Electrochemistry Unit Name: ______Page 16 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.
Purpose
In 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.
Materials
Small-scale pipets of the following solutions:
water (H2O) sodium sulfate (Na2SO4)
bromthymol blue (BTB) potassium iodide (KI)
starch sodium chloride (NaCl)
potassium bromide (KBr) copper(II) sulfate (CuSO4)
Equipment
small-scale reaction surface electrolysis apparatus (9volt Batter with cap.)
Experimental Page
Place 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 Analysis
Use 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?
2H2O + 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
MATERIALS
Macintosh or IBM-compatible computer / 100-mL beakerSerial Box Interface or ULI / distilled water
Logger Pro / 1.0 M NaCl solution
Vernier Conductivity Probe / 1.0 M CaCl2 solution
ring stand / 1.0 M AlCl3 solution
utility clamp / stirring rod
PROCEDURE
1. 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 .
• 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 , 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 , 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 . 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.
DATA TABLE
1.0 M NaCl / ______
1.0 M AlCl3 / ______
1.0 M CaCl2 / ______
PROCESSING THE DATA
1. 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
OBJECTIVES
In 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.
MATERIALS
computer / 0.10 M lead (II) nitrate, Pb(NO3)2, solution
Voltage Probe / 1.0 M copper (II) sulfate, CuSO4, solution
three 10 mL graduated cylinders / 0.050 M potassium iodide, KI, solution
24-well test plate / 1 M potassium nitrate, KNO3, solution
string / 0.10 M X nitrate solution
Cu and Pb electrodes / 0.10 M Y nitrate solution
two unknown electrodes, labeled X and Y / steel wool
150 mL beaker / plastic Beral pipets
PRE-LAB EXERCISE
Use 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°cellZn
Cu / 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
PROCEDURE
Part 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.
6. 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.