each electrode.d. Set up the voltaic cell again. Connect the Voltage Probe as before. Click immediately after

making the connection with the Voltage Probe. e. Remove the electrodes. Clean and polish each electrode again.f. Test the voltaic cell a third time. Click immediately after making the connection with the

Voltage Probe.g. Click to end data collection.h. Click the Statistics button, . Record the mean in your data table as the average potential and then

close the statistics box on the graph screen by clicking the X in the corner of the box.

12. Repeat Steps 8–11 using the unknown and its corresponding electrolyte solution labeled “Y”.

Electrochemistry Unit Name: _________________________________ of 16Part III Prepare and Test Two Concentration Cells

13. Set up and test a copper concentration cell.a. Prepare 20 mL of 0.050 M CuSO4 solution by mixing 1 mL of 1.0 M CuSO4 solution with 19 mL of

distilled water.b. Set up a concentration cell in two wells of the 24-well test plate by adding 5 mL of 0.050 M CuSO4

solution to one well and 5 mL of 1.0 M CuSO4 solution to a neighboring well. Use Cu metal electrodes in each well. Use a KNO3-soaked string as the salt bridge, as in Parts I and II.

c. Click to start data collection.i. Test and record the potential of the concentration cell in the same manner that you tested the voltaic

cells in Parts I and II.

14. Set up a concentration cell to determine the solubility product constant, Ksp, of PbI2.a. Prepare 10 mL of 0.050 M Pb(NO3)2 solution by mixing 5 mL of 0.10 M Pb(NO3)2 solution with 5 mL

of distilled water.b. Mix 9 mL of 0.050 M KI solution with 3 mL of 0.050 M Pb(NO3)2 solution in a small beaker. In this

reaction, most of the Pb2+ and I– will form the precipitate PbI2, but a small amount of the ions will remain dissolved.

c. Set up the half cells in neighboring wells of the 24-well test plate. Place 5 mL of 0.050 M Pb(NO3)2 solution in one half cell, and 5 mL of the PbI2 mixture, from the small beaker, into an adjacent half cell. Use Pb electrodes in each half cell. Use a KNO3-soaked string as the salt bridge.

j. Test and record the potential of the cell in the same manner that you tested the voltaic cells and the copper concentration cell.

15. Discard the electrodes and the electrolyte solutions as directed. Rinse and clean the 24-well plate. CAUTION: Handle these solutions with care. If a spill occurs, ask your instructor how to clean up safely.

DATA TABLE

Results of Parts I and II Cu/Pb X/Pb Y/Pb

Average cell potential (V)

Results of Part III Cu concentration Pb/PbI2

Average cell potential (V)

Electrochemistry Unit Name: _________________________________ of 16DATA ANALYSIS1. (Part I) Compare the average cell potential, for your Cu/Pb cell, with the E°

cell that you calculated in the pre-lab exercise. Explain why your cell potential is different from the text value.

2. (Part II) The unknown metals X and Y were either magnesium, silver, or zinc. Use the text value for the reduction potential of Pb and the measured cell potentials for the unknowns to identify X and Y.

3. (Part III) Use the Nernst equation to calculate the theoretical value of E of the copper-concentration cell and compare this value with the cell potential that you measured.

4. (Part III) Use the Nernst equation and the information that you collected about the Pb/PbI2 cell to complete the following calculations.a. Use the cell potential for the Pb-PbI2 cell and the known [Pb2+] to calculate the [Pb2+] in equilibrium

with PbI2.b. Use the original diluted [Pb2+] and [I–] to calculate the [I–] in solution.c. Use your data to calculate the Ksp of PbI2

k. The accepted value of the Ksp of PbI2 is 9.8 × 10–9. How does your experimental Ksp of PbI2 compare with the accepted value?

5. Write a paragraph about what you learned like Nernst equations, the difference in the two different type of cells, etc.


Electroplating (Email)

In this experiment, you will conduct, observe, and measure the process of electroplating. This process is used to deposit a layer of metal, such as chromium, copper, or gold, onto another metal. As a commercial process, electroplated coatings are used to improve appearance, resist corrosion, or improve hardness of metallic surfaces. This experiment describes one method of producing a copper coating on a brass key or other suitable metallic object.

You will prepare an electrochemical cell by using a copper strip as the cathode (positive terminal) and a brass key as the anode (negative terminal). The electrodes are immersed in a solution containing acidified copper (II) sulfate. As you apply a potential to the electrodes, you will be effectively transferring Cu atoms from the anode to the surface of the brass key.

In this experiment, you will use one application of Faraday’s law, stated in equation form below.

I is the current in amperes; t is the time that the current is applied, in seconds; MM is the molar mass of the element that is deposited; n is the number of moles of electrons/mol; and 96,500 is ₣, the Faraday constant.

Figure 1

OBJECTIVESIn this experiment, you will

Prepare and operate an electrochemical cell to plate copper onto a brass surface. Measure the amount of copper that was deposited in the electroplating process. Calculate the amount of energy used to complete the electroplating process.

Electrochemistry Unit Name: _________________________________ of 16MATERIALSVernier computer interface electrolyte solution (CuSO4 in H2SO4)computer vinegarVernier Current Probe 1 cm × 10 cm copper strip 1.5 volt DC power supply brass keyfour connecting wires with alligator clips solid sodium chloride, NaClsteel wool balance, 0.001 g precisiontwo 250 mL beakers bare copper wire, 20-22 gaugedistilled water

PROCEDURE1. Obtain and wear goggles.

2. Use steel wool to clean a brass key and a strip of copper, which will be the electrodes of the electrochemical cell.

3. Mix 3 g of NaCl with 15 mL of vinegar in a 250 mL beaker. Wash the key and the copper strip in this salt-vinegar solution. Rinse the key and copper strip with distilled water and dry each metal piece.

4. Use an analytical balance to determine the mass of the key and the mass of the copper strip. Record these two masses in your data table.

5. Fill a 250 mL beaker about ¾ full with the electrolyte solution. CAUTION: The electrolyte solution in this experiment is prepared in H2SO4 and should be handled with care.

6. Attach a 7 cm length of bare copper wire to the brass key to act as a handle. Connect the wire to the alligator clip for the anode, so that the key will be completely immersed in the electrolyte solution but the alligator clip will not be immersed. Connect the copper strip to the positive lead on the cell. Important: You will not place the electrodes in the cell until Step 10.

7. Obtain a DC power supply and a Vernier Current Probe. Use connecting wires, with alligator clips, to connect the DC power supply, Current Probe, and the electrodes. See Figure 1 for the proper setup of the wiring. Important: You will not place the electrodes in the cell until Step 10.

8. Connect the Current Probe to Channel 1 of the Vernier computer interface. Connect the interface to your computer with the proper cable.

9. Start the Logger Pro program on your computer. Open the file “21 Electroplating” from the Advanced Chemistry with Vernier folder.

10. Place the key and the copper strip into the electrolyte solution in the cell. Make sure that the key is completely immersed in the solution, and keep the two electrodes as far apart as possible.

11. Turn on the DC power source and check the sensor readings. The initial current should be in the 0.2–0.6 A range. If the current is not in this range, check with your instructor before proceeding.

12. Click to begin the data collection. Observe the electrolysis. Note the slow deposition of copper on the surface of the key. The data collection will run for 30 minutes.

13. When the data collection is complete, turn off the DC power source and carefully remove the copper strip and key from the electrolyte solution. Rinse the two metals with distilled water. Dry the copper strip and key so as not to remove copper.

14. Analyze the graph of your data to determine the average current that was applied during the electrolysis. Click the Statistics button, . Record the mean in your data table as the average current.

Electrochemistry Unit Name: _________________________________ of 1615. Measure and record the mass of the dry copper strip and key.

16. Discard the electrolyte solution and take care of the electrochemical cell as directed by your instructor.

DATA TABLE email the data tables and question answers to me

Initial mass of copper electrode (g)

Final mass of copper electrode (g)

Initial mass of key (g)

Final mass of key (g)

Average current (A)

Time of current application (s)

DATA ANALYSIS1. Calculate the number of coulombs of charge passed through the electrolytic cell.

2. Calculate the theoretical number of moles of copper that should have plated out onto the brass key.

3. Calculate the actual number of moles of copper that plated out.

4. Calculate the percent yield of copper.

5. Suggest the sources of error in your experiment.

6. Write the oxidation and reduction half-reactions for this process.

7. Write a conclusion paragraph.


Both Copper into Gold: The Alchemist’s Dream Take a picture of your Penny and write a summary of the reaction email to me.

A copper penny is placed in an evaporating dish and heated with a mixture. It turns silver. The penny is then heated on a hot plate, and it suddenly turns gold.

Procedure:Wear safety goggles, gloves, and a face shield. Note step 10: special disposal procedure.

1. Place approximately 2 g of mossy zinc (zinc pellets) in an evaporating dish.

2. Add enough NaOH 1M solution to cover the zinc and fill the dish to cover the penny.

3. Place the dish on a hot plate and heat until the solution is near boiling.

4. Prepare a copper penny by cleaning it thoroughly with a light abrasive (steel wool pads work well).

5. Using crucible tongs or tweezers, place the cleaned penny in the mixture in the dish.

6. Leave the penny in the dish for 3-4 mm. You will be able to tell when the silver coating is complete.

7. Remove the penny, rinse it, and blot dry with paper towels. (Do not rub as this would remove the zinc particles).

8. Using crucible tongs or tweezers, place the coated penny on the hot plate. The gold color appears almost immediately.

9. When the gold color forms, remove the coin, rinse it, and blot dry with paper towels.

Reactions

1. The first reaction is the plating of the copper with zinc: Zinc reacts with sodium hydroxide to form a sodium zincate, [Zn(OH)3(H2O)]Na+. This reaction gives the silver color to the penny. Write a balanced equation for this reaction in your conclusion paragraph.

2. The second reaction is the formation of the brass alloy. This alloy gives the penny the gold color. Heat causes a fusion of the zinc and copper.

3. Research ways that electroplating is use in modern industry.

electric current: the movement of charged...

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