Experiment O2(g) H2(g) 33 Electrolytic Cells, Avogadro’s Number A 1:2 mole (and volume) ratio of oxygen (left) to hydrogen (right) is produced from the electrolysis of water • To identify the reactions occurring at the anode and cathode during the electrolysis of various aqueous salt solutions • To determine Avogadro’s number and the Faraday constant Objectives The following techniques are used in the Experimental Procedure: Techniques Electrolysis processes are very important in achieving high standards of living The industrial production of metals such as aluminum and magnesium and nonmetals such as chlorine and uorine occurs in electrolytic cells The highly re ned copper metal required for electrical wiring is obtained through an electroplating process In an electrolytic cell, the input of an electric current causes an otherwise nonspontaneous oxidation–reduction reaction, a nonspontaneous transfer of electrons, to occur For example, sodium metal, a very active metal, and chlorine gas, a very toxic gas, are prepared industrially by the electrolysis of molten sodium chloride Electrical energy is supplied to a molten NaCl system (Figure 33.1) by a direct current (dc) power source (set at an appropriate voltage) across the electrodes of an electrolytic cell The electrical energy causes the reduction of the sodium ion, Naϩ, at the cathode and oxidation of the chloride ion, ClϪ, at the anode Because cations migrate to the cathode and anions migrate to the anode, the cathode is the negative electrode (opposite charges attract), and the anode is designated the positive electrode.1 Introduction cathode (Ϫ) reaction: anode (ϩ) reaction: Electrolysis: Use of electrical energy to cause a chemical reaction to occur Electroplating: the use of electrical current to deposit a metal onto an electrode Electrolytic cell: an apparatus used for an electrolysis reaction Naϩ(l) ϩ eϪ l Na(l) ClϪ(l) l Cl2(g) ϩ eϪ Electrolysis reactions also occur in aqueous solutions For example, in the electrolysis of an aqueous copper(II) bromide, CuBr2, solution, copper(II) ions, Cu2ϩ, are reduced at the cathode and bromide ions, BrϪ, are oxidized at the anode (Figure 33.2, page 364) cathode (Ϫ) reaction: Cu2ϩ(aq) ϩ eϪ l Cu(s) anode (ϩ) reaction: BrϪ(aq) l Br2(l) ϩ eϪ 2ϩ cell reaction: Cu (aq) ϩ BrϪ(aq) l Cu(s) ϩ Br2(l) Note that the cathode is “Ϫ” in an electrolytic cell but “ϩ” in a galvanic cell; the anode is “ϩ” in an electrolytic cell but “Ϫ” in a galvanic cell See Experiment 32 Figure 33.1 Schematic diagram of the electrolysis of molten sodium chloride Experiment 33 363 In an aqueous solution, however, the reduction of water at the cathode (the negative electrode) and the oxidation of water at the anode (the positive electrode) are also possible reactions Electrolysis of Aqueous Solution cathode (Ϫ) reaction for water: anode (ϩ) reaction for water: + H2O(l) ϩ eϪ l H2(g) ϩ OHϪ(aq) ϩ Ϫ H2O(l) l O2(g) ϩ H (aq) ϩ e (33.1) (33.2) – If the reduction of water occurs at the cathode, hydrogen gas is evolved and the solution near the cathode becomes basic as a result of the production of hydroxide ion If oxidation of water occurs at the anode, oxygen gas is evolved and the solution near the anode becomes acidic The acidity (or basicity) near the respective electrodes can be detected with pH paper or another acid–base indicator When two or more competing reduction reactions are possible at the cathode, the reaction that occurs most easily (the one with the higher reduction potential) is the one that usually occurs Conversely, for two or more competing oxidation reactions at the anode, the reaction that takes place most easily (the one with the higher oxidation potential or the lower reduction potential) is the one that usually occurs In the electrolysis of the aqueous copper(II) bromide solution, Cu2ϩ Figure 33.2 Electrolysis of a copper(II) has a higher reduction potential than H2O and is therefore preferentially bromide solution; the anode is on the left reduced at the cathode; BrϪ has a greater tendency to be oxidized than and the cathode is on the right water, and so BrϪ is oxidized at the anode In Part A of this experiment, a number of aqueous salt solutions using different electrodes are electrolyzed The anode and cathode are identi ed, and the products that are formed at each electrode are also identi ed Avogadro’s Number and the Faraday Constant In Part B, a quantitative investigation of the electrolytic oxidation of copper metal is used to determine Avogadro’s number and the Faraday constant: Cu(s) l Cu2ϩ(aq) ϩ eϪ faraday ϭ mol eϪ ϭ 96,485 coulombs Two moles of electrons (or faradays) are released for each mole of Cu(s) oxidized; therefore, a mass measurement of the copper anode before and after the electrolysis determines the moles of copper that are oxidized This in turn is used to calculate the moles of electrons that pass through the cell: moles of electrons ϭ mass Cu ϫ Coulomb: SI base unit for electrical charge (33.3) mol Cu mol eϪ ϫ mol Cu 63.54 g (33.4) The actual number of electrons that pass through the cell is calculated from the electrical current, measured in amperes (ϭ coulombs/second), that passes through the cell for a recorded time period (seconds) The total charge (coulombs, C) that passes through the cell is number of coulombs ϭ coulombs ϫ seconds second (33.5) As the charge of one electron equals 1.60 ϫ 10Ϫ19 C, the number of electrons that pass through the cell can be calculated: electron (33.6) 1.60 ϫ 10Ϫ19 C Therefore, since the number of electrons (equation 33.6) and the moles of electrons (equation 33.4) can be separately determined, Avogadro’s number is calculated as number of electrons ϭ number of coulombs ϫ Avogadro’s number ϭ 364 Electrolytic Cells, Avogadro’s Number number of electrons mole of electrons (33.7) In addition, the number of coulombs (equation 33.5) per mole of electrons (equation 33.4) equals the Faraday constant With the available data, the Faraday constant can also be calculated: Faraday constant ϭ number of coulombs mole of electrons (33.8) Procedure Overview: The products that result from the electrolysis of various salt solutions are observed and identi ed; these are qualitative measurements An experimental setup is designed to measure quantitatively the ow of current and consequent changes in mass of the electrodes in an electrolytic cell; from these data, experimental constants are calculated The electrolysis apparatus may be designed differently than the one described in this experiment Ask your instructor Set up the electrolysis apparatus Connect two wire leads (different colors) attached to alligator clips to a direct current (dc) power supply.2 Clean and mount the glass U-tube on a ring stand (see Figure 33.3) Connect the alligator clips to the corresponding electrodes, listed in Table 33.1 Electrolyze the solutions Fill the U-tube three-fourths full with solution from Table 33.1 Insert the corresponding electrodes into the solution and electrolyze for ϳ5 minutes During the electrolysis, watch for any evidence of a reaction in the anode and cathode chambers Experimental Procedure A Electrolysis of Aqueous Salt Solutions • Does the pH of the solution change at each electrode? Test each chamber with litmus or pH paper.3 Compare the color with a pH test on the original solution • Is a gas evolved at either or both electrodes? • Look closely at each electrode Is a metal depositing on the electrode or is the metal electrode slowly disappearing? Account for your observations Write the equations for the reactions occurring at the anode and cathode and for the cell reaction Repeat for solutions 2–5 Disposal: Discard the salt solutions into the Waste Salts container CLEANUP: Rinse the U-tube twice with tap water and twice with deionized water before preparing the next solution for electrolysis Discard each rinse in the sink Table 33.1 Electrolytic Cells for Study Solution No Solution* Electrodes (Cathode and Anode) g NaCl/100 mL g NaBr/100 mL g KI/100 mL 0.1 M CuSO4 0.1 M CuSO4 Carbon (graphite) Carbon (graphite) Carbon (graphite) Carbon (graphite) Polished copper metal strips *Try other solutions and electrodes as suggested by your laboratory instructor Figure 33.3 Electrolysis apparatus The dc power supply can be a 9-V transistor battery Several drops of universal indicator can be added to the solution in both chambers to detect pH changes Experiment 33 365 Set up the apparatus Refer to Figure 33.4 The U-tube from Part A can again be used The dc power supply must provide 3–5 volts (two or three ashlight batteries in series or a lantern battery); the ammeter or multimeter must read from 0.2 to 1.0 A Polish two copper metal strips (to be used as the electrodes) with steel wool or sandpaper Brie y dip each electrode (use the fume hood) into M HNO3 (Caution: not allow skin contact) for further cleaning, and then rinse with deionized water Add 100 mL of 1.0 M CuSO4 (in 0.1 M H2SO4) to the 150-mL beaker (or ll the U-tube) Set the electrodes Rinse the electrodes with ethanol if available When dry, label the two electrodes because the mass of each will be determined before and after the electrolysis Measure the mass (ע0.001 g, preferably ע0.0001 g) of each labeled electrode The copper electrode with the lesser mass is to serve as the anode (ϩ terminal), and the other is to serve as the cathode (Ϫ terminal) for the electrolytic cell Connect the cathode (through the variable resistor and ammeter/multimeter) to the negative terminal of the dc power supply Before electrolysis begins, obtain your instructor’s approval of the complete apparatus Electrolyze the CuSO4 solution Adjust the variable resistance to its maximum value.4 Be ready to start timing (a stopwatch is ideal) Attach the anode to the positive terminal of the dc power supply and START TIME During the electrolysis, not move the electrodes; this changes current ow Adjust the current with the variable resistor to about 0.5 A and, periodically during the course of the electrolysis, readjust the current to 0.5 A.5 Discontinue the electrolysis after 20–30 minutes Record the exact time (minutes and seconds) of the electrolysis process Dry and measure the mass Carefully remove the electrodes (be careful not to loosen the electroplated copper metal from the cathode); carefully dip each electrode into a 400-mL beaker of deionized water to rinse the electrodes (followed by ethanol if available) Air-dry, measure the mass (ע0.001 g, preferably ע0.0001 g) of each electrode, and record Repeat the electrolysis If time allows, repeat Part B using the same copper electrodes (with new mass measurements!) and 1.0 M CuSO4 solution B Determination of Avogadro’s Number and the Faraday Constant Disposal: Discard the copper(II) sulfate solution into the Waste Salts container CLEANUP: Rinse the beaker or U-tube twice with tap water and twice with deionized water Discard each rinse as directed by your instructor Figure 33.4 Setup for determining Avogadro’s number and the Faraday constant The Next Step Electroplating of metals such as nickel, chromium, silver, and copper is a common industrial process Research a speci c process and design an apparatus and procedure for depositing quantitative amounts of metal to a cathode If a variable resistor is unavailable, record the current at 1-minute intervals and then calculate an average current over the entire electrolysis time period If the current is greater or less than 0.5 A, vary the time of electrolysis proportionally 366 Electrolytic Cells, Avogadro’s Number Experiment 33 Prelaboratory Assignment Electrolytic Cells, Avogadro’s Number Date Lab Sec Name Desk No The standard reduction potential for the Cu2ϩ/Cu redox couple is ϩ0.34 V; that for H2O/H2, OHϪ at a pH of is Ϫ0.41 V For the electrolysis of a neutral 1.0 M CuSO4 solution, write the equation for the half-reaction occurring at the cathode at standard conditions In an electrolytic cell, a oxidation occurs at the (name of electrode) _ b the cathode is the (sign) electrode _ c cations ow toward the (name of electrode) _ d electrons ow from the (name of electrode) to (name of electrode) _ e the anode should be connected to the (positive/negative) terminal of the dc power supply _ _ Experimental Procedure, Part A.2 Describe the proper technique for testing the pH of a solution with litmus or pH paper a Identify a chemical test(s) to determine if water is oxidized at the anode in an electrolytic cell b Similarly, identify a chemical test(s) to determine if water is reduced at the cathode in an electrolytic cell Very pure copper metal is produced by the electrolytic re ning of blister (impure) copper In the cell at right, label the anode, the cathode, and the polarity (ϩ, Ϫ) of each Experiment 33 367 a When a solution of sodium sulfate, Na2SO4, adjusted to a pH of is electrolyzed, red litmus remains red and a gas is evolved in the anodic chamber In the cathodic chamber, red litmus turns blue and a gas is also evolved (i) Write a balanced equation for the half-reaction occurring at the anode (ii) Write a balanced equation for the half-reaction occurring at the cathode b When a solution of nickel(II) sulfate adjusted to a pH of is electrolyzed, the green color of the solution becomes less intense in the cathodic chamber and gas bubbles are detected in the anodic chamber (i) Write a balanced equation for the half-reaction occurring at the anode (ii) Write a balanced equation for the half-reaction occurring at the cathode A 1.0 M CuSO4 solution was electrolyzed for 28 minutes and 22 seconds using copper electrodes The average current owing through the cell over the time period was 0.622 A The mass of the copper anode before the electrolysis was 2.4852 g and 2.1335 g afterwards a Calculate the number of moles of copper oxidized and moles of electrons that passed through the cell (see equation 33.4) b Calculate the total charge (coulombs) that passed through the cell (see equation 33.5) c How many electrons passed through the cell during the 28.0 minute, 22 second time period (see equation 33.6)? d From the data, calculate Avogadro’s number (see equation 33.7) 368 Electrolytic Cells, Avogadro’s Number Experiment 33 Report Sheet Electrolytic Cells, Avogadro’s Number Date Lab Sec Name Desk No A Electrolysis of Aqueous Salt Solutions Solution Electrodes Litmus Test NaCl C(gr) Gas Evolved? Balanced Equations for Reactions Anode Cathode _ Cell NaBr C(gr) Anode Cathode _ Cell KI C(gr) Anode Cathode _ Cell CuSO4 C(gr) Anode Cathode _ Cell CuSO4 Cu(s) Anode Cathode _ Cell Experiment 33 369 B Determination of Avogadro’s Number and the Faraday Constant Trial Trial Initial mass of copper anode (g) Initial mass of copper cathode (g) Data Instructor’s approval of apparatus Time of electrolysis (s) Current (or average current) (A) Final mass of copper anode (g) Final mass of copper cathode (g) Mass of copper oxidized at anode (g) Moles of copper oxidized (mol) Moles of electrons transferred (mol e ) Coulombs passed through cell (C) Electrons passed through cell (e ) Avogadro’s number (eϪ/mol eϪ) Data Analysis Ϫ Ϫ Average value of Avogadro’s number Literature value of Avogadro’s number Percent error 10 Faraday constant (C/mol eϪ) Ϫ Appendix B 11 Average Faraday constant (C/mol e ) 12 Literature value of Faraday constant 13 Percent error Appendix B Laboratory Questions Circle the questions that have been assigned Part A.2 If zinc electrodes are used instead of the graphite electrodes, the reaction occurring at the anode may be different, but the reaction occurring at the cathode would remain unchanged Explain Part A.2 Nitrate ions, NO3Ϫ, being anions, migrate to the anode in an electrolytic cell Explain why you would expect water rather than nitrate ions to be oxidized at the anode Hint: Consider the oxidation state of nitrogen in the nitrate ion Part B Repeat the calculation of Avogadro’s number, using the mass gain of the cathode instead of the mass loss of the anode Account for any difference in the calculated values Part B.2 If the current is recorded as being less than it actually is, would Avogadro’s number be calculated as too high or too low, or would it be unaffected? Explain Part B.4 Because of an impure copper anode (see Prelaboratory Assignment question 5), the measured mass loss is greater than the actual mass of copper oxidized As a result, will Avogadro’s number be calculated as too high or too low? Explain *6 The electrolytic re ning of copper involves the oxidaton of impure copper containing such metals as iron and nickel (oxidized to copper(II), iron(II), and nickel(II) ions) at the anode and then reduction of the copper(II) ion to copper metal at the cathode Explain why the iron(II) and nickel(II) ions are not deposited on the cathode 370 Electrolytic Cells, Avogadro’s Number