Balance each redox reaction occurring in basic aqueous

Is a spontaneous redox reaction obtained by pairing any reduction half-reaction with one listed above it or with one listed below it in Table 18.1 ?

How can Table 18.1 be used to predict whether or not a metal will dissolve in HCl? In HNO3?

Explain why , , and K are all interrelated.

Does a redox reaction with a small equilibrium constant (K < 1) have a positive or a negative ? Does it have a positive or a negative ?

Problem 26E Explain how a fuel-cell breathalyzer works.

Problem 27E List some applications of electrolysis.

Problem 28E The anode of an electrolytic cell must be connected to which terminal—positive or negative—of the power source?

Problem 29E What species is oxidized and what species is reduced in the electrolysis of a pure molten salt?

Problem 30E If an electrolytic cell contains a mixture of species that can be oxidized, how do you determine which species will actually be oxidized? If it contains a mixture of species that can be reduced, how do you determine which one will actually be reduced?

Problem 31E Why does the electrolysis of an aqueous sodium chloride solution produce hydrogen gas at the cathode?

Problem 33E How is the amount of current flowing through an electrolytic cell related to the amount of product produced in the redox reaction?

Problem 34E What is corrosion? Why is corrosion only a problem for some metals (such as iron)?

Problem 35E Explain the role of each of the following in promoting corrosion: moisture, electrolytes, and acids.

Problem 36E How can the corrosion of iron be prevented?

Balance each redox reaction occurring in acidic aqueous solution.

Balance each redox reaction occurring in acidic aqueous solution.

Balance each redox reaction occurring in basic aqueous solution.

Balance each redox reaction occurring in basic aqueous solution.

Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow.

Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow.

Problem 25E What is a fuel cell? What is the most common type of fuel cell and what reactions occur at its anode and cathode?

Problem 23E What are the anode and cathode reactions in a lead–acid storage battery? What happens when the battery is recharged?

What are the anode and cathode reactions in a common dry-cell battery? In an alkaline battery?

Problem 21E What is a concentration electrochemical cell?

Use the Nernst equation to show that = under standard conditions.

How does depend on the concentrations of the reactants and products in the redox reaction occurring in the cell? What effect does increasing the concentration of a reactant have on ? Increasing the concentration of a product?

How is the cell potential of an electrochemical cell () related to the potentials of the half-cells?

Problem 12E Describe the standard hydrogen electrode (SHE) and explain its use in determining standard electrode potentials.

Problem 10E Describe the basic features of a cell diagram (or line notation) for an electrochemical cell.

What is the definition of the standard cell potential ()? What does a large positive standard cell potential imply about the spontaneity of the redox reaction occurring in the cell? What does a negative standard cell potential imply about the reaction?

Problem 8E What unit is used to measure the magnitude of electrical current? What unit is used to measure the magnitude of a potential difference? Explain how electrical current and potential difference differ.

Problem 6E What reaction (oxidation or reduction) occurs at the cathode of a voltaic cell? What is the sign of the cathode? Do electrons flow toward or away from the cathode?

What reaction (oxidation or reduction) occurs at the anode of a voltaic cell? What is the sign of the anode? Do electrons flow toward or away from the anode?

Problem 4E Explain the difference between a voltaic (or galvanic) electrochemical cell and an electrolytic one.

Problem 3E Give the basic definitions of oxidation and reduction and explain the basic procedure for balancing redox reactions.

Problem 2E In electrochemistry, what kind of reaction can be driven by electricity?

Problem 1E In electrochemistry, spontaneous redox reactions are used for what purpose?

Consider the following half-reactions: MnO2 4 (aq) 1 8H1(aq) 1 5e2 Mn21(aq) 1 4H2O(l) NO2 3 (aq) 1 4H1(aq) 1 3e2 NO(g) 1 2H2O(l) Predict whether NO3 2 ions will oxidize Mn21 to MnO4 2 under standard-state conditions

Predict whether the following reactions would occur spontaneously in aqueous solution at 25C. Assume that the initial concentrations of dissolved species are all 1.0 M. (a) Ca(s) 1 Cd21(aq) Ca21(aq) 1 Cd(s) (b) 2Br2(aq) 1 Sn21(aq) Br2(l) 1 Sn(s) (c) 2Ag(s) 1 Ni21(aq) 2Ag1(aq) 1 Ni(s) (d) Cu1(aq) 1 Fe31(aq) Cu21(aq) 1 Fe21(aq)

Which species in each pair is a better oxidizing agent under standard-state conditions? (a) Br2 or Au31, (b) H2 or Ag1, (c) Cd21 or Cr31, (d) O2 in acidic media or O2 in basic media

Which species in each pair is a better reducing agent under standard-state conditions? (a) Na or Li, (b) H2 or I2, (c) Fe21 or Ag, (d) Br2 or Co21

Consider the electrochemical reaction Sn21 1 X S Sn 1 X21. Given that E cell 5 0.14 V, what is the E for the X21/X half-reaction?

The E cell for the following cell is 1.54 V at 25C U(s) 0 U31(aq) 0 0 Ni21(aq) 0 Ni(s) Calculate the standard reduction potential for the U31/U half-cell

Write the equations relating DG and K to the standard emf of a cell. Define all the terms

The E value of one cell reaction is positive and that of another cell reaction is negative. Which cell reaction will proceed toward the formation of more products at equilibrium?

What is the equilibrium constant for the following reaction at 25C? Mg(s) 1 Zn21(aq) Mg21(aq) 1 Zn(s)

The equilibrium constant for the reaction Sr(s) 1 Mg21(aq) Sr21(aq) 1 Mg(s) is 2.69 3 1012 at 25C. Calculate E for a cell made up of Sr/Sr21 and Mg/Mg21 half-cells

Use the standard reduction potentials to find the equilibrium constant for each of the following reactions at 25C: (a) Br2(l) 1 2I2(aq) 2Br2(aq) 1 I2(s) (b) 2Ce41(aq) 1 2Cl2(aq) Cl2(g2 1 2Ce31(aq) (c) 5Fe21(aq) 1 MnO4 2(aq) 1 8H1(aq) Mn21(aq) 1 4H2O(l) 1 5Fe31(aq)

Calculate DG and Kc for the following reactions at 25C: (a) Mg(s) 1 Pb21(aq) Mg21(aq) 1 Pb(s) (b) Br2(l) 1 2I2(aq) 2Br2(aq) 1 I2(s) (c) O2(g2 1 4H1(aq) 1 4Fe21(aq) 2H2O(l) 1 4Fe31(aq) (d) 2Al(s) 1 3I2(s) 2Al31(aq) 1 6I2(aq)

Under standard-state conditions, what spontaneous reaction will occur in aqueous solution among the ions Ce41, Ce31, Fe31, and Fe21? Calculate DG and Kc for the reaction

Given that E 5 0.52 V for the reduction Cu1(aq) 1 e2 S Cu(s), calculate E, DG, and K for the following reaction at 25C: 2Cu1(aq) Cu21(aq) 1 Cu(s)

Write the Nernst equation and explain all the terms.

Write the Nernst equation for the following processes at some temperature T: (a) Mg(s) 1 Sn21(aq) Mg21(aq) 1 Sn(s) (b) 2Cr(s) 1 3Pb21(aq) 2Cr31(aq) 1 3Pb(s)

What is the potential of a cell made up of Zn/Zn21 and Cu/Cu21 half-cells at 25C if [Zn21] 5 0.25 M and [Cu21] 5 0.15 M?

Calculate E, E, and DG for the following cell reactions. (a) Mg(s) 1 Sn21(aq) Mg21(aq) 1 Sn(s) [Mg21] 5 0.045 M, [Sn21] 5 0.035 M (b) 3Zn(s) 1 2Cr31(aq) 3Zn21(aq) 1 2Cr(s) [Cr31] 5 0.010 M, [Zn21] 5 0.0085 M

Calculate the standard potential of the cell consisting of the Zn/Zn21 half-cell and the SHE. What will the emf of the cell be if [Zn21] 5 0.45 M, PH2 5 2.0 atm, and [H1] 5 1.8 M?

What is the emf of a cell consisting of a Pb21/Pb half-cell and a Pt/H1/H2 half-cell if [Pb21] 5 0.10 M, [H1] 5 0.050 M, and PH2 5 1.0 atm?

Referring to the arrangement in Figure 18.1, calculate the [Cu21]/[Zn21] ratio at which the following reaction is spontaneous at 25C: Cu(s) 1 Zn21(aq) Cu21(aq) 1 Zn(s)

Calculate the emf of the following concentration cell: Mg(s) 0 Mg21(0.24 M) 0 0 Mg21(0.53 M) 0 Mg(s)

Explain the differences between a primary galvanic cellone that is not rechargeableand a storage cell (for example, the lead storage battery), which is rechargeable.

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

The hydrogen-oxygen fuel cell is described in Section 18.6. (a) What volume of H2(g), stored at 25C 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 (liters) of air at 25C 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.

Calculate the standard emf of the propane fuel cell discussed on p. 836 at 25C, given that DGf for propane is 223.5 kJ/mol.

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

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

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 Ag2S(s) 1 2e 2 S 2Ag(s) 1 S22(aq) is 20.71 V.]

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

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

Describe the electrolysis of an aqueous solution of KNO3

The half-reaction at an electrode is Mg21(molten) 1 2e2 Mg(s) Calculate the number of grams of magnesium that can be produced by supplying 1.00 F to the electrode

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?

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

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 necessary to produce (a) 10.0 tons of aluminum, (b) 30.0 tons of sodium, (c) 50.0 tons of calcium?

One of the half-reactions for the electrolysis of water is 2H2O (l) O2(g) 1 4H1(aq) 1 4e2 If 0.076 L of O2 is collected at 25C and 755 mmHg,how many moles of electrons had to pass through thesolution?

How many moles of electrons are required to produce (a) 0.84 L of O2 at exactly 1 atm and 25C from aqueous H2SO4 solution; (b) 1.50 L of Cl2 at 750 mmHg and 20C from molten NaCl; (c) 6.0 g of Sn from molten SnCl2?

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

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 Ag1. (b) What is the probable oxidation halfreaction? (c) Calculate the quantity of electricity used, in coulombs

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.

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 are 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?

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 3 103 A? The anode efficiency for the oxidation of Cl2 is 93.0 percent.

Chromium plating is applied by electrolysis to objects suspended in a dichromate solution, according to the following (unbalanced) halfreaction: Cr2O22 7 (aq) 1 e2 1 H1(aq) Cr(s) 1 H2O(l) How long (in hours) would it take to apply a chromium plating 1.0 3 1022 mm thick to a car bumper with a surface area of 0.25 m2 in an electrolytic cell carrying a current of 25.0 A? (The density of chromium is 7.19 g/cm3 .)

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.

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 AgNO3 solution), while 0.120 g of an unknown metal X was deposited in another cell (containing anaqueous XCl3 solution) in series with the AgNO3 cell. Calculate the molar mass of X.

One of the half-reactions for the electrolysis of water is 2H1(aq) 1 2e2 H2(g) If 0.845 L of H2 is collected at 25C and 782 mmHg, how many moles of electrons had to pass through the solution?

A steady current of 10.0 A is passed through three electrolytic cells for 10.0 min. Calculate the mass of the metals formed if the solutions are 0.10 M AgNO3, 0.10 M Cu(NO3)2, and 0.10 M Au(NO3)3.

Industrially, copper metal can be purified electrolytically according to the following arrangement. The anode is made of the impure Cu electrode and the cathode is the pure Cu electrode. The electrodes are immersed in a CuSO4 solution. (a) Write the half-cell reactions at the electrodes. (b) Calculate the mass (in g) of Cu purified after passing a current of 20 A for 10 h. (c) Explain why impurities such as Zn, Fe, Au, and Ag are not deposited at the electrodes. Pure copper cathode Cu2+ SO4 2 Impure copper anode Battery

A Daniell cell consists of a zinc electrode in 1.00 L of 1.00 M ZnSO4 and a Cu electrode in 1.00 L of 1.00 M CuSO4 at 25C. A steady current of 10.0 A is drawn from the cell. Calculate the Ecell after 1.00 h. Assume volumes to remain constant

A concentration cell is constructed having Cu electrodes in two CuSO4 solutions A and B. At 25C, the osmotic pressures of the two solutions are 48.9 atm and 4.89 atm, respectively. Calculate the Ecell, assuming no ion-pair formation

For each of the following redox reactions, (i) write the half-reactions, (ii) write a balanced equation for the whole reaction, (iii) determine in which direction the reaction will proceed spontaneously under standard-state conditions: (a) H2(g) 1 Ni21(aq) H1(aq) 1 Ni(s)(b) MnO42(aq) 1 Cl2(aq) Mn21(aq) 1 Cl2(g) (in acid solution)(c) Cr(s) 1 Zn21(aq) Cr31(aq) 1 Zn(s)

The oxidation of 25.0 mL of a solution containing Fe21 requires 26.0 mL of 0.0250 M K2Cr2O7 in acidic solution. Balance the following equation and calculate the molar concentration of Fe21: Cr2O7 22 1 Fe21 1 H1 Cr31 1 Fe31

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: 5SO2 1 2MnO4 2 1 2H2O 5SO22 4 1 2Mn21 1 4H1 Calculate the number of grams of SO2 in a sample of air if 7.37 mL of 0.00800 M KMnO4 solution are required for the titration.

A sample of iron ore weighing 0.2792 g was dissolved in an excess of a dilute acid solution. All the iron was first converted to Fe(II) ions. The solution then required 23.30 mL of 0.0194 M KMnO4 for oxidation to Fe(III) ions. Calculate the percent by mass of iron in the ore

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: MnO4 2 1 H2O2 O2 1 Mn21 (a) Balance the above equation. (b) If 36.44 mL of a 0.01652 M KMnO4 solution are required to completely oxidize 25.00 mL of a H2O2 solution, calculate the molarity of the H2O2 solution.

Oxalic acid (H2C2O4) is present in many plants and vegetables. (a) Balance the following equation in acid solution: MnO2 4 1 C2O22 4 Mn21 1 CO2 (b) If a 1.00-g sample of H2C2O4 requires 24.0 mL of 0.0100 M KMnO4 solution to reach the equivalence point, what is the percent by mass of H2C2O4 in the sample?

Complete the following table. State whether the cell reaction is spontaneous, nonspontaneous, or at equilibrium. E DG Cell Reaction . 0 . 0 5 0

Calcium oxalate (CaC2O4) is insoluble in water. This property has been used to determine the amount of Ca21 ions in blood. The calcium oxalate isolated from blood is dissolved in acid and titrated against a standardized KMnO4 solution as described in Problem 18.72. 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 3 1024 M KMnO4 for titration. Calculate the number of milligrams of calcium per milliliter of blood.

From the following information, calculate the solubility product of AgBr: Ag1(aq) 1 e2 Ag(s) E 5 0.80 V AgBr(s) 1 e2 Ag(s) 1 Br2(aq) E 5 0.07 V

Consider a galvanic cell composed of the SHE and a half-cell using the reaction Ag1(aq) 1 e2 S 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 [H1] in the hydrogen electrode is changed to (i) 1.0 3 1022 M and (ii) 1.0 3 1025 M, all other reagents being held at standard-state conditions. (d) Based on this cell arrangement, suggest a design for a pH meter.

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 25C. (b) A current is drawn from the cell until 1.20 g of silver have 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.

Calculate the emf of the following concentration cell at 25C: Cu(s) 0 Cu21(0.080 M) 0 0 Cu21(1.2 M) 0 Cu(s)

The cathode reaction in the Leclanch cell is given by 2MnO2(s) 1 Zn21(aq) 1 2e2 ZnMn2O4(s) If a Leclanch cell produces a current of 0.0050 A, calculate how many hours this current supply will last if there are initially 4.0 g of MnO2 present in the cell. Assume that there is an excess of Zn21 ions.

Suppose you are asked to verify experimentally the electrode reactions shown in Example 18.8. In addition to the apparatus and the solution, you are also given two pieces of litmus paper, one blue and the other red. Describe what steps you would take in this experiment.

For a number of years it was not clear whether mercury(I) ions existed in solution as Hg1 or as Hg2 21. To distinguish between these two possibilities, we could set up the following system: Hg(l) 0 soln A 0 0 soln B0 Hg(l) where soln A contained 0.263 g mercury(I) nitrate per liter and soln B contained 2.63 g mercury(I) nitrate per liter. If the measured emf of such a cell is 0.0289 V at 18C, what can you deduce about the nature of the mercury(I) ions?

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 what you would observe at the anode and the cathode. (Hint: Molecular iodine is only slightly soluble in water, but in the presence of I2 ions, it forms the brown color of I3 2 ions. See Problem 12.102.)

A piece of magnesium metal weighing 1.56 g is placed in 100.0 mL of 0.100 M AgNO3 at 25C. Calculate [Mg21] and [Ag1] in solution at equilibrium. What is the mass of the magnesium left? The volume remains constant

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.

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 3 103 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 3 10219 C, calculate Avogadros number. Assume that copper is oxidized to Cu21 ions.

In a certain electrolysis experiment involving Al31 ions, 60.2 g of Al is recovered when a current of 0.352 A is used. How many minutes did the electrolysis last?

Consider the oxidation of ammonia: 4NH3(g) 1 3O2(g) 2N2(g) 1 6H2O(l) (a) Calculate the DG for the reaction. (b) If this reaction were used in a fuel cell, what would the standard cell potential be?

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 23C and 747 mmHg. (b) What is the current used if the electrolytic process took 2.00 h?

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, she 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?

People living in cold-climate countries where there is plenty of snow are advised not to heat their garages in the winter. What is the electrochemical basis for this recommendation?

Given that 2Hg21(aq) 1 2e2 Hg2 21(aq) E 5 0.92 V Hg2 21(aq) 1 2e2 2Hg(l) E 5 0.85 V calculate DG and K for the following process at 25C: Hg2 21(aq) Hg21(aq) 1 Hg(l) (The preceding reaction is an example of a disproportionation reaction in which an element in one oxidation state is both oxidized and reduced.)

A galvanic cell with E cell 5 0.30 V can be constructed using an Fe electrode in a 1.0 M Fe(NO3)2 solution, and either a Sn electrode in a 1.0 M Sn(NO3)2 solution, or a Cr electrode in a 1.0 M Cr(NO3)3 solution, even though Sn21/Sn and Cr31/Cr have different standard reduction potentials. Explain

Shown here is a galvanic cell connected to an electrolytic cell. Label the electrodes (anodes and cathodes) and show the movement of electrons along the wires and cations and anions in solution. For simplicity, the salt bridge is not shown for the galvanic cell. Galvanic cell Electrolytic cell

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 24.0C and 1.2 atm after electrolyzing the solution for 15 h at a current of 502 A

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 CuSO4 solution. During electrolysis, copper at the anode enters the solution as Cu21 while Cu21 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 Zn and Ag. Explain what happens to these impurities during electrolysis. (c) How many hours will it take to obtain 1.00 kg of Cu at a current of 18.9 A?

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 are formed at the cathode. Calculate the charge on the Pt ions in this solution

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.

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.

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 HCl?

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?

The table here shows the standard reduction potentials of several half-reactions: Half-Reactions E8 (V) A21 1 2e2 A 21.46 B2 1 2e2 2B2 0.33 C31 1 3e2 C 1.13 D1 1 e2 D 20.87 (a) Which is the strongest oxidizing agent andwhich is the strongest reducing agent? (b) Which substances can be oxidized by B2? (c) Which substances can be reduced by B2? (d) Write the overall equation for a cell that delivers a voltage of 2.59 Vunder standard-state conditions

Consider a concentration cell made of the following two compartments: Cl2(0.20 atm) 0 Cl2(1.0 M) and Cl2(2.0 atm) 0 Cl2(1.0 M). Platinum is used as the inert electrodes. Draw a cell diagram for the cell and calculate the emf of the cell at 25C.

A silver rod and a SHE are dipped into a saturated aqueous solution of silver oxalate, Ag2C2O4, at 25C. 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

Zinc is an amphoteric metal; that is, it reacts with both acids and bases. The standard reduction potential is 21.36 V for the reaction Zn(OH)22 4 (aq) 1 2e2 Zn(s) 1 4OH2(aq) Calculate the formation constant (Kf) for the reaction Zn21(aq) 1 4OH2(aq) Zn(OH24 22(aq)

Use the data in Table 18.1 to determine whether or not hydrogen peroxide will undergo disproportionation in an acid medium: 2H2O2 S 2H2O 1 O2.

The magnitudes (but not the signs) of the standard reduction potentials of two metals X and Y are Y21 1 2e2 Y 0 E0 5 0.34 V X21 1 2e2 X 0 E0 5 0.25 V where the 0 0 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?

A galvanic cell is constructed as follows. One halfcell consists of a platinum wire immersed in a solution containing 1.0 M Sn21 and 1.0 M Sn41; the other half-cell has a thallium rod immersed in a solution of 1.0 M Tl1. (a) Write the half-cell reactions and the overall reaction. (b) What is the equilibrium constant at 25C? (c) What is the cell voltage if the T11 concentration is increased tenfold? (E Tl1/Tl 5 20.34 V.)

Given the standard reduction potential for Au31 in Table 18.1 and Au1(aq) 1 e2 Au(s) E 5 1.69 V answer the following questions. (a) Why does gold not tarnish in air? (b) Will the following disproportionationoccur spontaneously? 3Au1(aq) Au31(aq) 1 2Au(s) (c) Predict the reaction between gold and fluorine gas

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

When 25.0 mL of a solution containing both Fe21 and Fe31 ions is titrated with 23.0 mL of 0.0200 M KMnO4 (in dilute sulfuric acid), all of the Fe21 ions are oxidized to Fe31 ions. Next, the solution is treated with Zn metal to convert all of the Fe31 ions to Fe21 ions. Finally, 40.0 mL of the same KMnO4 solution are added to the solution in order to oxidize the Fe21 ions to Fe31. Calculate the molar concentrations of Fe21 and Fe31 in the original solution.

Consider the Daniell cell in Figure 18.1. 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

Use the data in Table 18.1 to show that the decomposition of H2O2 (a disproportionation reaction) is spontaneous at 25C: 2H2O2(aq) 2H2O(l) 1 O2(g)

Consider two electrolytic cells A and B. Cell A contains a 0.20 M CoSO4 solution and platinum electrodes. Cell B differs from cell A only in that cobalt metals are used as electrodes. In each case, a current of 0.20 A is passed through the cell for 1.0 h. (a) Write equations for the half-cell and overall cell reactions for these cells. (b) Calculate the products formed (in grams) at the anode and cathode in each case.

A galvanic cell consists of a Mg electrode in a 1 M Mg(NO3)2 solution and another metal electrode X in a 1 M X(NO3)2 solution. Listed here are the E cell values of four such galvanic cells. In each case, identify X from Table 18.1. (a) E cell 5 2.12 V, (b) E cell 5 2.24 V, (c) E cell 5 1.61 V, (d) E cell 5 1.93 V

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 5 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 amperes?

Consider a Daniell cell operating under nonstandardstate conditions. Suppose that the cells 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) ln Q, and (e) n?

An electrolysis cell was constructed similar to the one shown in Figure 18.18, except 0.1 M MgCl2(aq) was used as the electrolyte solution. Under these conditions, a clear gas was formed at one electrode and a very pale green gas was formed at the other electrode in roughly equal volumes. (a) What gases are formed at these electrodes? (b) Write balanced half-reactions for each electrode. Account for any deviation from the normally expected results.

Comment on whether F2 will become a stronger oxidizing agent with increasing H1 concentration.

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.

Calculate the pressure of H2 (in atm) required to maintain equilibrium with respect to the following reaction at 25C: Pb(s) 1 2H1(aq) Pb21(aq) 1 H2(g) Given that [Pb21] 5 0.035 M and the solution is buffered at pH 1.60

A piece of magnesium ribbon and a copper wire are partially immersed in a 0.1 M HCl 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. (b) What visual evidence would you seek to show that Cu is not oxidized to Cu21? (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 Zn(s)11 2O2(g) SZnO(s). (a) Write the half-reactions at the zinc-air electrodes and calculate the standard emf of the battery at 25C. (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 kg of the metal) of the zinc electrode? (d) If a current of 2.1 3 105 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 25C and the partial pressure of oxygen is 0.21 atm.

Calculate E for the reactions of mercury with (a) 1 M HCl and (b) 1 M HNO3. Which acid will oxidize Hg to Hg2 21 under standard-state conditions? Can you identify which test tube shown contains HNO3 and Hg and which contains HCl and Hg?

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 Li1(aq) 1 1 2H2(g) Li(s) 1 H1(aq) Using the appropriate equation presented in this chapter and the thermodynamic data in Appendix 3, calculate E for Li1(aq) 1 e2 S Li(s) at 298 K. Compare your result with that listed in Table 18.1. (See back endpaper for the Faraday constant.)

A galvanic cell using Mg/Mg21 and Cu/Cu21 halfcells operates under standard-state conditions at 25C 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 [Cu21] remaining?

Given the following standard reduction potentials, calculate the ion-product, Kw, for water at 25C: 2H1(aq) 1 2e2 H2(g) E 5 0.00 V 2H2O(l) 1 2e2 H2(g) 1 2OH2(aq) E 5 20.83 V 18.129 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

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.

Lead storage batteries are rated by ampere hours, that is, the number of amperes they can deliver in an hour. (a) Show that 1 A ? h 5 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 of the lead will be used up in the electrochemical reaction and refer to the electrode reactions on p. 833.) (c) Calculate E cell and DG for the battery.

Use Equations (17.10) and (18.3) to calculate the emf values of the Daniell cell at 25C and 80C. Comment on your results. What assumptions are used in the derivation? (Hint: You need the thermodynamic data in Appendix 3.)

A construction company is installing an iron culvert (a long cylindrical tube) that is 40.0 m long with a radius of 0.900 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 Zn21 ions, using graphite as the anode and the iron sheet as the cathode. If the voltage is 3.26 V, what is the cost of electricity for depositing a layer 0.200 mm thick if the efficiency of the process is 95 percent? The electricity rate is $0.12 per kilowatt hour (kWh), where 1 W 5 1 J/s and the density of Zn is 7.14 g/cm3 .

A 9.00 3 102 -mL 0.200 M MgI2 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 26C and 779 mmHg was 1.22 3 103 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.

Based on the following standard reduction potentials: Fe21(aq) 1 2e2 Fe(s) E 1 5 20.44 V Fe31(aq) 1 e2 Fe21(aq) E 2 5 0.77 V calculate the standard reduction potential for the half-reaction Fe31(aq) 1 3e2 Fe(s) E 3 5 ?

A galvanic cell is constructed by immersing a piece of copper wire in 25.0 mL of a 0.20 M CuSO4 solution and a zinc strip in 25.0 mL of a 0.20 M ZnSO4 solution. (a) Calculate the emf of the cell at 25C and predict what would happen if a small amount of concentrated NH3 solution were added to (i) the CuSO4 solution and (ii) the ZnSO4 solution. Assume that the volume in each compartment remains constant at 25.0 mL. (b) In a separate experiment, 25.0 mL of 3.00 M NH3 are added to the CuSO4 solution. If the emf of the cell is 0.68 V, calculate the formation constant (Kf) of Cu(NH3)4 21

Calculate the equilibrium constant for the following reaction at 298 K: Zn(s) 1 Cu21(aq) Zn21(aq) 1 Cu(s)

To remove the tarnish (Ag2S) on a silver spoon, a student carried out the following steps. First, she placed the spoon in a large pan filled with water so the spoon was totally immersed. Next, she added a few tablespoonful of baking soda (sodium bicarbonate), which readily dissolved. Finally, she placed some aluminum foil at the bottom of the pan in contact with the spoon and then heated the solution to about 80C. After a few minutes, the spoon was removed and rinsed with cold water. The tarnish was gone and the spoon regained its original shiny appearance. (a) Describe with equations the electrochemical basis for the procedure. (b) Adding NaCl instead of NaHCO3 would also work because both compounds are strong electrolytes. What is the added advantage of using NaHCO3? (Hint: Consider the pH of the solution.) (c) What is the purpose of heating the solution? (d) Some commercial tarnish removers contain a fluid (or paste) that is a dilute HCl solution. Rubbing the spoon with the fluid will also remove the tarnish. Name two disadvantages of using this procedure compared to the one described above

The nitrite ion (NO2 2) in soil is oxidized to nitrate ion (NO3 2) by the bacteria Nitrobacter agilis in the presence of oxygen. The half-reduction reactions are NO2 3 1 2H1 1 2e2 NO2 2 1 H2O E 5 0.42 V O2 1 4H1 1 4e2 2H2O E 5 1.23 V Calculate the yield of ATP synthesis per mole of nitrite oxidized. (Hint: See Section 17.7.)

The diagram here shows an electrolytic cell consisting of a Co electrode in a 2.0 M Co(NO3)2 solution and a Mg electrode in a 2.0 M Mg(NO3)2 solution. (a) Label the anode and cathode and show the half-cell reactions. Also label the signs (1 or 2) on the battery terminals. (b) What is the minimum voltage to drive the reaction? (c) After the passage of 10.0 A for 2.00 h the battery is replaced with a voltmeter and the electrolytic cell now becomes a galvanic cell. Calculate Ecell. Assume volumes to remain constant at 1.00 L in each compartment. Battery Salt bridge

Fluorine is a highly reactive gas that attacks water to form HF and other products. Follow the procedure in Problem 18.126 to show how you can determine indirectly the standard reduction for fluorine as shown in Table 18.1

Show a sketch of a galvanic concentration cell. Each compartment consists of a Co electrode in a Co(NO3)2 solution. The concentrations in the compartments are 2.0 M and 0.10 M, respectively. Label the anode and cathode compartments. Show the direction of electron flow. (a) Calculate the Ecell at 25C. (b) What are the concentrations in the compartments when the Ecell drops to 0.020 V? Assume volumes to remain constant at 1.00 L in each compartment.

The emf of galvanic cells varies with temperature (either increases or decreases). Starting with Equation (18.3), derive an equation that expresses E cell in terms of DH and DS. Predict whether E cell will increase or decrease if the temperature of a Daniell cell increases. Assume both DH and DS to be temperature independent.

A concentration cell ceases to operate when the concentrations of the two cell compartments are equal. At this stage, is it possible to generate an emf from the cell by adjusting another parameter without changing the concentrations? Explain.

It has been suggested that a car can be powered from the hydrogen generated by reacting aluminum soda cans with a solution of lye (sodium hydroxide) according to the following reaction: 2Al(s) 1 2OH2(aq) 1 6H2O(l) 2 Al(OH)2 4 (aq) 1 3H2(g) How many aluminum soda cans would be required to generate the same amount of chemical energy as contained in one tank of gasoline? Read the Chemistry in Action on aluminum recycling in Chapter 21 (p. 950), and comment on the cost and environmental impact of powering a car with aluminum cans

Estimate how long it would take to electroplate a teaspoon with silver from a solution of AgNO3, assuming a constant current of 2 A

The potential for a cell based on the standard hydrogen electrode and the half-reaction Mn1(aq) 1 ne2 M(s) was measured at several concentrations of Mn1(aq), giving the following plot. What is the value of n in the half-reaction? 0.635 0.8 Ecell (V) 0.60.4 log (1/[Mn1]) 0.20 0.640 0.645 0.650 0.655

Balance the redox reaction equation (occurring in acidic solution) and choose the correct coefficients for each reactant and product.

Problem 2SAQ Which statement is true for voltaic cells? a) Electrons flow from the anode to the cathode. b) Electrons flow from the more negatively charged electrode to the more positively charged electrode. c) Electrons flow from higher potential energy to lower potential energy. d) All of the above

Use data from Table 18.1 to calculate for the reaction. a) 1.77 V b) 2.03 V c) 0.82 V d) 1.08 V

Use data from Table 18.1 to determine which statement is true of the voltaic cell picture here. a) Sn is the anode; Ag is the cathode; electrons flow from left to right b) Sn is the cathode; Ag is the anode; electrons flow from left to right c) Sn is the anode; Ag is the cathode; electrons flow from right to left d) Sn is the cathode; Ag is the anode; electrons flow from right to left

Use data from Table 18.1 to determine which metal does not dissolve in hydrochloric acid (HCl). a) Zn b) Cd c) Cu d) Fe

The Zn/Zn2+ electrode has a standard electrode potential of EO = - 0.76 V. How does the relative potential energy of an electron at the Zn/Zn2+ electrode compare to the potential energy of an electron at the standard hydrogen electrode? a) An electron at the Zn/Zn2+ electrode has a higher potential energy than an electron at the standard hydrogen electrode. b) An electron at the Zn/Zn2+ electrode has a lower potential energy than an electron at the standard hydrogen electrode. c) An electron at the Zn/Zn2+ electrode has the same potential energy as an electron at the standard hydrogen electrode. d) Nothing can be concluded about the relative potential energy of an electron at the standard electrode potential.

Explain the purpose of a salt bridge in an electrochemical cell.

Use data from Table 18.1 to calculate \(\Delta G^{\circ}\) for the reaction. \(2\ \mathrm{MnO}_4^-(aq)+\mathrm{Cd}(s)\ \rightarrow\ 2\mathrm{\ MnO}_4^{2-}(aq)+\mathrm{Cd}^{2+}(aq)\) a. +30.9 kJ b. -30.9 kJ c. -185 kJ d. +185 kJ

A redox reaction has an = - 0.56 V. What can you conclude about the equilibrium constant () for the reaction? a) K < 1 b) K > 1 c) K = 0 d) Nothing can be concluded about from .

Find Ecell for an electrochemical cell based on the following reaction with [MnO4-] = 2.0 M, [H+] = 1.0 M, and [Ag+] = 0.010 M. for the reaction is + 0.88 V. a) 0.83 V b) 1.00 V c) 0.76 V d) 0.93 V

In an electrochemical cell, = 0.010 and = 855. What can you conclude about and ? a) is positive and is negative. b) is negative and is positive. c) and are both negative. d) and are both positive.

Problem 11E Why do some electrochemical cells employ inert electrodes such as platinum?

Which reaction occurs at the anode of a lead storage battery?

Which reaction could be used to generate electricity in a voltaic electrochemical cell?

Which reaction occurs at the cathode of an electrolytic cell containing a mixture of molten KCl and ZnCl2?

Problem 14E Does a large positive electrode potential indicate a strong oxidizing agent or a strong reducing agent? What about a large negative electrode potential?

Copper is plated onto the cathode of an electrolytic cell containing CuCl2(aq). How long does it take to plate 111 mg of copper with a current of 3.8 A? a) 1.3 x 103 s b) 44 s c) 89 s d) 22 s

Problem 15SAQ Which metal can be used as a sacrificial electrode to prevent the rusting of an iron pipe? a) Au b) Ag c) Cu d) Mn

Problem 24E What are the three common types of portable rechargeable batteries and how does each one work?

Problem 32E What is overvoltage in an electrochemical cell? Why is it important?

Balance each redox reaction occurring in acidic aqueous solution.

Balance each redox reaction occurring in acidic aqueous solution.

Calculate the standard cell potential for each of the electrochemical cells in Problem 43.

Calculate the standard cell potential for each of the electrochemical cells in Problem 44.

Consider the voltaic cell: a. Determine the direction of electron flow and label the anode and the cathode. b. Write a balanced equation for the overall reaction and calculate . c. Label each electrode as negative or positive. d. Indicate the direction of anion and cation flow in the salt bridge.

Consider the voltaic cell: a. Determine the direction of electron flow and label the anode and the cathode. b. Write a balanced equation for the overall reaction and calculate . c. Label each electrode as negative or positive. d. Indicate the direction of anion and cation flow in the salt bridge.

Use line notation to represent each electrochemical cell in Problem 43.

Use line notation to represent each electrochemical cell in Problem 44.

Make a sketch of the voltaic cell represented by the line notation. Write the overall balanced equation for the reaction and calculate .

Make a sketch of the voltaic cell represented by the line notation. Write the overall balanced equation for the reaction and calculate .

Determine whether or not each redox reaction occurs spontaneously in the forward direction.

Determine whether or not each redox reaction occurs spontaneously in the forward direction.

Which metal could you use to reduce Mn2+ ions but not Mg2+ ions?

Which metal can be oxidized with an Sn2+ solution but not with an Fe2+ solution?

Problem 57E Determine whether or not each metal dissolves in 1 M HCl. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Al b. Ag c. Pb

Problem 58E Determine whether or not each metal dissolves in 1 M HCl. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Cu b. Fe c. Au

Determine whether or not each metal dissolves in 1 M HNO3. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Cu b. Au

Determine whether or not each metal dissolves in 1 M HIO3. For those metals that do dissolve, write a balanced redox equation for the reaction that occurs. a. Au b. Cr

Calculate for each balanced redox reaction and determine if the reaction is spontaneous as written.

Calculate for each balanced redox reaction and determine if the reaction is spontaneous as written.

Which metal cation is the best oxidizing agent? a. \(\mathrm{Pb}^{2+}\) b. \(\mathrm{Cr}^{3+}\) c. \(\mathrm{Fe}^{2+}\) d. \(\mathrm{Sn}^{2+}\)

Problem 64E Which metal is the best reducing agent? a. Mn b. Al c. Ni d. Cr

Use tabulated electrode potentials to calculate \(\Delta G_{\mathrm{rxn}}^{\circ}\) for each reaction at \(25\ ^{\circ}\mathrm{C}\). a. \(\mathrm{Pb}^{2+}(a q)+\mathrm{Mg}(s) \longrightarrow \mathrm{Pb}(s)+\mathrm{Mg}^{2+}(a q)\) b. \(\mathrm{Br}_2(l)+2\mathrm{\ Cl}^-(aq)\longrightarrow2\mathrm{\ Br}^-(aq)+\mathrm{Cl}_2(g)\) c. \(\mathrm{MnO}_2(s)+4\mathrm{\ H}^+(aq)+\mathrm{Cu}(s)\longrightarrow\mathrm{Mn}^{2+}(aq)+2\mathrm{\ H}_2\mathrm{O}(l)+\mathrm{Cu}^{2+}(aq)\)

Use tabulated electrode potentials to calculate for each reaction at 25 oC.

Calculate the equilibrium constant for each of the reactions in Problem 65.

Calculate the equilibrium constant for each of the reactions in Problem 66.

Calculate the equilibrium constant for the reaction between Ni2+(aq) and Cd(s) (at 25 oC).

Calculate the equilibrium constant for the reaction between Fe2+(aq) and Zn(s) (at 25 oC).

Calculate and for a redox reaction with n = 2 that has an equilibrium constant of K = 25 (at 25 oC).

Calculate and for a redox reaction with n = 3 that has an equilibrium constant of K = 0.050 (at 25 oC).

A voltaic cell employs the following redox reaction: Calculate the cell potential at 25 oC under each set of conditions. a. standard conditions b. [Sn2+] = 0.0100 M; [Mn2+] = 2.00 M c. [Sn2+] = 2.00 M; [Mn2+] = 0.0100 M

A voltaic cell employs the redox reaction: Calculate the cell potential at 25 oC under each set of conditions. a. standard conditions b. [Fe3+] = 1.0 x 10-3 M; [Mg2+] = 2.50 M c. [Fe3+] = 2.00 M; [Mg2+] = 1.5 x 10-3 M

An electrochemical cell is based on these two half-reactions: Calculate the cell potential at 25 oC.

An electrochemical cell is based on these two half-reactions: Calculate the cell potential at 25 oC.

A voltaic cell consists of a Zn/Zn2+ half-cell and a Ni/Ni2+ half-cell at 25 oC. The initial concentrations of Ni2+ and Zn2+ are 1.50 M and 0.100 M, respectively. a. What is the initial cell potential? b. What is the cell potential when the concentration of Ni2+ has fallen to 0.500 M? c. What are the concentrations of Ni2+ and Zn2+ when the cell potential falls to 0.45 V?

A voltaic cell consists of a Pb/Pb2+ half-cell and a Cu/Cu2+ half-cell at 25 oC. The initial concentrations of Pb2+ and Cu2+ are 0.0500 M and 1.50 M, respectively. a. What is the initial cell potential? b. What is the cell potential when the concentration of Cu2+ has fallen to 0.200 M? c. What are the concentrations of Pb2+ and Cu2+ when the cell potential falls to 0.35 V?

Consider the concentration cell: a. Label the anode and cathode. b. Indicate the direction of electron flow. c. Indicate what happens to the concentration of Pb2+ in each half-cell.

Make a sketch of a concentration cell employing two Zn/Zn2+ half-cells. The concentration of Zn2+ in one of the half-cells is 2.0 M and the concentration in the other half-cell is 1.0 x 10-3 M. Label the anode and the cathode and indicate the half-reaction occurring at each electrode. Also indicate the direction of electron flow

A concentration cell consists of two Sn/Sn2+ half-cells. The cell has a potential of 0.10 V at 25 0C. What is the ratio of the Sn2+ concentrations in the two half-cells?

A Cu/Cu2+ concentration cell has a voltage of 0.22 V at 25 oC. The concentration of Cu2+ in one of the half-cells is 1.5 x 10-3 M. What is the concentration of Cu2+ in the other half-cell? (Assume the concentration in the unknown cell to be the lower of the two concentrations.)

Determine the optimum mass ratio of \(\mathrm{Zn}\) to \(\mathrm{MnO}_2\) in an alkaline battery.

Problem 84E What mass of lead sulfate is formed in a lead–acid storage battery when 1.00 g of Pb undergoes oxidation?

Refer to the tabulated values of in Appendix IIB to calculate for a fuel cell that employs the reaction between methane gas (CH4) and oxygen to form carbon dioxide and gaseous water.

Determine whether or not each metal, if coated onto iron, would prevent the corrosion of iron. (a) Zn (b) Sn (c) Mn

Refer to the tabulated values of in Appendix IIB to calculate for the fuel cell breathalyzer, which employs the following reaction. ( for HC2H3O2(g) = -374.2 kJ/mol.)

Problem 88E Determine whether or not each metal, if coated onto iron, would prevent the corrosion of iron. a. Mg b. Cr c. Cu

Consider the electrolytic cell: a. Label the anode and the cathode and indicate the half reactions occurring at each. b. Indicate the direction of electron flow. c. Label the terminals on the battery as positive or negative and calculate the minimum voltage necessary to drive the reaction.

Draw an electrolytic cell in which Mn2+ is reduced to Mn and Sn is oxidized to Sn2+. Label the anode and cathode, indicate the direction of electron flow, and write an equation for the half reaction occurring at each electrode. What minimum voltage is necessary to drive the reaction?

Problem 91E Write equations for the half-reactions that occur in the electrolysis of molten potassium bromide.

Problem 92E What products are obtained in the electrolysis of molten NaI?

Problem 93E Write equations for the half-reactions that occur in the electrolysis of a mixture of molten potassium bromide and molten lithium bromide.

Write equations for the half-reactions that occur at the anode and cathode for the electrolysis of each aqueous solution: a. NaBr(aq) b. PbI2(aq) c. Na2SO4(aq)

What products are obtained in the electrolysis of a molten mixture of KI and KBr?

Write equations for the half-reactions that occur at the anode and cathode for the electrolysis of each aqueous solution: a. Ni(NO3)2(aq) b. KCl(aq) c. CuBr2(aq)

Problem 97E Make a sketch of an electrolysis cell that electroplates copper onto other metal surfaces. Label the anode and the cathode and indicate the reactions that occur at each.

Problem 98E Make a sketch of an electrolysis cell that electroplates nickel onto other metal surfaces. Label the anode and the cathode and indicate the reactions that occur at each.

Copper can be electroplated at the cathode of an electrolysis cell by the half-reaction: How much time would it take for 325 mg of copper to be plated at a current of 5.6 A?

Silver can be electroplated at the cathode of an electrolysis cell by the half-reaction: What mass of silver would plate onto the cathode if a current of 6.8 A flowed through the cell for 72 min?

A major source of sodium metal is the electrolysis of molten sodium chloride. What magnitude of current produces 1.0 kg of sodium metal in 1 hour?

Problem 102E What mass of aluminum metal can be produced per hour in the electrolysis of a molten aluminum salt by a current of 25 A?

Consider the unbalanced redox reaction: \(\mathrm{MnO}_4{ }^{-}(a q)+\mathrm{Zn}(s) \longrightarrow \mathrm{Mn}^{2+}(a q)+\mathrm{Zn}^{2+}(a q)\) Balance the equation and determine the volume of a \(0.500 \mathrm{M} \mathrm{KMnO}_4\) solution required to completely react with \(2.85 \mathrm{~g}\) of \(\mathrm{Zn}\).

Consider the unbalanced redox reaction: Balance the equation and determine the volume of a 0.850 M K2Cr2O7 solution required to completely react with 5.25 g of Cu.

Consider the molecular views of an Al strip and \(\mathrm{Cu}^{2+}\) solution. Draw a similar sketch showing what happens to the atoms and ions after the Al strip is submerged in the solution for a few minutes.

Consider the molecular view of an electrochemical cell involving the overall reaction: Draw a similar sketch of the cell after it has generated a substantial amount of electrical current.

Problem 107E Determine whether HI can dissolve each metal sample. If it can, write a balanced chemical reaction showing how the metal dissolves in HI and determine the minimum volume of 3.5 M HI required to completely dissolve the sample. a. 2.15 gAl b. 4.85 gCu c. 2.42 gAg

Determine if HNO3 can dissolve each metal sample. If it can, write a balanced chemical reaction showing how the metal dissolves in HNO3 and determine the minimum volume of 6.0 M HNO3 required to completely dissolve the sample. a. 5.90 g Au b. 2.55 g Cu c. 4.83 g Sn

The cell potential of this electrochemical cell depends on the pH of the solution in the anode half-cell. What is the pH of the solution if is 355 mV?

The cell potential of this electrochemical cell depends on the gold concentration in the cathode half-cell. What is the concentration of Au3+ in the solution if is 1.22 V?

A friend wants you to invest in a new battery she has designed that produces 24 V in a single voltaic cell. Why should you be wary of investing in such a battery?

Problem 112E What voltage can theoretically be achieved in a battery in which lithium metal is oxidized and fluorine gas is reduced? Why might such a battery be difficult to produce?

A battery relies on the oxidation of magnesium and the reduction of Cu2+. The initial concentrations of Mg2+ and Cu2+ are 1.0 x 10-4 M and 1.5 M, respectively, in 1.0-liter half-cells. a. What is the initial voltage of the battery? b. What is the voltage of the battery after delivering 5.0 A for 8.0 h? c. How long can the battery deliver 5.0 A before going dead?

A rechargeable battery is constructed based on a concentration cell constructed of two Ag/Ag+ half-cells. The volume of each half-cell is 2.0 L and the concentrations of Ag+ in the half-cells are 1.25 M and 1.0 x 10-3 M. a. How long can this battery deliver 2.5 A of current before it goes dead? b. What mass of silver is plated onto the cathode by running at 3.5 A for 5.5 h? c. Upon recharging, how long would it take to redissolve 1.00 x 102 g of silver at a charging current of 10.0 amps?

If a water electrolysis cell operates at a current of 7.8 A, how long will it take to generate 25.0 L of hydrogen gas at a pressure of 25.0 atm and a temperature of 25 oC?

When a suspected drunk driver blows 188 mL of his breath through the fuel-cell breathalyzer described in Section 18.7 , the breathalyzer produces an average of 324 mA of current for 10 s. Assuming a pressure of 1.0 atm and a temperature of 25 o C, what percent (by volume) of the driver’s breath is ethanol?

The Ksp of Zn(OH)2 is 1.8 x 10-14. Find for the half reaction:

The Ksp of CuI is 1.1 x 10-12. Find for the cell:

Calculate \(\Delta G_{\mathrm{rxn}}^{\circ}\) and K for each reaction. a. The disproportionation of \(\mathrm{Mn}^{2+}(a q)\) to Mn(s) and \(\mathrm{MnO}_{2}(s)\) in acid solution at \(25\ ^{\circ}\mathrm{C}\). b. The disproportionation of \(\mathrm{MnO}_{2}(s)\) to \(\mathrm{Mn}^{2+}(a q)\) and \(\mathrm{MnO}_4^{\ -}(aq)\) in acid solution at \(25\ ^{\circ}\mathrm{C}\).

Calculate and K for each reaction. a. The reaction of Cr2+(aq) with Cr2O72-(aq) in acid solution to form Cr3+(aq). b. The reaction of Cr3+(aq) and Cr(s) to form Cr2+(aq). [The electrode potential of Cr2+(aq) to Cr(s) is -0.91 V. ]

The molar mass of a metal (M) is 50.9 g/mol; it forms a chloride of unknown composition. Electrolysis of a sample of the molten chloride with a current of 6.42 A for 23.6 minutes produces 1.20 g of M at the cathode. Determine the empirical formula of the chloride.

A metal forms the fluoride MF3. Electrolysis of the molten fluoride by a current of 3.86 A for 16.2 minutes deposits 1.25 g of the metal. Calculate the molar mass of the metal.

A sample of impure tin of mass 0.535 g is dissolved in strong acid to give a solution of Sn2+. The solution is then titrated with a 0.0448 M solution of NO3-, which is reduced to NO(g). The equivalence point is reached upon the addition of 0.0344 L of the NO3- solution. Find the percent by mass of tin in the original sample, assuming that it contains no other reducing agents.

A 0.0251 L sample of a solution of Cu+ requires 0.0322 L of 0.129 M KMnO4 solution to reach the equivalence point. The products of the reaction are Cu2+ and Mn2+. What is the concentration of the Cu2+ solution?

A current of 11.3 A is applied to 1.25 L of a solution of 0.552 M HBr converting some of the H+ to H2(g) , which bubbles out of solution. What is the pH of the solution after 73 minutes?

Problem 126E A 215 mL sample of a 0.500 M NaCl solution with an initial pH of 7.00 is subjected to electrolysis. After 15.0 minutes, a 10.0 mL portion (or aliquot) of the solution was removed from the cell and titrated with 0.100 M HCl solution. The endpoint in the titration was reached upon addition of 22.8 mL of HCl. Assuming constant current, what was the current (in A) running through the cell?

An MnO2(s)/Mn2+(aq) electrode in which the pH is 10.24 is prepared. Find the [Mn2+] necessary to lower the potential of the half-cell to 0.00 V (at 25 oC).

Problem 128E To what pH should you adjust a standard hydrogen electrode to get an electrode potential of -0.122 V ? (Assume that the partial pressure of hydrogen gas remains at 1 atm.)

Suppose a hydrogen–oxygen fuel-cell generator produces electricity for a house. Use the balanced redox reactions and the standard cell potential to predict the volume of hydrogen gas (at STP) required each month to generate the electricity. Assume the home uses 1.2 x 103 kWh of electricity per month.

A voltaic cell designed to measure [Cu2+] is constructed of a standard hydrogen electrode and a copper metal electrode in the Cu2+ solution of interest. If you want to construct a calibration curve for how the cell potential varies with the concentration of copper(II), what do you plot in order to obtain a straight line? What is the slope of the line?

The surface area of an object to be gold plated is 49.8 cm2 and the density of gold is 19.3 g/cm3. A current of 3.25 A is applied to a solution that contains gold in the +3 oxidation state. Calculate the time required to deposit an even layer of gold 1.00 x 10-3 cm thick on the object.

To electrodeposit all the Cu and Cd from a solution of CuSO4 and CdSO4 required 1.20 F of electricity (1 F = 1 mol e-). The mixture of Cu and Cd that was deposited had a mass of 50.36 g. What mass of CuSO4 was present in the original mixture?

Sodium oxalate, Na2C2O4, in solution is oxidized to CO2(g) by MnO4-, which is reduced to Mn2+. A 50.1 mL volume of a solution of MnO4- is required to titrate a 0.339 g sample of sodium oxalate. This solution of MnO4 - is used to analyze uranium- containing samples. A 4.62 g sample of a uranium containing material requires 32.5 mL of the solution for titration. The oxidation of the uranium can be represented by the change UO2+ UO22+. Calculate the percentage of uranium in the sample.

Problem 134E Three electrolytic cells are connected in a series. The electrolytes in the cells are aqueous copper(II) sulfate, gold(III) sulfate, and silver nitrate. A current of 2.33 A is applied, and after some time 1.74 g Cu is deposited. How long was the current applied? What mass of gold and silver were deposited?

The cell Pt(s) | Cu+(1 M), Cu2+ (1 M) | | Cu+ (1 M) | Cu(s) has Eo = 0.364 V. The cell Cu(s)!Cu2+(1 M)! !Cu+(1 M)!Cu(s) has E! = 0.182 V. Write the cell reaction for each cell and explain the differences in Eo. Calculate for each cell reaction to help explain these differences.

An electrochemical cell has a positive standard cell potential but a negative cell potential. Which statement is true for the cell? a. K > 1; Q > K b. K < 1; Q > K c. K > 1; Q < K d. K < 1; Q < K

Which oxidizing agent will oxidize Br- but not Cl-? a. K2Cr2O7 (in acid) b. KMnO4 (in acid) c. HNO3

A redox reaction employed in an electrochemical cell has a negative \(\Delta G_{\mathrm{rxn}}^{\circ}\). Which statement is true? a. \(E_{\text {cell }}^{\circ}\) is positive; K < 1 b. \(E_{\text {cell }}^{\circ}\) is positive; K > 1 c. \(E_{\text {cell }}^{\circ}\) is negative; K > 1 d. \(E_{\text {cell }}^{\circ}\) is negative; K < 1

A redox reaction has an equilibrium constant of K = 0.055. What is true of and for this reaction?