Bài giảng hoá phân tích galvanic cells

Experiment 32 Galvanic Cells, the Nernst Equation Copper metal spontaneously oxidizes to copper(II) ion in a solution containing silver ion Silver metal crystals form on the surface of the copper metal • To measure the relative reduction potentials for a number of redox couples • To develop an understanding of the movement of electrons, anions, and cations in a galvanic cell • To study factors affecting cell potentials • To estimate the concentration of ions in solution using the Nernst equation Objectives The following techniques are used in the Experimental Procedure: Techniques Electrolyic cells are of two types, galvanic and electrolysis, both employing the principle of oxidation–reduction (redox) reactions In galvanic (or voltaic) cells (this experiment), redox reactions occur spontaneously as is common with all portable batteries of which we are very familiar Electric cars, ashlights, watches, and power tools operate because of a speci c spontaneous redox reaction Electrolysis cells (Experiment 33) are driven by nonspontaneous redox reactions, reactions that require energy to occur The recharging of batteries, electroplating and re ning of metals, and generation of various gases all require the use of energy to cause the redox reaction to proceed Experimentally, when copper wire is placed into a silver ion solution (see opening photo), copper atoms spontaneously lose electrons (copper atoms are oxidized) to the silver ions (which are reduced) Silver ions migrate to the copper atoms to pick up electrons and form silver atoms at the copper metal–solution interface; the copper ions that form then move into the solution away from the interface The overall reaction that occurs at the interface is: Introduction Cu(s) ϩ Agϩ(aq) l Ag(s) ϩ Cu2ϩ(aq) (32.1) This redox reaction can be divided into an oxidation and a reduction half-reaction Each half-reaction, called a redox couple, consists of the reduced state and the oxidized state of the substance: Cu(s) l Cu2ϩ(aq) ϩ eϪ ϩ Ϫ Ag (aq) ϩ e l Ag(s) Interface: the boundary between two phases; in this case, the boundary that separates the solid metal from the aqueous solution oxidation half-reaction (redox couple) (32.2) Redox couple: an oxidized and reduced form of an ion/substance appearing in a reduction or oxidation half-reaction, generally associated with galvanic cells reduction half-reaction (redox couple) (32.3) Experiment 32 351 Figure 32.1 Schematic diagram of a galvanic cell Half-cell: a part of the galvanic cell that hosts a redox couple External circuit: the movement of charge as electrons through a wire connecting the two half-cells, forming one-half of the electrical circuit in a galvanic cell Salt bridge: paper moistened with a salt solution, or an inverted tube containing a salt solution, that bridges two half-cells to complete the solution part of an electrical circuit Internal circuit: the movement of charge as ions through solution from one half-cell to the other, forming onehalf of the electrical circuit in a galvanic cell Cell Potentials 352 A galvanic cell is designed to take advantage of this spontaneous transfer of electrons Instead of electrons being transferred at the interface of the copper metal and the silver ions in solution, a galvanic cell separates the copper metal from the silver ions to force the electrons to pass externally through a wire, an external circuit Figure 32.1 is a schematic diagram of a galvanic cell setup for these two redox couples The two redox couples are placed in separate compartments called half-cells Each half-cell consists of an electrode, usually the metal (reduced state) of the redox couple, and a solution containing the corresponding cation (oxidized state) of the redox couple The electrodes of the half-cells are connected by a wire through which the electrons ow, providing current for the external circuit A salt bridge that connects the two half-cells completes the construction of the galvanic cell (and the circuit) The salt bridge permits limited movement of ions from one half-cell to the other, the internal circuit, so that when the cell operates, electrical neutrality is maintained in each half-cell For example, when copper metal is oxidized to copper(II) ions in the Cu2ϩ/Cu half-cell, either NO3Ϫ anions must enter or copper(II) ions must leave the half-cell to maintain neutrality Similarly, when silver ions are reduced to form silver metal in its half-cell, either NO3Ϫ anions must leave or cations must enter its half-cell to maintain neutrality The electrode at which reduction occurs is called the cathode; the electrode at which oxidation occurs is called the anode Because oxidation releases electrons to the electrode to provide a current in the external circuit, the anode is designated the negative electrode in a galvanic cell The reduction process draws electrons from the circuit and supplies them to the ions in solution; the cathode is the positive electrode This sign designation allows us to distinguish the anode from the cathode in a galvanic cell Different metals, such as copper and silver, have different tendencies to oxidize; similarly, their ions have different tendencies to undergo reduction The cell potential of a galvanic cell is due to the difference in tendencies of the two metals to oxidize (lose electrons) or of their ions to reduce (gain electrons) Commonly, a measured reduction potential, the tendency for an ion (or molecule) to gain electrons, is the value used to identify the relative ease of reduction for a half-reaction A potentiometer or multimeter, placed in the external circuit between the two electrodes, measures the cell potential, Ecell, a value that represents the difference between the tendencies of the metal ions in their respective half-cells to undergo reduction (i.e., the difference between the reduction potentials of the two redox couples) Galvanic Cells, the Nernst Equation For the copper and silver redox couples, we can represent their reduction potentials as ECu2ϩ,Cu and EAgϩ,Ag, respectively The cell potential being the difference of the two reduction potentials is therefore Ecell ϭ EAgϩ,Ag Ϫ ECu2ϩ,Cu (32.4) Experimentally, silver ion has a greater tendency than copper ion does to be in the reduced (metallic) state; therefore, Agϩ has a greater (more positive) reduction potential Since the cell potential, Ecell, is measured as a positive value, EAgϩ,Ag is placed before ECu2ϩ,Cu in equation 32.4 The measured cell potential corresponds to the standard cell potential when the concentrations of all ions are mol/L and the temperature of the solutions is 25ЊC The standard reduction potential for the Agϩ(1 M)/Ag redox couple, EЊAgϩ,Ag, is ϩ0.80 V, and the standard reduction potential for the Cu2ϩ(1 M)/Cu redox couple, EЊCu2ϩ,Cu, is ϩ0.34 V Theoretically, a potentiometer (or multimeter) would show the difference between these two potentials, or, at standard conditions, EЊcell ϭ EЊAgϩ,Ag Ϫ EЊCu2ϩ,Cu ϭ ϩ0.80 V Ϫ (ϩ0.34 V) ϭ ϩ0.46V Silver jewelry is longer lasting than copper jewelry; therefore silver has a higher tendency to be in the reduced state, a higher reduction potential (32.5) Deviation from the theoretical value may be the result of surface activity at the electrodes or activity of the ions in solution In Part A of this experiment, several cells are “built” from a selection of redox couples and data are collected From an analysis of the data, the relative reduction potentials for the redox couples are determined and placed in an order of decreasing reduction potentials In Part B, the formations of the complex [Cu(NH3)4]2ϩ and the precipitate CuS are used to change the concentration of Cu2ϩ(aq) in the Cu2ϩ/Cu redox couple The observed changes in the cell potentials are interpreted Measure Cell Potentials The Nernst equation is applicable to redox systems that are not at standard conditions, most often when the concentrations of the ions in solution are not mol/L At 25ЊC, the measured cell potential, Ecell, is related to EЊcell and ionic concentrations by Measure Nonstandard Cell Potentials 0.0592 (32.6) n log Q where n represents the moles of electrons exchanged according to the cell reaction For the copper–silver cell, n ϭ 2; two electrons are lost per copper atom and two electrons are gained per two silver ions (see equations 32.1–32.3) For dilute ionic concentrations, the reaction quotient, Q, equals the mass action expression for the cell reaction For the copper–silver cell (see equation 32.1): Nernst equation: Ecell ϭ EЊcell Ϫ Qϭ [Cu2ϩ] [Agϩ]2 In Part C of this experiment, we study in depth the effect that changes in concentration of an ion have on the potential of the cell The cell potentials for a number of zinc–copper redox couples are measured in which the copper ion concentrations are varied but the zinc ion concentration is maintained constant Mass action expression: the product of the molar concentrations of the products divided by the product of the molar concentrations of the reactants, each concentration raised to the power of its coefficient in the balanced cell equation Zn(s) ϩ Cu2ϩ(aq) l Cu(s) ϩ Zn2ϩ(aq) The Nernst equation for this reaction is Ecell ϭ EЊcell Ϫ [Zn2ϩ] 0.0592 log [Cu2ϩ] (32.7) Rearrangement of this equation (where EЊcell and [Zn2ϩ] are constants in the experiment) yields an equation for a straight line: Experiment 32 353 To simplify, pCu ϭ Ϫlog [Cu2ϩ] 0.0592 pCu (32.9) A plot of Ecell versus pCu for solutions of known copper ion concentrations has a negative slope of 0.0592/2 and an intercept b that includes not only the constants in equation 32.8 but also the inherent characteristics of the cell and potentiometer (Figure 32.2) The Ecell of a solution with an unknown copper ion concentration is then measured; from the linear plot, its concentration is determined Ecell ϭ constant Ϫ Experimental Procedure Procedure Overview: The cell potentials for a number of galvanic cells are measured and the redox couples are placed in order of decreasing reduction potentials The effects of changes in ion concentrations on cell potentials are observed and analyzed Perform the experiment with a partner At each circled superscript 1–12 in the procedure, stop and record your observation on the Report Sheet Discuss your observations with your lab partner and your instructor A Reduction Potentials of Several Redox Couples The apparatus for the voltaic cell described in the Experimental Procedure may be different in your laboratory Consult with your instructor Collect the electrodes, solutions, and equipment Obtain four small (ϳ50 mL) beakers and ll them three-fourths full of the 0.1 M solutions as shown in Figure 32.3 Share these solutions with other chemists/groups of chemists in the laboratory Polish strips of copper, zinc, magnesium, and iron metal with steel wool or sandpaper, rinse brie y with dilute ( ϳ1 M) HNO3 (Caution!), and rinse with deionized water These polished metals, used as electrodes, should be bent to extend over the lip of their respective beakers Check out a multimeter (Figure 32.4) (or a voltmeter) with two electrical wires (preferably a red and black wire) attached to alligator clips Set up the copper–zinc cell Place a Cu strip (electrode) in the CuSO4 solution and a Zn strip (electrode) in the Zn(NO3)2 solution Roll and atten a piece of lter paper; wet the lter paper with a 0.1 M KNO3 solution Fold and insert the ends of the lter paper into the solutions in the two beakers; this is the salt bridge shown in Figures 32.1 and 32.3 Set the multimeter to the 2000-mV range or as appropriate Connect one electrode to the negative terminal of the multimeter and the other to the positive terminal.1 Chemists often use the “red, right, plus” rule in connecting the red wire to the right-side positive electrode (cathode) of the galvanic cell E Determine the copper–zinc cell potential If the multimeter reads a negative potential, reverse the connections to the electrodes Read and record the (positive) cell potential Identify the metal strips that serve as the cathode (positive terminal) Figure 32.2 The variation of Ecell versus the pCu 354 Figure 32.3 Setup for measuring the cell potentials of six galvanic cells You have now combined two half-cells to form a galvanic cell Galvanic Cells, the Nernst Equation and the anode Write an equation for the half-reaction occurring at each electrode Combine the two half-reactions to write the equation for the cell reaction Repeat for the remaining cells Determine the cell potentials for all possible galvanic cells that can be constructed from the four redox couples Refer to the Report Sheet for the various galvanic cells Prepare a new salt bridge for each galvanic cell Determine the relative reduction potentials Assuming the reduction potential of the Zn2+(0.1 M)/Zn redox couple is Ϫ0.79 V, determine the reduction potentials of all other redox couples.2 Determine the reduction potential of the unknown redox couple Place a 0.1 M solution and electrode obtained from your instructor in a small beaker Determine the reduction potential, relative to the Zn2ϩ(0.1 M)/Zn redox couple, for your unknown redox couple Effect of different molar concentrations Set up the galvanic cell shown in Figure 32.5, using M CuSO4 and 0.001 M CuSO4 solutions Immerse a polished copper electrode in each solution Prepare a salt bridge (Part A.2) to connect the two half-cells Measure the cell potential Determine the anode and the cathode Write an equation for the reaction occurring at each electrode Effect of complex formation Add 2–5 mL of M NH3 to the 0.001 M CuSO4 solution until any precipitate redissolves.3 (Caution: Do not inhale NH3.) Observe and record any changes in the half-cell and the cell potential Effect of precipitate formation Add 2–5 mL of 0.2 M Na2S to the 0.001 M CuSO4 solution now containing the added NH3 What is observed in the half-cell and what happens to the cell potential? Record your observations B Effect of Concentration Changes on Cell Potential Prepare the diluted solutions Prepare solutions through as shown in Figure 32.6 using a 1-mL pipet and 100-mL volumetric asks Be sure to rinse the pipet with the more concentrated solution before making the transfer Use deionized water for dilution to the mark in the volumetric asks Calculate the molar concentration of the Cu2ϩ ion for each solution and record Measure and calculate the cell potential for solution Set up the experiment as shown in Figure 32.7, page 356, using small (ϳ50 mL) beakers The Zn2ϩ/Zn redox couple is the reference half-cell for this part of the experiment Connect the two half-cells with a new salt bridge Reset the multimeter to the C The Nernst Equation and an Unknown Concentration M Figure 32.4 A modern multimeter M Figure 32.5 Setup for measuring the cell potential of a Cu2ϩ concentration cell Figure 32.6 Successive quantitative dilution, starting with 0.1 M CuSO4 Note: These are not standard reduction potentials because M concentrations of cations at 25ЊC are not used Copper ion forms a complex with ammonia: Cu2ϩ(aq) ϩ NH3(aq) l [Cu(NH3)4]2ϩ(aq) Share these prepared solutions with other chemists/groups of chemists in the laboratory Experiment 32 355 Figure 32.7 Setup to measure the effect that diluted solutions have on cell potentials Appendix C lowest range (ϳ200 mV) Connect the electrodes to the multimeter and record the potential difference, Ecell, expt Calculate the theoretical cell potential Ecell, calc (Use a table of standard reduction potentials and the Nernst equation.) 10 Measure and calculate the cell potentials for solutions and Repeat Part C.2 with solutions and 2, respectively A freshly prepared salt bridge is required for each cell Plot the data Plot Ecell, expt and Ecell, calc (ordinate) versus pCu (abscissa) on the same piece of linear graph paper (page 362) or by using appropriate software for the four concentrations of CuSO4 (see data from Part A.3 for the potential of solution 1) Have your instructor approve your graph.11 Determine the concentration of the unknown Obtain a CuSO4 solution with an unknown copper ion concentration from your instructor and set up a like galvanic cell Determine Ecell as in Part C.2 Using the graph, determine the unknown copper(II) ion concentration in the solution.12 Disposal: Dispose of the waste zinc, copper, magnesium, and iron solutions in the Waste Metal Solutions container Return the metals to appropriately marked containers CLEANUP: Rinse the beakers twice with tap water and twice with deionized water Discard the rinses in the Waste Metal Solutions container The Next Step 356 Galvanic cells are the basis for the design of speci c ion electrodes, electrodes that sense the relative concentration of a speci c ion (e.g., hydrogen ion) relative to the electrode that has a xed concentration Part C of this experiment could be the apparatus for measuring concentrations of Cu2ϩ in other samples According to equation 32.8, the pCu (negative log of [Cu2ϩ]) is proportional to the Ecell! Research speci c ion electrodes, their design, and their application Design an experiment in which a speci c ion electrode, other than the pH electrode, can be used to systematically study an ion of interest Galvanic Cells, the Nernst Equation Experiment 32 Prelaboratory Assignment Galvanic Cells, the Nernst Equation Date Lab Sec Name Desk No In a galvanic cell, a reduction occurs at the (name of electrode) b the anode is the (sign) electrode c anions ow in solution toward the (name of electrode) d electrons ow from the (name of electrode) to (name of electrode) _ _ a What is the purpose of a salt bridge? Explain b How is the salt bridge prepared in this experiment? Experimental Procedure, Part C.1 A 1-mL pipet is used to transfer 1.0 mL of a 0.10 M CusO4 solution to a 100-mL volumetric ask The volumetric ask is then lled to the mark with deionized water What is the molar concentration of the diluted solution? Show calculations expressing the concentration with the correct number of signi cant guers Refer to Figure 32.2 and equations 32.8 and 32.9 a What is the value of the cell constant? b What is the [Cu2ϩ] if the measured cell potential is 0.96 V? c What should be the cell potential if the [Cu2ϩ] is 1.0 ϫ 10Ϫ3 mol/L? Experiment 32 357 Consider a galvanic cell consisting of the following two redox couples: Agϩ(0.010 M) ϩ eϪ l Ag(s) Ϫ Cr (0.010 M) ϩ e l Cr(s) 3ϩ EЊ ϭ ϩ0.80 V EЊ ϭ Ϫ0.74 V a Write the equation for the half-reaction occurring at the cathode Cr3+(0.010 M ) b Write the equation for the half-reaction occurring at the anode M Cr c Write the equation for the cell reaction d What is the standard cell potential, EЊcell, for the cell? e Realizing the nonstandard concentrations, what is the actual cell potential, Ecell, for the cell? See equation 32.6 Hint: What is the value of n in the Nernst equation? *6 The extent of corrosion in the steel reinforcing rods (rebar) of concrete is measured by the galvanic cell shown in the diagram of the instrument The half-cell of the probe is usually a AgCl/Ag redox couple: AgCl ϩ eϪ l Ag ϩ ClϪ (1.0 M) EЊ ϭ ϩ0.23 V Corrosion is said to be severe if the cell potential is measured at greater than 0.41 V Under these conditions, what is the iron(II) concentration on the rebar? See equation 32.6 Fe2ϩ ϩ eϪ l Fe EЊ ϭ Ϫ0.44 V Rebar 358 Galvanic Cells, the Nernst Equation Concrete Experiment 32 Report Sheet Galvanic Cells, the Nernst Equation Date Lab Sec Name Desk No A Reduction Potentials of Several Redox Couples Fill in the following table with your observations and interpretations from the galvanic cells Galvanic Ecell Equation for Equation for Cell Measured Anode Anode Reaction Cathode Cathode Reaction _ _ Cu–Zn Cu–Mg _ _ _ _ Cu–Fe _ _ Zn–Mg _ _ Fe–Mg _ _ Zn–Fe _ _ Write balanced equations for the six cell reactions What is the oxidizing agent in the Zn–Mg cell? Compare the sum of the Cu–Zn and Zn–Mg cell potentials with the Cu–Mg cell potential Explain Compare the sum of the Zn–Fe and Zn–Mg cell potentials with the Fe–Mg cell potential Explain Experiment 32 359 Complete the table as follows: • Ecell Measured: Re-enter the values from Part A, Column • Reduction potential (experimental): Enter the reduction potential for each redox couple relative to Ϫ0.79V for the Zn2ϩ (0.1 M)/Zn redox couple Use EM2ϩ/M ϭ Ecell, measured ϩ (Ϫ0.79V), assuming Zn as the anode • Reduction potential (theoretical) Enter the reduction potential for each redox couple (M2ϩ/M) as calculated from a table of standard reduction potentials and the Nernst equation (equation 32.7) for [M2+] = 0.10 M • % Error See Appendix B Galvanic Cell _ Ecell Measured For the Redox Couple Reduction Potential (experimental) _ Reduction Potential (theoretical) _ % Error _ Cu–Zn _ Cu2ϩ/Cu _ _ _ Zn–Fe _ Fe2ϩ/Fe _ _ _ Zn–Zn _ Zn2ϩ/Zn –0.79 V _ –0.79 _ _ Zn–Mg _ Mg2ϩ/Mg _ _ _ Zn–unknown, X _ X2ϩ, X _ _ _ Reduction potential of the unknown redox couple: B Effect of Concentration Changes on Cell Potential Cell potential of concentration cell: Anode reaction: _ Cathode reaction: Explain why a potential is recorded Cell potential from complex formation: Observation of solution in half-cell Explain why the potential changes as it does with the addition of NH3(aq) Cell potential from precipitate formation: Observation of solution in half-cell Explain why the potential changes as it does with the addition of Na2S 360 Galvanic Cells, the Nernst Equation C The Nernst Equation and an Unknown Concentration Complete the following table with the concentrations of the Cu(NO3)2 solutions and the measured cell potentials Use equation 32.9 to determine Ecell, calc Solution Number 11 Concentration of Cu(NO3)2 Ecell, experimental Ϫlog [Cu2ϩ], pCu 10 Ecell, calculated 0.1 mol/L Instructor’s approval of graph: Account for any signi cant difference between the measured and calculated Ecell values 12 Ecell for the solution of unknown concentration: Molar concentration of Cu2ϩ in the unknown: Laboratory Questions Circle the questions that have been assigned Part A.3 The lter paper salt bridge is not wetted with the 0.1 M KNO3 solution As a result, will the measured potential of the cell be too high, too low, or unaffected? Explain Part A.3 A positive potential is recorded when the copper electrode is the positive electrode Is the copper electrode the cathode or the anode of the cell? Explain Part A.5 The measured reduction potentials are not equal to the calculated reduction potentials Give two reasons why this might be observed Part B.2 Would the cell potential be higher or lower if the NH3(aq) had been added to the M CuSO4 solution instead of the 0.001 M CuSO4 solution of the cell? Explain Part B.3 The cell potential increased (compared to Part B.2) with the addition of the Na2S solution to the 0.001 M CuSO4 solution Explain Part C As the concentration of the copper(II) ion increased from solution to solution 1, did the measured cell potentials increase or decrease? Explain why the change occurred Part C Suppose the 0.1 M Zn2ϩ solution had been diluted (instead of the Cu2ϩ solution), Would the measured cell potentials have increased or decreased? Explain why the change occurred Part C How would you increase or decrease the Cu2ϩ concentration and/or increase or decrease the Zn2ϩ concentration to maximize the cell potential? Explain how the change for each ion would maximize the cell potential Experiment 32 361 362 Galvanic Cells, the Nernst Equation 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 [...]... increase or decrease the Zn2ϩ concentration to maximize the cell potential? Explain how the change for each ion would maximize the cell potential Experiment 32 361 362 Galvanic Cells, the Nernst Equation 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... anode (ϩ) reaction: 2 BrϪ(aq) l Br2(l) ϩ 2 eϪ 2ϩ cell reaction: Cu (aq) ϩ 2 BrϪ(aq) l Cu(s) ϩ Br2(l) 1 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 ... 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