Chapter 20

Write balanced equations for the following half-reactions. Specify whether each is an oxidation or reduction. (a) \(\operatorname{Cr}(\mathrm{s}) \longrightarrow \mathrm{Cr}^{3+}(\mathrm{aq}) \quad\) (in acid) (b) \(\mathrm{AsH}_{3}(\mathrm{g}) \longrightarrow \mathrm{As}(\mathrm{s}) \quad\) (in acid) (c) \(\mathrm{VO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{V}^{2+}(\mathrm{aq}) \quad\) (in acid) (d) \(\mathrm{Ag}(\mathrm{s}) \longrightarrow \mathrm{Ag}_{2} \mathrm{O}(\mathrm{s})\) (in base)

Write balanced equations for the following half-reactions. Specify whether each is an oxidation or reduction. (a) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) \longrightarrow \mathrm{O}_{2}(\mathrm{g}) \quad\) (in acid) (b) \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \longrightarrow \mathrm{CO}_{2}(\mathrm{g}) \quad\) (in acid) (c) \(\mathrm{NO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{NO}(\mathrm{g})\) (in acid) (d) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq}) \longrightarrow \mathrm{MnO}_{2}(\mathrm{s})\) (in base)

Balance the following redox equations. All occur in acid solution. (a) \(\mathrm{Ag}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{NO}_{2}(\mathrm{g})+\mathrm{Ag}^{+}(\mathrm{aq})\) (b) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{HSO}_{3}^{-}(\mathrm{aq}) \longrightarrow\) \(\mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})\) (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{N}_{2} \mathrm{O}(\mathrm{g})\) (d) \(\operatorname{Cr}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{NO}(\mathrm{g})\)

Balance the following redox equations. All occur in acid solution. (a) \(\mathrm{Sn}(\mathrm{s})+\mathrm{H}^{+}(\mathrm{aq}) \longrightarrow \mathrm{Sn}^{2+}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (b) \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{Fe}^{3+}(\mathrm{aq})\) (c) \(\mathrm{MnO}_{2}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{Cl}_{2}(\mathrm{g})\) (d) \(\mathrm{CH}_{2} \mathrm{O}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \longrightarrow \mathrm{HCO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})\)

Balance the following redox equations. All occur in basic solution. (a) \(\mathrm{Al}(\mathrm{s})+\mathrm{OH}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Al}(\mathrm{OH})_{4}^{-}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (b) \(\mathrm{CrO}_{4}^{2-}(\mathrm{aq})+\mathrm{SO}_{3}^{2-}(\mathrm{aq}) \longrightarrow\) \(\mathrm{Cr}(\mathrm{OH})_{3}(\mathrm{s})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})\) (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}(\mathrm{OH})_{2}(\mathrm{s}) \longrightarrow\left[\mathrm{Zn}(\mathrm{OH})_{4}\right]^{2-}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})\) (d) \(\mathrm{HS}^{-}(\mathrm{aq})+\mathrm{ClO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{S}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq})\)

Balance the following redox equations. All occur in basic solution. (a) \(\operatorname{Fe}(\text { OH })_{3}(s)+\operatorname{Cr}(s) \longrightarrow \operatorname{Cr}(\text { OH })_{3}(s)+\operatorname{Fe}(\text { OH })_{2}(s)\) (b) \(\mathrm{NiO}_{2}(\mathrm{s})+\mathrm{Zn}(\mathrm{s}) \longrightarrow \mathrm{Ni}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{s})\) (c) \(\mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{CrO}_{4}^{2-}(\mathrm{aq}) \longrightarrow\) \(\mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s})+\left[\mathrm{Cr}(\mathrm{OH})_{4}\right]^{-}(\mathrm{aq})\) (d) \(\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+\mathrm{Ag}_{2} \mathrm{O}(\mathrm{s}) \longrightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{Ag}(\mathrm{s})\)

A voltaic cell is constructed using the reaction of chromium metal and iron(II) ion. $$2 \mathrm{Cr}(\mathrm{s})+3 \mathrm{Fe}^{2+}(\mathrm{aq}) \longrightarrow 2 \mathrm{Cr}^{3+}(\mathrm{aq})+3 \mathrm{Fe}(\mathrm{s})$$ Complete the following sentences: Electrons in the external circuit flow from the _____ electrode to the _____ electrode. Negative ions move in the salt bridge from the _____ half-cell to the half-cell. The half-reaction at the anode is _____ and that at the cathode is ____.

voltaic cell is constructed using the reaction \(\mathrm{Mg}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq}) \longrightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (a) Write equations for the oxidation and reduction halfreactions. (b) Which half-reaction occurs in the anode compartment and which occurs in the cathode compartment? (c) Complete the following sentences: Electrons in the external circuit flow from the ___ electrode to the ___ electrode. Negative ions move in the salt bridge from the ___ half-cell to the ___ half-cell. The half-reaction at the anode is ___ and that at the cathode is ___.

The half-cells \(\mathrm{Fe}^{2+}(\text { aq }) | \mathrm{Fe}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}) | \mathrm{H}_{2} \mathrm{O}\) (in acid solution) are linked to create a voltaic cell. (a) Write equations for the oxidation and reduction halfreactions and for the overall (cell) reaction. (b) Which half-reaction occurs in the anode compartment and which occurs in the cathode compartment? (c) Complete the following sentences: Electrons in the external circuit flow from the ___ electrode to the ___ electrode. Negative ions move in the salt bridge from the ___ half-cell to the ___ half-cell.

The half-cells \(\mathrm{Ag}^{+}(\mathrm{aq}) | \mathrm{Ag}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{g}) | \mathrm{Cl}^{-}(\mathrm{aq})\) are linked to create a voltaic cell. (a) Write equations for the oxidation and reduction halfreactions and for the overall (cell) reaction. (b) Which half-reaction occurs in the anode compartment and which occurs in the cathode compartment? (c) Complete the following sentences: Electrons in the external circuit flow from the ___ electrode to the ___ electrode. Negative ions move in the salt bridge from the ___ half-cell to the ___ half-cell.

What are the similarities and differences between dry cells, alkaline batteries, and ni-cad batteries?

What reactions occur when a lead storage battery is recharged?

Balance each of the following unbalanced equations, then calculate the standard potential, \(E^{\circ},\) and decide whether each is product-favored as written. (All reactions occur in acid solution.) (a) \(\mathrm{Sn}^{2+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s}) \longrightarrow \mathrm{Sn}(\mathrm{s})+\mathrm{Ag}^{+}(\mathrm{aq})\) (b) \(\mathrm{Al}(\mathrm{s})+\mathrm{Sn}^{4+}(\mathrm{aq}) \longrightarrow \mathrm{Sn}^{2+}(\mathrm{aq})+\mathrm{Al}^{3+}(\mathrm{aq})\) (c) \(\mathrm{ClO}_{3}^{-}(\mathrm{aq})+\mathrm{Ce}^{3+}(\mathrm{aq}) \longrightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Ce}^{4+}(\mathrm{aq})\) (d) \(\mathrm{Cu}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{NO}(\mathrm{g})\)

Balance each of the following unbalanced equations, then calculate the standard potential, \(E^{\circ},\) and decide whether each is product-favored as written. (All reactions occur in acid solution.) (a) \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{Br}^{-}(\mathrm{aq}) \longrightarrow \mathrm{I}^{-}(\mathrm{aq})+\mathrm{Br}_{2}(\ell)\) (b) \(\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Cu}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Fe}^{3+}(\mathrm{aq})\) (c) \(\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq}) \longrightarrow \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Cr}^{3+}(\mathrm{aq})\) (d) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{HNO}_{2}(\mathrm{aq}) \longrightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{NO}_{3}^{-}(\mathrm{aq})\)

Which of the following elements is the best reducing agent? (a) Cu (b) \(\mathrm{Zn}\) (c) \(\mathrm{Fe}\) (d) \(\mathrm{Ag}\) (e) \(\mathrm{Cr}\)

From the following list, identify those elements that are easier to oxidize than \(\mathrm{H}_{2}(\mathrm{g})\) (a) \(\mathrm{Cu}\) (b) \(\mathrm{Zn}\) (c) \(\mathrm{Fe}\) (d) \(\mathrm{Ag}\) (e) \(\mathrm{Cr}\)

Which of the following ions is most easily reduced? (a) \(\mathrm{Cu}^{2+}(\mathrm{aq})\) (b) \(\mathrm{Zn}^{2+}(\mathrm{aq})\) (c) \(\mathrm{Fe}^{2+}(\mathrm{aq})\) (d) \(\mathrm{Ag}^{+}(\mathrm{aq})\) (e) \(\mathrm{Al}^{3+}(\mathrm{aq})\)

(a) Which halogen is most easily reduced: \(\mathrm{F}_{2}, \mathrm{Cl}_{2}, \mathrm{Br}_{2}\) or \(\mathrm{I}_{2}\) in acidic solution. (b) Identify the halogens that are better oxidizing agents than \(\mathrm{MnO}_{2}(\mathrm{s})\) in acidic solution.

One half-cell in a voltaic cell is constructed from a copper wire dipped into a \(4.8 \times 10^{-3} \mathrm{M}\) solution of \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\) The other half-cell consists of a zinc electrode in a \(0.40 \mathrm{M}\) solution of \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2} .\) Calculate the cell potential.

Which product, \(\mathrm{O}_{2}\) or \(\mathrm{F}_{2}\), is more likely to form at the anode in the electrolysis of an aqueous solution of KF? Explain your reasoning.

Which product, Ca or \(\mathrm{H}_{2}\), is more likely to form at the cathode in the electrolysis of \(\mathrm{CaCl}_{2} ?\) Explain your reasoning.

An aqueous solution of KBr is placed in a beaker with two inert platinum electrodes. When the cell is attached to an external source of electrical energy, electrolysis occurs. (a) Hydrogen gas and hydroxide ion form at the cathode. Write an equation for the half-reaction that occurs at this electrode. (b) Bromine is the primary product at the anode. Write an equation for its formation.

An aqueous solution of \(\mathrm{Na}_{2} \mathrm{S}\) is placed in a beaker with two inert platinum electrodes. When the cell is attached to an external battery, electrolysis occurs. (a) Hydrogen gas and hydroxide ion form at the cathode. Write an equation for the half-reaction that occurs at this electrode. (b) Sulfur is the primary product at the anode. Write an equation for its formation.

In the electrolysis of a solution containing \(\mathrm{Ni}^{2+}\) (aq), metallic \(\mathrm{Ni}(\mathrm{s})\) deposits on the cathode. Using a current of 0.150 A for 12.2 min, what mass of nickel will form?

Electrolysis of a solution of \(\mathrm{CuSO}_{4}(\mathrm{aq})\) to give copper metal is carried out using a current of 0.66 A. How long should electrolysis continue to produce \(0.50 \mathrm{g}\) of copper?

Electrolysis of a solution of \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\) to give zinc metal is carried out using a current of 2.12 A. How long should electrolysis continue to prepare \(2.5 \mathrm{g}\) of zinc?

A voltaic cell can be built using the reaction between Al metal and \(\mathrm{O}_{2}\) from the air. If the Al anode of this cell consists of \(84 \mathrm{g}\) of aluminum, how many hours can the cell produce 1.0 A of electricity, assuming an unlimited supply of \(\mathrm{O}_{2} ?\)

Assume the specifications of a Ni-Cd voltaic cell include delivery of 0.25 A of current for \(1.00 \mathrm{h}\). What is the minimum mass of the cadmium that must be used to make the anode in this cell?

Write balanced equations for the following half-reactions. (a) \(\mathrm{UO}_{2}^{+}(\mathrm{aq}) \longrightarrow \mathrm{U}^{4+}(\mathrm{aq}) \quad\) (acid solution) (b) \(\mathrm{ClO}_{3}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Cl}^{-}(\mathrm{aq}) \quad\) (acid solution) (c) \(\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq}) \longrightarrow \mathrm{N}_{2}(\mathrm{g})\) (basic solution) (d) \(\mathrm{ClO}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Cl}^{-}(\mathrm{aq}) \quad\) (basic solution)

Balance the following equations. (a) \(\mathrm{Zn}(\mathrm{s})+\mathrm{VO}^{2+}(\mathrm{aq}) \longrightarrow\) $$\mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{V}^{3+}(\mathrm{aq}) \text { (acid solution) }$$ (b) \(\mathrm{Zn}(\mathrm{s})+\mathrm{VO}_{3}^{-}(\mathrm{aq}) \longrightarrow\) $$\mathrm{V}^{2+}(\mathrm{aq})+\mathrm{Zn}^{2+}(\mathrm{aq}) \text { (acid solution) }$$ (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{ClO}^{-}(\mathrm{aq}) \longrightarrow\) $$\mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \text { (basic solution) }$$ (d) \(\mathrm{ClO}^{-}(\mathrm{aq})+\left[\mathrm{Cr}(\mathrm{OH})_{4}\right]^{-}(\mathrm{aq}) \longrightarrow\) $$\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{CrO}_{4}^{2-}(\mathrm{aq}) \text { (basic solution) }$$

Magnesium metal is oxidized and silver ions are reduced in a voltaic cell using \(\mathrm{Mg}^{2+}(\mathrm{aq}, 1 \mathrm{M}) | \mathrm{Mg}\) and \(\mathrm{Ag}^{+}(\mathrm{aq}, 1 \mathrm{M}) |\) Ag half-cells. (IMAGE NOT COPY) (a) Label each part of the cell. (b) Write equations for the half-reactions occurring at the anode and the cathode, and write an equation for the net reaction in the cell. (c) Trace the movement of electrons in the external circuit. Assuming the salt bridge contains \(\mathrm{NaNO}_{3},\) trace the movement of the \(\mathrm{Na}^{+}\) and \(\mathrm{NO}_{3}^{-}\) ions in the salt bridge that occurs when a voltaic cell produces current. Why is a salt bridge required in a cell?

The reaction occurring in the cell in which \(\mathrm{Al}_{2} \mathrm{O}_{3}\) and aluminum salts are electrolyzed is \(\mathrm{Al}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \longrightarrow\) \(\mathrm{Al}(\mathrm{s}) .\) If the electrolysis cell operates at \(5.0 \mathrm{V}\) and \(1.0 \mathrm{X}\) \(10^{5} \mathrm{A},\) what mass of aluminum metal can be produced in a 24-h day?

A potential of \(+0.146 \mathrm{V}\) is recorded (under standard conditions) for a voltaic cell constructed using the following half-reactions: Anode: \(\mathrm{Ag}(\mathrm{s}) \longrightarrow \mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{e}^{-}\) Cathode: \(\mathrm{Ag}_{2} \mathrm{SO}_{4}(\mathrm{s})+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Ag}(\mathrm{s})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})\) (a) What is the standard reduction potential for the cathode reaction? (b) Calculate the solubility product, \(K_{\mathrm{sp}},\) for \(\mathrm{Ag}_{2} \mathrm{SO}_{4}\)

The standard voltage, \(E^{\circ},\) for the reaction of \(\mathrm{Zn}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{g})\) is \(+2.12 \mathrm{V}\). What is the standard free energy change, \(\overline{\Delta G}^{\circ},\) for the reaction?

The standard potential for the reaction of \(\mathrm{Mg}(\mathrm{s})\) with \(\mathrm{I}_{2}(\mathrm{s})\) is \(+2.91 \mathrm{V} .\) What is the standard free energy change, \(\Delta G^{\circ}\) for the reaction?

An electrolysis cell for aluminum production operates at \(5.0 \mathrm{V}\) and a current of \(1.0 \times 10^{5} \mathrm{A}\). Calculate the number of kilowatt- hours of energy required to produce 1 metric ton \(\left(1.0 \times 10^{3} \mathrm{kg}\right)\) of aluminum. \((1 \mathrm{kWh}=\) \(3.6 \times 10^{6} \mathrm{J}\) and \(1 \mathrm{J}=1 \mathrm{C} \cdot \mathrm{V} .\)

A current of \(0.0100 \mathrm{A}\) is passed through a solution of rhodium sulfate, causing reduction of the metal ion to the metal. After \(3.00 \mathrm{h}, 0.038 \mathrm{g}\) of Rhas been deposited. What is the charge on the rhodium ion, \(R h^{n+} ?\) What is the formula for rhodium sulfate?

A current of \(0.44 \mathrm{A}\) is passed through a solution of ruthenium nitrate causing reduction of the metal ion to the metal. After \(25.0 \mathrm{min}, 0.345 \mathrm{g}\) of Ru has been deposited. What is the charge on the ruthenium ion, \(\mathrm{Ru}^{n+} ?\) What is the formula for ruthenium nitrate?

Chlorine gas is obtained commercially by electrolysis of brine (a concentrated aqueous solution of \(\mathrm{NaCl}\) ). If the electrolysis cells operate at \(4.6 \mathrm{V}\) and \(3.0 \times 10^{5} \mathrm{A},\) what mass of chlorine can be produced in a 24 -h day?

An old method of measuring the current flowing in a circuit was to use a "silver coulometer." The current passed first through a solution of \(\mathrm{Ag}^{+}(\mathrm{aq})\) and then into another solution containing an electroactive species. The amount of silver metal deposited at the cathode was weighed. From the mass of silver, the number of atoms of silver was calculated. since the reduction of a silver ion requires one electron, this value equalled the number of electrons passing through the circuit. If the time was noted, the average current could be calculated. If, in such an experiment, \(0.052 \mathrm{g}\) of \(\mathrm{Ag}\) is deposited during \(450 \mathrm{s}\), what was the current flowing in the circuit?

A "silver coulometer" (Study Question 72) was used in the past to measure the current flowing in an electrochemical cell. Suppose you found that the current flowing through an electrolysis cell deposited \(0.089 \mathrm{g}\) of \(\mathrm{Ag}\) metal at the cathode after exactly 10 min. If this same current then passed through a cell containing gold(III) ion in the form of \(\left(\mathrm{AuCl}_{4}\right)^{-}\), how much gold was deposited at the cathode in that electrolysis cell?

A Write balanced equations for the following reduction half-reactions involving organic compounds. (a) \(\mathrm{HCO}_{2} \mathrm{H} \longrightarrow \mathrm{CH}_{2} \mathrm{O} \quad\) (acid solution) (b) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H} \longrightarrow \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CH}_{3} \quad\) (acid solution) (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CHO} \longrightarrow \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH} \quad\) (acid solution) (d) \(\mathrm{CH}_{3} \mathrm{OH} \longrightarrow \mathrm{CH}_{4} \quad\) (acid solution)

A Balance the following equations involving organic compounds. (a) \(\mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CHO}(\mathrm{aq}) \longrightarrow\) \(\mathrm{Ag}(\mathrm{s})+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H}(\mathrm{aq}) \quad\) (acid solution) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq}) \longrightarrow\) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Cr}^{3+}(\mathrm{aq}) \quad\) (acid solution)

A voltaic cell is constructed in which one half-cell consists of a silver wire in an aqueous solution of AgNO \(_{3} .\) The other half-cell consists of an inert platinum wire in an aqueous solution containing \(\mathrm{Fe}^{2+}(\mathrm{aq})\) and \(\mathrm{Fe}^{3+}(\mathrm{aq})\) (a) Calculate the voltage of the cell, assuming standard conditions. (b) Write the net ionic equation for the reaction occurring in the cell. (c) In this voltaic cell, which electrode is the anode and which is the cathode? (d) If \(\left[\mathrm{Ag}^{+}\right]\) is \(0.10 \mathrm{M},\) and \(\left[\mathrm{Fe}^{2+}\right]\) and \(\left[\mathrm{Fe}^{3+}\right]\) are both 1.0 M, what is the cell voltage? Is the net cell reaction still that used in part (a)? If not, what is the net reaction under the new conditions?

A solution of KI is added dropwise to a pale blue solution of \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2} .\) The solution changes to a brown color and a precipitate forms. In contrast, no change is observed if solutions of KCl and KBr are added to aqueous \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2} .\) Consult the table of standard reduction potentials to explain the dissimilar results seen with the different halides. Write an equation for the reaction that occurs when solutions of KI and \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\) are mixed.

Four metals, \(A, B, C,\) and \(D,\) exhibit the following properties: (a) Only A and C react with 1.0 M hydrochloric acid to give \(\mathrm{H}_{2}(\mathrm{g})\) (b) When \(\mathrm{C}\) is added to solutions of the ions of the other metals, metallic \(\mathbf{B}, \mathbf{D},\) and \(\mathbf{A}\) are formed. (c) Metal D reduces \(B^{n+}\) to give metallic \(B\) and \(D^{n+}\) Based on this information, arrange the four metals in order of increasing ability to act as reducing agents.

A hydrogen-oxygen fuel cell operates on the simple reaction $$\mathbf{H}_{2}(\mathrm{g})+\frac{1}{2} \mathrm{O}_{2}(\mathrm{g}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(\ell)$$ If the cell is designed to produce 1.5 A of current, and if the hydrogen is contained in a 1.0 -L tank at \(200 .\) atm pressure at \(25^{\circ} \mathrm{C},\) how long can the fuel cell operate before the hydrogen runs out? (Assume there is an unlimited supply of \(\left.\mathbf{O}_{2 \cdot}\right)\)

living organisms derive energy from the oxidation of food, typified by glucose. $$\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(\mathrm{aq})+6 \mathrm{O}_{2}(\mathrm{g}) \longrightarrow 6 \mathrm{CO}_{2}(\mathrm{g})+6 \mathrm{H}_{2} \mathrm{O}(\ell)$$ Electrons in this redox process are transferred from glucose to oxygen in a series of at least 25 steps. It is instructive to calculate the total daily current flow in a typical organism and the rate of energy expenditure (power). (See T.P. Chirpich: Journal of Chemical Education, Vol. 52 , p. \(99,1975 .)\) (a) The molar enthalpy of combustion of glucose is \(-2800 \mathrm{kJ} .\) If you are on a typical daily diet of \(2400 \mathrm{Cal}\) (kilocalories), what amount of glucose (in moles) must be consumed in a day if glucose is the only source of energy? What amount of \(\mathrm{O}_{2}\) must be consumed in the oxidation process? (b) How many moles of electrons must be supplied to reduce the amount of \(\mathrm{O}_{2}\) calculated in part (a)? (c) Based on the answer in part (b), calculate the current flowing, per second, in your body from the combustion of glucose. (d) If the average standard potential in the electron transport chain is \(1.0 \mathrm{V},\) what is the rate of energy expenditure in watts?