A sample of iron ore weighing 0.2792 g was dissolved in an

Problem 131AP As discussed in Section 19.5. the potential of a concentration cell diminishes as the cell operates and the concentrations in the two compartments approach each other. When the concentrations in both compartments are the same, the cell ceases to operate. At this stage, is it possible to generate a cell potential by adjusting a parameter other than concentration? Explain.

Problem 1QP Balance the following redox equations by the half-reaction method:

Problem 1VC In the first scene of the animation, when a zinc bar is immersed in an aqueous copper sulfate solution, solid copper deposits on the bar. What reaction would take place if a copper bar were immersed in an aqueous zinc sulfate solution? (a) No reaction would take place. ________________ (b) Solid copper would still deposit on the bar. ________________ (c) Solid zinc would deposit on the bar.

Problem 2QP Balance the following redox equations by the half-reaction method:

Problem 2VC What causes the change in the potential of the galvanic cell in Figure 19.1 as the cell operates? (a) Changes in the sizes of the zinc and copper electrodes. ________________ (b) Changes in the concentrations of zinc and copper ions. ________________ (c) Changes in the volumes of solutions in the half-cells.

Problem 3QP Define the following terms: anode, cathode, cell voltage, electromotive force, standard reduction potential.

Why does the color of the blue solution in the galvanic cell (Figure 19.1) fade as the cell operates? (a) Blue \(\mathrm{Cu}^{2+}\) ions are replaced by colorless \(\mathrm{Zn}^{2+}\) ions solution. (b) Blue \(\mathrm{Cu}^{2+}\) ions are removed from solution by reduction. (c) Blue \(\mathrm{Cu}^{2+}\) ions are removed from solution by oxidation.

Problem 4QP Describe the basic features of a galvanic cell. Why are the two components of the cell separated from each other?

Problem 4VC What happens to the mass of the copper electrode in the galvanic cell in Figure 19.1 as the cell operates? (a) It increases. ________________ (b) It decreases. ________________ (c) It does not change.

Problem 5QP What is the function of a salt bridge? What kind of electrolyte should be used in a salt bridge?

Problem 7QP What is tile difference between the half-reactions discussed in redox processes in Chapter 4 and the half-cell reactions discussed in Section 19.2?

Problem 8QP Discuss the spontaneity of an electrochemical reaction in terms of its standard cmf

Problem 9QP After operating a Daniel cell (see Figure 19.1) for a few minutes, the cell cmf begins to drop. Explain.

Problem 10QP Calculate the standard cmf of a cell that uses the Mg/Mg2+ and Cu/Cu2 half-cell reactions at 25ºC. Write the equation for the cell reaction that occurs under standard-state conditions.

Problem 6QP What is a cell diagram? Write the cell diagram for a galvanic cell consisting of an A1 electrode placed in a 1 M Al(NO3)3 solution and an Ag electrode placed in a 1 M AgNO3, solution.

Problem 11QP Calculate the standard cmf of a cell that uses Ag/Ag+ and Al/Al3- half-cell reactions. Write the cell reaction that occurs under standard-state conditions.

Problem 12QP Predict whether Fe3+ can oxidize I? to I2 under standard-state conditions.

Which of the following reagents can oxidize \(\mathrm{H}_{2} \mathrm{O}\) to \(\mathrm{O}_{2}\)(g) under standard-state conditions: \(\mathrm{H}^{+}\)(aq), \(\mathrm{Cl}^{-}\)(aq), \(\mathrm{Cl}_{2}\)(g), \(\mathrm{Cu}^{2+}\)(aq), \(\mathrm{Pb}^{2+}\)(aq), \(\mathrm{MnO}_{4}^{-}\)(aq) (in acid)?

Problem 14QP Consider the following half-reactions: Predict whether ions will oxidize Mn2+ to MnO4 under standard-state conditions.

Predict whether the following reactions would occur spontaneously in aqueous solution at 25°C. Assume that the initial concentrations of dissolved species are all 1.0 M. (a) \\mathrm{Ca}(s)+\mathrm{Cd}^{2+}(a q) \longrightarrow \mathrm{Ca}^{2+}(a q)+\mathrm{Cd}(s)\) (b) \(2 \mathrm{Br}^{-}(a q)+\mathrm{Sn}^{2+}(a q) \longrightarrow \mathrm{Br}_{2}(l)+\mathrm{Sn}(s)\) (c) \(2 \mathrm{Ag}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow 2 \mathrm{Ag}^{+}(a q)+\mathrm{Ni}(s)\) (d) \(\mathrm{Cu}^{+}(a q)+\mathrm{Fe}^{3+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{Fe}^{2+}(a q)\)

Which species in each pair is a better oxidizing agent under standard-state conditions: (a) \(\mathrm{Br}_{2}\) or \(\mathrm{Au}^{3+}\), (b) \(\mathrm{H}_{2}\) or \(\mathrm{Ag}^{+}\), (c) \(\mathrm{Cd}^{2+}\) or \(\mathrm{Cr}^{3+}\), (d) \(\mathrm{O}_{2}\) in acidic media or \(\mathrm{O}_{2}\) in basic media?

Problem 17QP Which species in each pair is a better reducing agent under standard-state conditions: (a) Na or Li, (b) H2 or I2, (c) Fc2+ or Ag, (d) Br? or Co2+?

Problem 19QP Compare the ease of measuring the equilibrium constant electrochemically with that by chemical means (sec Equation 18.14).

Problem 20QP Use the information in Table 2.1. and calculate the Faraday, constant.

Problem 18QP Write the equations relating ?G° and K to the standard cmf of a cell Define all the terms.

Problem 21QP What is the equilibrium constant for the following reaction at 25°C?

Problem 22QP The equilibrium constant for the reaction is 2.69 × 1012 at 25°C. Calculate Eo for a cell made up of Sr/Sr2+ and Mg/Mg2+ half-cells.

Problem 23QP Use the standard reduction potentials to find the equilibrium constant for each of the following reactions at 25°C:

Problem 25QP Under standard-state conditions, what spontaneous reaction will occur in aqueous solution among the ions Ce4+, Ce3+, Fe3+, and Fe2+? Calculate ?Go and Kc for the reaction.

Problem 24QP Calculate ?G° and Kc for the following reactions at 25°C:

Problem 26QP Given that Eo = 0.52 V for the reduction Cu+ (aq) + e? Cu(s), calculate Eº, ?G°, and K for the following reaction at 25°C:

Problem 27QP Balance (in acidic medium) the equation for the oxidation of tin from an amalgam filling when it comes into contact with aluminum foil, and calculate the standard cell potential for the reaction.

Problem 29QP Write the Nernst equation, and explain all the terms.

Problem 28QP Calculate the standard free-energy change and the equilibrium constant at 25°C for the reaction in Problem 19.27.

Problem 30QP Write the Nernst equation for the following processes at some temperature T:

Problem 31QP What is the potential of a cell made up of Zn/Zn2+ and Cu/Cu2+ half-cells at 25°C if [Zn2+ ] = 0.25 M and [Cu2+] = 0.15 M?

Problem 33QP Calculate the standard potential of the cell consisting of the Zn/Zn2+ half-cell and the SHE. What will the cmf of the cell be if [Zn2+] = 0.45 M, = 2.0 atm, and [H+] = 1.8 M?

Calculate \(E^{\circ}\), E, and \(\Delta G\) for the following cell reactions. (a) \(\mathrm{Mg}(s)+\mathrm{Sn}^{2+}(a q) \rightleftarrows \mathrm{Mg}^{2+}(a q)+\mathrm{Sn}(s) \left[\mathrm{Mg}^{2+}\right]=0.045 \mathrm{M},\left[\mathrm{Sn}^{2+}\right]=0.035 \mathrm{M}\) (b) \(3 \mathrm{Zn}(s)+2 \mathrm{Cr}^{3+}(a q) \rightleftarrows 3 \mathrm{Zn}^{2+}(a q)+2 \mathrm{Cr}(s) \left[\mathrm{Cr}^{3+}\right]=0.010 M,\left[\mathrm{Zn}^{2+}\right]=0.0085 M\)

Problem 35QP Referring to the arrangement in Figure 19.1, calculate the [Cu2]/[Zn2+] ratio at which the following reaction is spontaneous at 25°C:

Problem 34QP What is the cmf of a cell consisting of a Pb2+/Pb half-cell and a Pt/H /H2 half-cell if [Pb2+] = 0.10 M. [H] = 0.050 M, and = 1.0 atm?

Problem 36QP Calculate the cmf of the following concentration cell:

Problem 38QP Explain the differences between a primary galvanic cell—one that is not rechargeable—and a storage cell (for example, the lead storage battery), which is rechargeable.

Problem 39QP Discuss the advantages and disadvantages of fuel cells over conventional power plants in producing electricity.

Problem 40QP The hydrogen-oxygen fuel cell is described in Section 19.6. (a) What volume of H2(g), stored at 25°C at a pressure of 155 atm, would be needed to run an electric motor drawing a current of 8.5 A for 3.0 h? (b) What volume (in liters) of air at 25°C and 1.00 atm will have to pass into the cell per minute to run the motor? Assume that air is 20 percent O2 by volume and that all the O2 is consumed in the cell. The other components of air do not affect the fuel-cell reactions. Assume ideal gas behavior.

Problem 41QP Calculate the standard emf of the propane fuel cell discussed on page 830 at 25°C, given that ? for propane is -23.5 kJ/mol.

Problem 37QP What is a battery? Describe several types of batteries.

Problem 42QP What is the difference between a galvanic cell (such as a Daniell cell) and an electrolytic cell?

Problem 43QP Define the term overvoltage. How does overvoltage affect electrolytic processes?

Problem 44QP Calculate the number of grams of copper metal that can be produced by supplying 1.00 F to a solution of Cu2+ ions.

Problem 45QP The half-reaction at an electrode is Calculate the number of grams of magnesium that can be produced by supplying 1.00 F to the electrode.

Problem 46QP Consider the electrolysis of molten barium chloride (BaCl2). (a) Write the half-reactions. (b) How many grams of barium metal can be produced by supplying 0.50 A for 30 min?

Problem 47QP Considering only the cost of electricity, would it be cheaper to produce a ton of sodium or a ton of aluminum by electrolysis?

Problem 48QP If the cost of electricity to produce magnesium by the electrolysis of molten magnesium chloride is $155 per ton of metal, what is the cost (in dollars) of the electricity neccssary to produce (a) 10.0 tons of aluminum, (b) 30.0 tons of sodium, and (c) 50.0 tons of calcium?

One of the half-reactions for the electrolysis of water is \(2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)+4 e^{-}\) If 0.076 L of \(\mathrm{O}_{2}\) is collected at \(25^{\circ} \mathrm{C}\) and 755 mmHg, how many faradays of electricity had to pass through the solution?

Problem 50QP How many faradays of electricity are required to produce (a) 0.84 L of O2 at exactly 1 atm and 25°C from aqueous H2SO4 solution, (b) 1.50 L of Cl2 at 750 mmHg and 20°C from molten NaCl, and (c) 6.0 g of Sn from molten SnCl2?

Problem 51QP Calculate the amounts of Cu and Br2 produced in 1.0 h at inert electrodes in a solution of CuBr2 by a current of 4.50 A.

Problem 52QP In the electrolysis of an aqueous AgNO3 solution. 0.67 g of Ag is deposited after a certain period of time, (a) Write the half-reaction for the reduction of Ag+. (b) What is the probable oxidation half-reaction? (c) Calculate the quantity of electricity used (in coulombs).

Problem 53QP A steady current was passed through molten CoSO4 until 2.35 g of metallic cobalt was produced. Calculate the number of coulombs of electricity used.

Problem 54QP A constant electric current flows for 3 75 h through two electrolytic cells connected in series. One contains a solution of AgNO3 and the second a solution of CuCl2. During this time, 2.00 g of silver is deposited in the first cell. (a) How many grams of copper are deposited in the second cell? (b) What is the current flowing (in amperes)?

Problem 55QP What is the hourly production rate of chlorine gas (in kg) from an electrolytic cell using aqueous NaCl electrolyte and carrying a current of 1.500 x 103 A? The anode efficiency for the oxidation of Cl? is 93.0 percent.

Problem 56QP Chromium plating is applied by electrolysis to objects suspended in a dichromate solution, according to the following (unbalanced) half-reaction: How long (in hours) would it take to apply a chromium plating 1.0 × 10?2 mm thick to a car bumper with a surface area of 0.25 nr in an electrolytic cell carrying a current of 25.0 A? (The density of chromium is 7.19 g/cm3.)

Problem 57QP The passage of a current of 0.750 A for 25.0 min deposited 0.369 g of copper from a CuSO4 solution. From this information, calculate the molar mass of copper.

Problem 58QP A quantity of 0.300 g of copper was deposited from a CuSO4 solution by passing a current of 3.00 A through the solution for 304 s. Calculate the value of the Faraday constant.

In a certain electrolysis experiment, 1.44 g of Ag were deposited in one cell (containing an aqueous \(\mathrm{AgNO}_{3}\) solution), while 0.120 g of an unknown metal X was deposited in another cell (containing an aqueous \(\mathrm{XCl}_{3}\) solution) in series with the \(\mathrm{AgNO}_{3}\) cell. Calculate the molar mass of X.

Problem 60QP One of the half-reactions for the electrolysis of water is If 0.845 L of H2 is collected at 25°C and 782 mmHg. how many faradays of electricity had to pass through the solution?

Problem 61QP Steel hardware, including nuts and bolts, is often coated with a thin plating of cadmium. Explain the function of the cadmium layer.

Problem 62QP “Galvanized iron” is steel sheet that has been coated with zinc; "tin" cans are made of steel sheet coated with tin. Discuss the functions of these coatings and the electrochemistry of the corrosion reactions that occur if an electrolyte contacts the scratched surface of a galvanized iron sheet or a tin can.

Problem 64QP How does the tendency of iron to rust depend on the pH of the solution?

Problem 63QP Tarnished silver contains Ag2S. The tarnish can be removed by placing silverware in an aluminum pan containing an inert electrolyte solution, such as NaCl. Explain the electrochemical principle for this procedure. (The standard reduction potential for the half-cell reaction is ?0.71 V.]

Problem 66AP The oxidation of 25.0 mL of a solution containing Fc2+ requires 26.0 mL of 0.0250 M K2Cr2O7 in acidic solution. Balance the following equation, and calculate the molar concentration of Fc2+:

Problem 65AP For each of the following redox reactions, (i) write the half-reactions. (ii) write a balanced equation for the whole reaction, (in) determine in which direction the reaction will proceed spontaneously under standard-state conditions:

Problem 67AP The SO2 present in air is mainly responsible for the phenomenon of acid rain. The concentration of SO2 can be determined by titrating against a standard permanganate solution as follows: Calculate the number of grams of SO2 in a sample of air if 7.37 mL of 0.00800 M KMnO4 solution is required for the titration.

A sample of iron ore weighing \(0.2792 \mathrm{~g}\) was dissolved in an excess of a dilute acid solution. All the iron was first converted to \(\mathrm{Fe}\) (II) ions. The solution then required \(23.30 \mathrm{~mL}\) of \(0.0194 M\) \(\mathrm{KMnO}_4\) for oxidation to \(\mathrm{Fe}\) (III) ions. Calculate the percent by mass of iron in the ore.

Problem 69AP The concentration of a hydrogen peroxide solution can be conveniently determined by titration against a standardized potassium permanganate solution in an acidic medium according to the following unbalanced equation: (a) Balance this equation, (b) If 36.44 mL of a 0.01652 M KMnO4 solution is required to completely oxidize 25.00 mL of an H2O2 solution, calculate the molarity of the H2O2 solution.

Problem 70AP Oxalic acid (H2C2O4) is present in many plants and vegetables. (a) Balance the following equation in acid solution: ________________ (b) If a 1.00-g sample of plant matter requires 24.0 mL of 0.0100 M KMnO4 solution to reach the equivalence point, what is the percent by mass of H2C4 in the sample?

Problem 72AP Complete the following table. State whether the cell reaction is spontaneous, nonspontancous, or at equilibrium. E ?G Cell Reaction > 0 >0 = 0

Problem 73AP From the following information, calculate the solubility product of AgBr:

Problem 71AP Calcium oxalate (CaC2O4) is insoluble in water. This properly has been used to determine the amount of Ca2+ ions in blood. The calcium oxalate isolated from blood is dissolved in acid and titrated against a standardized KMnO4 solution as described in Problem 19.68. In one test it is found that the calcium oxalate isolated from a 10.0-mL sample of blood requires 24.2 mL of 9.56 × 10?4 M KMnO4 for titration. Calculate the number of milligrams of calcium per milliliter of blood.

Problem 74AP Consider a galvanic cell composed of the SHE and a half-cell using the reaction Ag (aq) + e? Ag(s). (a) Calculate the standard cell potential, (b) What is the spontaneous cell reaction under standard-state conditions? (c) Calculate the cell potential when [H+] in the hydrogen electrode is changed to (i) 1.0 x 10?2 M and (ii) 1.0 × 10?5 M, all other reagents being held at standard-state conditions, (d) Based on this cell arrangement, suggest a design for a pH meter.

Problem 75AP A galvanic cell consists of a silver electrode in contact with 346 mL of 0.100 M AgNO3 solution and a magnesium electrode in contact with 288 mL of 0.100 M Mg(NO3)2 solution. (a) Calculate E for the cell at 25°C. (b) A current is drawn from the cell until 1.20 g of silver has been deposited at the silver electrode. Calculate E for the cell at this stage of operation.

Explain why chlorine gas can be prepared by electrolyzing an aqueous solution of NaCl but fluorine gas cannot be prepared by electrolyzing an aqueous solution of NaF

Problem 77AP Calculate the cmf of the following concentration cell at 25°C:

Problem 78AP The cathode reaction in the Leelanché cell is given by If a Leclanché cell produces a current of 0.0050 A, calculate how many hours this current supply will last if there is initially 4.0 g of MnO2 present in the cell. Assume that there is an excess of Zn2+ ions.

Problem 79AP For a number of years, it was not clear whether mercury(I) ions existed in solution as Hg+ or as . To distinguish between these two possibilities, we could set up the following system: where soln A conlained 0.263 g mcrcury(I) nitrate per liter and soln B contained 2.63 g mercury(I) nitrate per liter. If the measured cmf of such a cell is 0.0289 V at 18°C, what can you deduce about the nature of the mercury(I) ions?

Problem 80AP An aqueous KI solution to which a few drops of phenolphthalein have been added is electrolyzed using an apparatus like the one shown here: Describe w hat you would observe at the anode and the cathode. (Hint: Molecular iodine is only slightly soluble in water, but in the presence of I? ions, it forms the brown color of ions. See Problem 13.114.)

Problem 81AP A piece of magnesium metal weighing 1.56 g is placed in 100.0 mL of 0.100 M AgNO3 at 25°C. Calculate [Mg2+ ] and [Ag+] in solution at equilibrium. What is the mass of the magnesium left? The volume remains constant.

Problem 82AP Describe an experiment that would enable you to determine which is the cathode and which is the anode in a galvanic cell using copper and zinc electrodes.

Problem 83AP An acidified solution was electrolyzed using copper electrodes. A constant current of 1.18 A caused the anode to lose 0.584 g after 1.52 × 103 s. (a) What is the gas produced at the cathode, and what is its volume at STP? (b) Given that the charge of an electron is 1.6022 x 10?19 C, calculate Avogadro's number. Assume that copper is oxidized to Cu2+ ions.

Problem 84AP In a certain electrolysis experiment involving Al3+ ions, 60.2 g of A1 is recovered when a current of 0.352 A is used. How many minutes did the electrolysis last?

Problem 85AP Consider the oxidation of ammonia: (a) Calculate the ?Gº for the reaction, (b) If this reaction were used in a fuel cell, what would the standard cell potential be?

Problem 86AP When an aqueous solution containing gold(III) salt is electrolyzed, metallic gold is deposited at the cathode and oxygen gas is generated at the anode. (a) If 9.26 g of Au is deposited at the cathode, calculate the volume (in liters) of O2 generated at 23°C and 747 mmHg. (b) What is the current used if the electrolytic process took 2.00 h?

Problem 87AP In an electrolysis experiment, a student passes the same quantity of electricity through two electrolytic cells, one containing a silver salt and the other a gold salt. Over a certain period of time, the student finds that 2.64 g of Ag and 1.61 g of Au are deposited at the cathodes. What is the oxidation state of gold in the gold salt?

Problem 88AP Consider the electrochemical cell represented by the following diagram: Determine the initial value of Ecell under the conditions shown in the cell diagram; and determine the initial value of Ecell if [Ag+] were increased by a factor of 4 at 25°C.

Problem 89AP Given that calculate ?Gº and K for the following process at 25°C: (The preceding reaction is an example of a disproportionation reaction in which an element in one oxidation state is both oxidized and reduced.)

Problem 90AP Fluorine (F2) is obtained by the electrolysis of liquid hydrogen fluoride (HF) containing potassium fluoride (KF). (a) Write the half-cell reactions and the overall reaction for the process. (b) What is the purpose of KF? (c) Calculate the volume of F2 (in liters) collected at 24.0°C and 1.2 aim after electrolyzing the solution for 15 h at a current of 502 A.

Problem 93AP An aqueous solution of a platinum salt is electrolyzed at a current of 2.50 A for 2.00 h. As a result. 9.09 g of metallic Pt is formed at the cathode. Calculate the charge on the Pt ions in this solution.

Problem 91AP A 300-mL solution of NaCl was electrolyzed for 6.00 min. If the pH of the final solution was 12.24, calculate the average current used.

Industrially, copper is purified by electrolysis. The impure copper acts as the anode, and the cathode is made of pure copper. The electrodes are immersed in a \(\mathrm{CuSO}_4\) solution. During electrolysis, copper at the anode enters the solution as \(\mathrm{Cu}^{2+}\) while \(\mathrm{Cu}^{2+}\) ions are reduced at the cathode. (a) Write half-cell reactions and the overall reaction for the electrolytic process. (b) Suppose the anode was contaminated with \(\mathrm{Zn}\) and \(\mathrm{Ag}\). Explain what happens to these impurities during electrolysis. (c) How many hours will it take to obtain \(1.00 \mathrm{~kg}\) of \(\mathrm{Cu}\) at a current of \(18.9 \mathrm{~A}\)?

Problem 94AP Consider a galvanic cell consisting of a magnesium electrode in contact with 1.0 M Mg(NO3)2 and a cadmium electrode in contact with 1.0 M Cd(NO3)2. Calculate E° for the cell, and draw a diagram showing the cathode, anode, and direction of electron flow.

Problem 95AP A current of 6.00 A passes through an electrolytic cell containing dilute sulfuric acid for 3.40 h. If the volume of O2 gas generated at the anode is 4.26 L (at STP), calculate the charge (in coulombs) on an electron.

Problem 96AP Gold will not dissolve in either concentrated nitric acid or concentrated hydrochloric acid. However, the metal does dissolve in a mixture of the acids (one part HNO3 and three parts HCl by volume), called aqua regia. (a) Write a balanced equation for this reaction. (Hint: Among the products are HAuCl4 and NO2) (b) What is the function of HCI?

Problem 97AP Explain why most useful galvanic cells give voltages of no more than 1.5 to 2.5 V. What are the prospects for developing practical galvanic cells with voltages of 5 V or more?

Problem 98AP A silver rod and a SHE are dipped into a saturated aqueous solution of silver oxalate (Ag2C2O4), at 25°C. The measured potential difference between the rod and the SHE is 0.589 V, the rod being positive. Calculate the solubility product constant for silver oxalate.

Problem 99AP Zinc is an amphoteric metal; that is, it reacts with both acids and bases. The standard reduction potential is ?1.36 V for the reaction Calculate the formation constant (Kf) for the reaction

Use the data in Table 19.1 to determine whether or not hydrogen peroxide will undergo disproportionation in an acid medium: \(2 \mathrm{H}_{2} \mathrm{O}_{2} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\)

The magnitudes (but not the signs) of the standard reduction potentials of two metals X and Y are \(\mathrm{Y}^{2+}+2 e^{-} \longrightarrow \mathrm{Y} \quad\left|E^{\circ}\right|=0.34 \mathrm{~V}\) \(\mathrm{X}^{2+}+2 e^{-} \longrightarrow \mathrm{X} \quad\left|E^{\circ}\right|=0.25 \mathrm{~V}\) where the || notation denotes that only the magnitude (but not the sign) of the E° value is shown. When the half-cells of X and Y are connected, electrons flow from X to Y. When X is connected to a SHE, electrons flow from X to SHE. (a) Are the E° values of the half- reactions positive or negative? (b) What is the standard emf of a cell made up of X and Y?

Problem 102AP A galvanic cell is constructed as follows. One half-cell consists of a platinum wire immersed in a solution containing 1.0 M Sn2+ and 1.0 M Sn4+ ; the other half-cell has a thallium rod immersed in a solution of 1.0 M Tl+. (a) Write the half-cell reactions and the overall reaction. (b) What is the equilibrium constant at 25°C? (c) What is the cell voltage if the Tl+ concentration is increased 10-fold?

Problem 104AP The ingestion of a very small quantity of mercury is not considered loo harmful. Would this statement still hold if the gastric juice in your stomach were mostly nitric acid instead of hydrochloric acid?

Given the standard reduction potential for \(\mathrm{Au}^{3+}\) in Table 19.1 and \(\mathrm{Au}^{+}(a q)+e^{-} \longrightarrow \mathrm{Au}(s) \quad E^{\circ}=1.69 \mathrm{V}\) answer the following questions. (a) Why does gold not tarnish in air? (b) Will the following disproportionation occur spontaneously? \(3 \mathrm{Au}^{+}(a q) \longrightarrow \mathrm{Au}^{3+}(a q)+2 \mathrm{Au}(s)\) (c) Predict the reaction between gold and fluorine gas.

Problem 105AP When 25.0 mL of a solution containing both Fe2+ and Fe3+ ions is titrated with 23.0 mL of 0.0200 M MnO4 (in dilute sulfuric acid), all the Fe2+ ions are oxidized to Fe3+ ions. Next, the solution is treated with Zn metal to convert all the Fe3+ ions to Fe2+ ions. Finally, 40.0 mL of the same KMnO4 solution is added to the solution to oxidize the Fe2+ ions to Fe3+. Calculate the molar concentrations of Fe2+ and Fe3+ in the original solution.

Problem 106AP Consider the Daniell cell in Figure 19.1. When viewed externally, the anode appears negative and the cathode positive (electrons are flowing from the anode to the cathode). Yet in solution anions are moving toward the anode, which means that it must appear positive to the anions. Because the anode cannot simultaneously be negative and positive, give an explanation for this apparently contradictory situation.

Problem 107AP Use the data in Table 19.1 to show that the decomposition of H2O2: (a disproportionation reaction) is spontaneous at 25°C:

Problem 108AP The concentration of sulfuric acid in the lead-storage battery of an automobile over a period of time has decreased from 38.0 percent by mass (density = 1.29 g/mL) to 26.0 percent by mass (1.19 g/mL). Assume the volume of the acid remains constant at 724 mL. (a) Calculate the total charge in coulombs supplied by the battery, (b) How long (in hours) will it take to recharge the battery back to the original sulfuric acid concentration using a current of 22.4 A?

Problem 109AP Consider a Daniell cell operating under non-standard-state conditions. Suppose that the cell's reaction is multiplied by 2. What effect does this have on each of the following quantities in the Nernst equation: (a) E, (b) E°, (c) Q, (d) In Q, (c) n?

A spoon was silver-plated electrolytically in an \(\mathrm{AgNO}_{3}\) solution. (a) Sketch a diagram for the process. (b) If 0.884 g of Ag was deposited on the spoon at a constant current of 18.5 mA, how long (in min) did the electrolysis take?

Problem 111AP Comment on whether F2 will become a stronger oxidizing agent with increasing H+ concentration.

Problem 112AP In recent years there has been much interest in electric cars. List some advantages and disadvantages of electric cars compared to automobiles with internal combustion engines.

Problem 113AP Calculate the pressure of H2 (in aim) required to maintain equilibrium with respect to the following reaction at 25°C, given that [Pb2+] = 0.035 M and the solution is buffered at pH 1.60.

Problem 114AP A piece of magnesium ribbon and a copper wire are partially immersed in a 0.1 M HC1 solution in a beaker. The metals are joined externally by another piece of metal wire. Bubbles are seen to evolve at both the Mg and Cu surfaces, (a) Write equations representing the reactions occurring at the metals. What visual evidence would you seek to show that Cu is not oxidized to Cu2+? (c) At some stage, NaOH solution is added to the beaker to neutralize the HCl acid. Upon further addition of NaOH, a white precipitate forms. What is it?

The zinc-air battery shows much promise for electric cars because it is lightweight and rechargeable: The net transformation is \(\mathrm{Zn}(s)+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \longrightarrow \mathrm{ZnO}(s)\). (a) Write the half-reactions at the zinc-air electrodes, and calculate the standard emf of the battery at \(25^{\circ} \mathrm{C}\). (b) Calculate the emf under actual operating conditions when the partial pressure of oxygen is 0.21 atm. (c) What is the energy density (measured as the energy in kilojoules that can be obtained from \(1 \mathrm{~kg}\) of the metal) of the zinc electrode? (d) If a current of \(2.1 \times 10^5 \mathrm{~A}\) is to be drawn from a zinc-air battery system, what volume of air (in liters) would need to be supplied to the battery every second? Assume that the temperature is \(25^{\circ} \mathrm{C}\) and the partial pressure of oxygen is 0.21 atm.

Problem 116AP Calculate E° for the reactions of mercury with (a) 1 M HCl and (b) 1 M HNO3. Which acid will oxidize Hg to under standard-state conditions? Can you identify which pictured test tube contains HNO3 and Hg and which contains HCl and Hg?

Problem 117AP Because all alkali metals react with water, it is not possible to measure the standard reduction potentials of these metals directly as in the case of, say, zinc. An indirect method is to consider the following hypothetical reaction: Using the appropriate equation presented in this chapter and the thermodynamic data in Appendix 2, calculate E° for Li (aq) + e? Li(s) at 298 K. Use 96,485.338 C/mol e? for the Faraday constant. Compare your result with that listed in Table 19.1.

Problem 118AP A galvanic cell using Mg/Mg2? and Cu/Cu2+ half-cells operates under standard-state conditions at 25°C, and each compartment has a volume of 218 mL. The cell delivers 0.22 A for 31.6 h. (a) How many grams of Cu are deposited? (b) What is the [Cu2+ ] remaining?

Problem 120AP Compare the pros and cons of a fuel cell, such as the hydrogen-oxygen fuel cell, and a coal-fired power station for generating electricity.

Problem 119AP Given the following standard reduction potentials, calculate the ion-product, Kw, for water at 25°C:

Problem 121AP Lead storage batteries are rated by ampere-hours, that is, the number of amperes they can deliver in an hour. (a) Show that 1 Ah = 3600 C. (b) The lead anodes of a certain lead-storage battery have a total mass of 406 g. Calculate the maximum theoretical capacity of the battery in ampere-hours. Explain why in practice we can never extract this much energy from the battery. (Hint: Assume all the lead will be used up in the electrochemical reaction, and refer to the electrode reactions on page 839.) (c) Calculate and ?G° for the battery.

Problem 125AP Based on the following standard reduction potentials, calculate the standard reduction potential for the half-reaction

Problem 122AP Use Equations 18.l0and 19.3 to calculate the cmf values of the Daniell cell at 25°C and 80°C. Comment on your results. What assumptions are used in the derivation? (Hint: You need the thermodynamic data in Appendix 2.)

A construction company is installing an iron culvert (a long cylindrical tube) that is \(40.0 \mathrm{~m}\) long with a radius of \(0.900 \mathrm{~m}\). To prevent corrosion, the culvert must be galvanized. This process is carried out by first passing an iron sheet of appropriate dimensions through an electrolytic cell containing \(\mathrm{Zn}^{2+}\) ions, using graphite as the anode and the iron sheet as the cathode. If the voltage is \(3.26 \mathrm{~V}\), what is the cost of electricity for depositing a layer \(0.200 \mathrm{~mm}\) thick if the efficiency of the process is 95 percent? The electricity rate is $0.12 per kilowatt hour \((\mathrm{kWh})\), where \(1 \mathrm{~W}=1 \mathrm{~J} / \mathrm{s}\) and the density of \(\mathrm{Zn}\) is \(7.14 \mathrm{~g} / \mathrm{cm}^3\).

Problem 124AP A 9.00 × 102mL amount of 0.200 M Mgl2: solution was electrolyzed. As a result, hydrogen gas was generated at the cathode and iodine was formed at the anode. The volume of hydrogen collected at 26°C and 779 mmHg was 1.22 x 103 mL. (a) Calculate the charge in coulombs consumed in the process. (b) How long (in min) did the electrolysis last if a current of 7.55 A was used? (c) A white precipitate was formed in the process. What was it, and what was its mass in grams? Assume the volume of the solution was constant.

Calculate the equilibrium constant for the following reaction at \(298 \mathrm{~K}\) : \(\mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \rightleftarrows \mathrm{Zn}^{2+}(a q)+\mathrm{Cu}(s)\)

Problem 126AP Which of the components of dental amalgam (mercury, silver, tin, copper, or zinc) would be oxidized when a filling is brought into contact with lead?

Problem 128AP Cytochrome-c is a protein involved in biological electron transfer processes. The redox half-reaction is shown by the reduction of the Fe2+ ion to the Fe2+ ion: Calculate the number of moles of cyt c(Fe3+) formed from cyt c(Fe3+) with the Gibbs free energy derived from the oxidation of 1 mole of glucose.

Problem 129AP The nitrite ion in soil is oxidized to the nitrate ion by the bacterium Nitrobacter agilis in the presence of oxygen. The half-reactions are Calculate the yield of ATP synthesis per mole of nitrite oxidized. (Hint: Refer to Section 18.7.)

Problem 130AP Fluorine is a highly reactive gas that attacks water to form HF and other products. Follow the procedure in Problem 19.115 to show how you can determine indirectly the standard reduction potential for fluorine Compare your result with the value in Table 19.1.