Determine whether HI can dissolve each metal sample. If it can,write a balanced chemical

How can Table 20.1 be used to predict whether or not a metal will dissolve in HCl? In \(HNO_{3}? Text Transcription: HNO_3

Explain why \(E°_{cell}\), \(?G°_{rxn}\), and K are all interrelated. Text Transcription: E°_cell ?G°_rxn

Does a redox reaction with a small equilibrium constant (K 6 1) have a positive or a negative \(E°_{cell}\)? Does it have a positive or a negative \(?G°_{rxn}\)? Text Transcription: E°_cell ?G°_rxn

How does \(E_{cell}\) 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 \(E_{cell}\)? Increasing the concentration of a product? Text Transcription: E_cell

Use the Nernst equation to show that \(E_{cell} = E°_{cell}\) under standard conditions. Text Transcription: E_cell = E°_cell

Balance each redox reaction occurring in acidic aqueous solution. a. \(K(s) + Cr^{3+}(aq) \rightarrow Cr(s) + K+(aq)\) b. \(Al(s) + Fe^{2+}(aq) \rightarrow Al^{3+}(aq) + Fe(s)\) c. \(BrO_{3}^{-}(aq) + N_{2}H_{4}( g) \rightarrow Br^{-}(aq) + N_{2}( g)\) Text Transcription: K(s) + Cr^3+ (aq) Cr(s) + K^+ (aq) Al(s) + Fe^2+ (aq) Al^3+ (aq) + Fe(s) BrO_3^- (aq) + N_2H_4(g) Br^- (aq) + N_2(g)

Balance each redox reaction occurring in acidic aqueous solution. a. \(Zn(s) + Sn_{2}+(aq) \rightarrow Zn_{2}+(aq) + Sn(s)\) b. \(Mg(s) + Cr^{3+}(aq) \rightarrow Mg^{2+}(aq) + Cr(s)\) c. \(MnO_{4}^{-}(aq) + Al(s) \rightarrow Mn^{2+}(aq) + Al^{3+}(aq) \) Text Transcription: Zn(s) + Sn_2+(aq) Zn_2+(aq) + Sn(s) Mg(s) + Cr^3+ (aq) Mg^2+(aq) + Cr(s) MnO_4^-(aq) + Al(s) Mn^2+(aq) + Al^3+(aq)

Balance each redox reaction occurring in acidic aqueous solution. a. \(PbO_{2}(s) + I^{-}(aq) \rightarrow Pb^{2+}(aq) + I_{2}(s)\) b. \(SO_{3}^{2-}(aq) + MnO_{4}^{-}(aq) \rightarrow SO_{4}^{2-}(aq) + Mn^{2+}(aq) \) c. \(S_{2}O_{3}^{2-}(aq) + Cl_{2}( g) \rightarrow SO_{4}^{2-}(aq) + Cl^{-}(aq)\) Text Transcription: PbO_2(s) + I^-(aq) Pb^2+(aq) + I_2(s) SO_3^2-(aq) + MnO_4^-(aq) SO_4^{2-}(aq) + Mn^2+(aq) S_2O_3^2-(aq) + Cl_2( g) SO_4^2-(aq) + C^l-(aq)

Balance each redox reaction occurring in acidic aqueous solution. a. \(I^{-}(aq) + NO_{2}^{-}(aq) \rightarrow I_{2}(s) + NO( g)\) b. \(ClO_{4}^{-}(aq) + Cl^{-}(aq) \rightarrow ClO_{3}^{-}(aq) + Cl_{2}( g)\) c. \(NO_{3}^{-}(aq) + Sn^{2+}(aq) \rightarrow Sn^{4+}(aq) + NO( g)\) Text Transcription: I^-(aq) + NO_2^-(aq) I_2(s) + NO( g) ClO_4^-(aq) + Cl^-(aq) ClO_3^-(aq) + Cl_2( g) NO_3^-(aq) + Sn^2+(aq) Sn^4+(aq) + NO( g)

Balance each redox reaction occurring in basic aqueous solution. a. \(H_{2}O_{2}(aq) + ClO_{2}(aq) \rightarrow ClO_{2}^{-}(aq) + O_{2}( g) \) b. \(Al(s) + MnO_{4}^{-}(aq) \rightarrow MnO_{2}(s) + Al(OH)_{4}^{-}(aq)\) c. \(Cl_{2}( g) \rightarrow Cl^{-}(aq) + ClO^{-}(aq)\) Text Transcription: H_2O_2(aq) + ClO_2(aq) ClO_2^-(aq) + O_2( g) Al(s) + MnO_4^-(aq) MnO_2(s) + Al(OH)_4^-(aq) Cl_2( g) \rightarrow Cl^-(aq) + ClO^-(aq)

Balance each redox reaction occurring in basic aqueous solution. a. \(MnO_4^-(aq) + Br^-(aq) \rightarrow MnO_2(s) + BrO_3^-(aq)\) b. \(Ag(s) + CN^-(aq) + O_2( g) \rightarrow Ag(CN)_2^-(aq)\) c. \(NO_2^-(aq) + Al(s) \rightarrow NH_3( g) + AlO_2^-(aq)\) Text Transcription: MnO_4^-(aq) + Br^-(aq) MnO_2(s) + BrO_3^-(aq) Ag(s) + CN^-(aq) + O_2( g) Ag(CN)_2^-(aq) NO_2^-(aq) + Al(s) NH_3( g) + AlO_2^-(aq)

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. a. \(2 Ag^{+}(aq) + Pb(s) \rightarrow 2 Ag(s) + Pb^{2+}(aq)\) b. \(2 ClO_{2}( g) + 2 I^{-}(aq) \rightarrow 2 ClO_{2}^{-}(aq) + I_{2}(s)\) c. \(O_{2}( g) + 4 H^{+}(aq) + 2 Zn(s) \rightarrow 2 H_{2}O(l) + 2 Zn^{2+}(aq)\) Text Transcription: 2 Ag^+(aq) + Pb(s) 2 Ag(s) + Pb^2+(aq) 2 ClO_2( g) + 2 I^-(aq) 2 ClO_2^-(aq) + I_2(s) O_2( g) + 4 H^+(aq) + 2 Zn(s) 2 H_2O(l) + 2 Zn^2+(aq)

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. a. \(Ni^{2+}(aq) + Mg(s) \rightarrow Ni(s) + Mg^2+(aq)\) b. \(2 H^{+}(aq) + Fe(s) \rightarrow H_{2}( g) + Fe^{2+}(aq)\) c. \(2 NO_{3}^{-}(aq) + 8 H^{+}(aq) + 3 Cu(s) \rightarrow 2 NO( g) + 4 H_{2}O(l) + 3 Cu^{2+}(aq)\) Text Transcription: Ni^2+(aq) + Mg(s) Ni(s) + Mg^2+(aq) 2 H^+(aq) + Fe(s) H_2( g) + Fe^2+(aq) 2 NO_3^-(aq) + 8 H^+(aq) + 3 Cu(s) 2 NO( g) + 4 H_2O(l) + 3 Cu^2+(aq)

Calculate the standard cell potential for each of the electrochemical cells in Problem 43. a. \(2 Ag^{+}(aq) + Pb(s) \rightarrow 2 Ag(s) + Pb^{2+}(aq)\) b. \(2 ClO_{2}( g) + 2 I^{-}(aq) \rightarrow 2 ClO_{2}^{-}(aq) + I_{2}(s)\) c. \(O_{2}( g) + 4 H^{+}(aq) + 2 Zn(s) \rightarrow 2 H_{2}O(l) + 2 Zn^{2+}(aq)\) Text Transcription: 2 Ag^+(aq) + Pb(s) 2 Ag(s) + Pb^2+(aq) 2 ClO_2( g) + 2 I^-(aq) 2 ClO_2^-(aq) + I_2(s) O_2( g) + 4 H^+(aq) + 2 Zn(s) 2 H_2O(l) + 2 Zn^2+(aq)

Calculate the standard cell potential for each of the electrochemical cells in Problem 44. a. \(Ni^{2+}(aq) + Mg(s) \rightarrow Ni(s) + Mg^2+(aq)\) b. \(2 H^{+}(aq) + Fe(s) \rightarrow H_{2}( g) + Fe^{2+}(aq)\) c. \(2 NO_{3}^{-}(aq) + 8 H^{+}(aq) + 3 Cu(s) \rightarrow 2 NO( g) + 4 H_{2}O(l) + 3 Cu^{2+}(aq)\) Text Transcription: Ni^2+(aq) + Mg(s) Ni(s) + Mg^2+(aq) 2 H^+(aq) + Fe(s) H_2( g) + Fe^2+(aq) 2 NO_3^-(aq) + 8 H^+(aq) + 3 Cu(s) 2 NO( g) + 4 H_2O(l) + 3 Cu^2+(aq)

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 \(E°_{cell}\). c. Label each electrode as negative or positive. d. Indicate the direction of anion and cation flow in the salt bridge. Text Transcription: E°_cell

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 \(E°_{cell}\). c. Label each electrode as negative or positive. d. Indicate the direction of anion and cation flow in the salt bridge. Text Transcription: E°_cell

Use line notation to represent each electrochemical cell in Problem 43. a. \(2 Ag^{+}(aq) + Pb(s) \rightarrow 2 Ag(s) + Pb^{2+}(aq)\) b. \(2 ClO_{2}( g) + 2 I^{-}(aq) \rightarrow 2 ClO_{2}^{-}(aq) + I_{2}(s)\) c. \(O_{2}( g) + 4 H^{+}(aq) + 2 Zn(s) \rightarrow 2 H_{2}O(l) + 2 Zn^{2+}(aq)\) Text Transcription: 2 Ag^+(aq) + Pb(s) 2 Ag(s) + Pb^2+(aq) 2 ClO_2( g) + 2 I^-(aq) 2 ClO_2^-(aq) + I_2(s) O_2( g) + 4 H^+(aq) + 2 Zn(s) 2 H_2O(l) + 2 Zn^2+(aq)

Use line notation to represent each electrochemical cell in Problem 44. a. \(Ni^{2+}(aq) + Mg(s) \rightarrow Ni(s) + Mg^2+(aq)\) b. \(2 H^{+}(aq) + Fe(s) \rightarrow H_{2}( g) + Fe^{2+}(aq)\) c. \(2 NO_{3}^{-}(aq) + 8 H^{+}(aq) + 3 Cu(s) \rightarrow 2 NO( g) + 4 H_{2}O(l) + 3 Cu^{2+}(aq)\) Text Transcription: Ni^2+(aq) + Mg(s) Ni(s) + Mg^2+(aq) 2 H^+(aq) + Fe(s) H_2( g) + Fe^2+(aq) 2 NO_3^-(aq) + 8 H^+(aq) + 3 Cu(s) 2 NO( g) + 4 H_2O(l) + 3 Cu^2+(aq)

Make a sketch of the voltaic cell represented by the line notation. Write the overall balanced equation for the reaction and calculate \(E°_{cell}\). \(Sn(s) | Sn^{2+}(aq) || NO( g) NO_{3}^{-}(aq), H^{+}(aq) | Pt(s)\) Text Transcription: E°_cell Sn(s) I Sn^2+(aq) II NO( g) NO_3^-(aq), H^+(aq) | Pt(s)

Make a sketch of the voltaic cell represented by the line notation. Write the overall balanced equation for the reaction and calculate \(E°_{cell}\). \(Mn(s) | Mn^{2+}(aq) || ClO^{2-}(aq) ClO_2( g) | Pt(s)\) Text Transcription: E°_cell Mn(s) | Mn^2+(aq) || ClO^2-(aq) ClO_2( g) | Pt(s)

Determine whether or not each redox reaction occurs spontaneously in the forward direction. a. \(Ni(s) + Zn^{2+}(aq) \rightarrow Ni^{2+}(aq) + Zn(s)\) b. \(Ni(s) + Pb^{2+}(aq) \rightarrow Ni^{2+}(aq) + Pb(s)\) c. \(Al(s) + 3 Ag^{+}(aq) \rightarrow Al^{3+}(aq) + 3 Ag(s)\) d. \(Pb(s) + Mn^{2+}(aq) \rightarrow Pb^{2+}(aq) + Mn(s)\) Text Transcription: Ni(s) + Zn^2+(aq) Ni^2+(aq) + Zn(s) Ni(s) + Pb^2+(aq) Ni^2+(aq) + Pb(s) Al(s) + 3 Ag^+(aq) Al^3+(aq) + 3 Ag(s) Pb(s) + Mn^2+(aq) Pb^2+(aq) + Mn(s)

Determine whether or not each redox reaction occurs spontaneously in the forward direction. a. \(Ca^{2+}(aq) + Zn(s) \rightarrow Ca(s) + Zn^{2+}(aq)\) b. \(2 Ag^{+}(aq) + Ni(s) \rightarrow 2 Ag(s) + Ni^{2+}(aq)\) c. \(Fe(s) + Mn^{2+}(aq) \rightarrow Fe^{2+}(aq) + Mn(s)\) d. \(2 Al(s) + 3 Pb^{2+}(aq) \rightarrow 2 Al^{3+}(aq) + 3 Pb(s)\) Text Transcription: Ca^2+(aq) + Zn(s) Ca(s) + Zn^2+(aq) 2 Ag^+(aq) + Ni(s) 2 Ag(s) + Ni^2+(aq) Fe(s) + Mn^2+(aq) Fe^2+(aq) + Mn(s) 2 Al(s) + 3 Pb^2+(aq) 2 Al^3+(aq) + 3 Pb(s)

Which metal could you use to reduce \(Mn^{2+}\) ions but not \(Mg^{2+}\) ions? Text Transcription: Mn^2+ Mg^2+

Which metal can be oxidized with an \(Sn^{2+}\) solution but not with an \(Fe^{2+}\) solution? Text Transcription: Sn^2+ Fe^2+

Determine whether or not each metal dissolves in 1 M \(HNO_{3}\). For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Cu b. Au Text Transcription: HNO_3

Determine whether or not each metal dissolves in 1 M \(HNO_{3}\). For those metals that do dissolve, write a balanced redox equation for the reaction that occurs. a. Au b. Cr Text Transcription: HNO_3

Calculate \(E°_{cell}\) for each balanced redox reaction and determine if the reaction is spontaneous as written. a. \(2 Cu(s) + Mn^{2+}(aq) \rightarrow 2 Cu+(aq) + Mn(s)\) b. \(MnO_{2}(aq) + 4 H^{+}(aq) + Zn(s) \rightarrow Mn^{2+}(aq) + 2 H_{2}O(l) + Zn^{2+}(aq)\) c. \(Cl_{2}( g) + 2 F^{-}(aq) \rightarrow F_{2}( g) + 2 Cl^{-}(aq)\) Text Transcription: 2 Cu(s) + Mn^2+(aq) 2 Cu^+(aq) + Mn(s) MnO_2(aq) + 4 H^+(aq) + Zn(s) Mn^2+(aq) + 2 H_2O(l) + Zn^2+(aq) Cl_2( g) + 2 F^-(aq) F_2( g) + 2 Cl^-(aq)

Calculate \(E°_{cell}\) for each balanced redox reaction and determine if the reaction is spontaneous as written. a. \(O_{2}( g) + 2 H_{2}O(l) + 4 Ag(s) \rightarrow 4 OH^{-}(aq) + 4 Ag^{+}(aq)\) b. \(Br_{2}(l) + 2 I^{-}(aq) \rightarrow 2 Br^{-}(aq) + I_{2}(s)\) c. \(PbO_{2}(s) + 4 H^{+}(aq) + Sn(s) \rightarrow Pb^{2+}(aq) + 2 H_{2}O(l) + Sn^{2+}(aq)\) Text Transcription: O_2( g) + 2 H_2O(l) + 4 Ag(s) 4 OH^-(aq) + 4 Ag^+(aq) Br_2(l) + 2 I^-(aq) 2 Br^-(aq) + I_2(s) PbO_2(s) + 4 H^+(aq) + Sn(s) Pb^2+(aq) + 2 H_2O(l) + Sn^2+(aq)

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

Use tabulated electrode potentials to calculate \(?G°_{rxn}\) for each reaction at 25 °C. a. \(Pb^{2+}(aq) + Mg(s) \rightarrow Pb(s) + Mg^{2+}(aq)\) b. \(Br_{2}(l) + 2 Cl^{-}(aq) \rightarrow 2 Br^{-}(aq) + Cl_{2}( g)\) c. \(MnO_{2}(s) + 4 H^{+}(aq) + Cu(s) \rightarrow Mn^{2+}(aq) + 2 H_{2}O(l) + Cu^{2+}(aq)\) Text Transcription: ?G°_rxn Pb^2+(aq) + Mg(s) Pb(s) + Mg^2+(aq) Br_2(l) + 2 Cl^-(aq) 2 Br^-(aq) + Cl_2( g) MnO_2(s) + 4 H^+(aq) + Cu(s) Mn^2+(aq) + 2 H_2O(l) + Cu^2+(aq)

Use tabulated electrode potentials to calculate \(?G°_{rxn}\) for each reaction at 25 °C. a. \(2 Fe^{3+}(aq) + 3 Sn(s) \rightarrow 2 Fe(s) + 3 Sn^{2+}(aq)\) b. \(O_{2}( g) + 2 H_{2}O(l) + 2 Cu(s) \rightarrow 4 OH^{-}(aq) + 2 Cu^{2+}(aq)\) c. \(Br_{2}(l) + 2 I^{-}(aq) \rightarrow 2 Br^{-}(aq) + I_{2}(s)\) Text Transcription: ?G°_rxn 2 Fe^3+(aq) + 3 Sn(s) 2 Fe(s) + 3 Sn^2+(aq) O_2( g) + 2 H_2O(l) + 2 Cu(s) 4 OH^-(aq) + 2 Cu^2+(aq) Br_2(l) + 2 I^-(aq) 2 Br^-(aq) + I_2(s)

Calculate the equilibrium constant for each of the reactions in Problem 65. Use tabulated electrode potentials to calculate \(?G°_{rxn}\) for each reaction at 25 °C. a. \(Pb^{2+}(aq) + Mg(s) \rightarrow Pb(s) + Mg^{2+}(aq)\) b. \(Br_{2}(l) + 2 Cl^{-}(aq) \rightarrow 2 Br^{-}(aq) + Cl_{2}( g)\) c. \(MnO_{2}(s) + 4 H^{+}(aq) + Cu(s) \rightarrow Mn^{2+}(aq) + 2 H_{2}O(l) + Cu^{2+}(aq)\) Text Transcription: ?G°_rxn Pb^2+(aq) + Mg(s) Pb(s) + Mg^2+(aq) Br_2(l) + 2 Cl^-(aq) 2 Br^-(aq) + Cl_2( g) MnO_2(s) + 4 H^+(aq) + Cu(s) Mn^2+(aq) + 2 H_2O(l) + Cu^2+(aq)

Calculate the equilibrium constant for each of the reactions in Problem 66. Use tabulated electrode potentials to calculate \(?G°_{rxn}\) for each reaction at 25 °C. a. \(2 Fe^{3+}(aq) + 3 Sn(s) \rightarrow 2 Fe(s) + 3 Sn^{2+}(aq)\) b. \(O_{2}( g) + 2 H_{2}O(l) + 2 Cu(s) \rightarrow 4 OH^{-}(aq) + 2 Cu^{2+}(aq)\) c. \(Br_{2}(l) + 2 I^{-}(aq) \rightarrow 2 Br^{-}(aq) + I_{2}(s)\) Text Transcription: ?G°_rxn 2 Fe^3+(aq) + 3 Sn(s) 2 Fe(s) + 3 Sn^2+(aq) O_2( g) + 2 H_2O(l) + 2 Cu(s) 4 OH^-(aq) + 2 Cu^2+(aq) Br_2(l) + 2 I^-(aq) 2 Br^-(aq) + I_2(s)

Calculate the equilibrium constant for the reaction between \(Ni^{2+}(aq)\) and Cd(s) (at 25 °C). Text Transcription: Ni^2+(aq)

Calculate the equilibrium constant for the reaction between \(Fe^{2+}(aq)\) and Zn(s) (at 25 °C). Text Transcription: Fe^2+ (aq)

Calculate \(?G?_{rxn}\) and \(E°_{cell}\) for a redox reaction with n = 2 that has an equilibrium constant of K = 25 (at 25 °C). Text Transcription: ?G?_rxn E°_cell

Calculate \(?G?_{rxn}\) and \(E°_{cell}\) for a redox reaction with n = 3 that has an equilibrium constant of K = 0.050 (at 25 °C). Text Transcription: ?G?_rxn E°_cell

A voltaic cell employs the following redox reaction: \(Sn^{2+}(aq) + Mn(s) \rightarrow Sn(s) + Mn^{2+}(aq)\) Calculate the cell potential at 25 °C under each set of conditions. a. standard conditions b. \([Sn^{2+}] = 0.0100 M; [Mn^{2+}] = 2.00 M\) c. \([Sn^{2+}] = 2.00 M; [Mn^{2+}] = 0.0100 M\) Text Transcription: Sn^2+(aq) Sn(s) + Mn^2+(aq) [Sn^2+] = 0.0100 M; [Mn^2+] = 2.00 M [Sn^2+] = 2.00 M; [Mn^2+] = 0.0100 M

A voltaic cell employs the redox reaction: \(2 Fe_{3} +(aq) + 3 Mg(s) \rightarrow 2 Fe(s) + 3 Mg^{2+}(aq)\) Calculate the cell potential at 25 °C under each set of conditions. a. standard conditions b. \([Fe^{3+}] = 1.0 \times 10^{-3} M; [Mg^{2+}] = 2.50 M\) c. \([Fe^{3+}] = 2.00 M; [Mg^{2+}] = 1.5 \times 10^{-3} M\) Text Transcription: 2 Fe^3+(aq) + 3 Mg(s) 2 Fe(s) + 3 Mg^2+(aq) [Fe^3+] = 1.0 x 10^-3 M; [Mg^2+] = 2.50 M [Fe^3+] = 2.00 M; [Mg^2+] = 1.5 x 10^-3 M

An electrochemical cell is based on these two half-reactions: \(Ox: Pb(s) \rightarrow Pb^{2+}(aq, 0.10 M) + 2 e^{-}\) \(Red: MnO_{4}^{-}(aq, 1.50 M) + 4 H^{+}(aq, 2.0 M) + 3 e^{-} \rightarrow MnO_{2}(s) + 2 H_{2}O(l)\) Calculate the cell potential at 25 °C. Text Transcription: Ox: Pb(s) Pb^2+ (aq, 0.10M) + 2 e^- Red: MnO_4^-(aq, 1.50M) + 4H+(aq, 2.0M) + 3e^- MnO_2(s) + 2 H_2O (I)

An electrochemical cell is based on these two half-reactions: \(Ox: Sn(s) \rightarrow Sn^{2+}(aq, 2.00 M) + 2 e^{-}\) \(ClO_{2}( g, 0.100 atm) + e^{-} \rightarrow ClO_{2}^{-}(aq, 2.00 M)\) Calculate the cell potential at 25 °C. Text Transcription: Ox: Sn(s) Sn^2+ (aq, 2.00M) + 2 e^- Red: CIO_2(g, 0.100atm) + e^- MnO_2(s) + 2 H_2O(l)

A voltaic cell consists of a \(Zn/Zn^{2+}\) half-cell and a \(Ni/Ni^{2+}\) halfcell at 25 °C. The initial concentrations of \(Ni^{2+}\) and \(Zn^{2+}\) 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 \(Ni^{2+}\) has fallen to 0.500 M? c. What are the concentrations of \(Ni^{2+}\) and \(Zn^{2+}\) when the cell potential falls to 0.45 V? Text Transcription: Zn/Zn^2+ Ni/Ni^2+ Ni^2+ Zn^2+

A voltaic cell consists of a \(Pb/Pb^{2+}\) half-cell and a \(Cu/Cu^{2+}\) halfcell at 25 °C. The initial concentrations of \(Pb^{2+}\) and \(Cu^{2+}\) 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 \(Cu^{2+}\) has fallen to 0.200 M? c. What are the concentrations of \(Pb^{2+}\) and \(Cu^{2+}\) when the cell potential falls to 0.35 V? Text Transcription: Pb/Pb^2+ Cu/Cu^2+ Pb^2+ Cu^2+

Make a sketch of a concentration cell employing two \(Zn/Zn^{2+}\) halfcells. The concentration of \(Zn2+\) in one of the half-cells is 2.0 M, and the concentration in the other half-cell is \(1.0 \times 10^{-3}\) M. Label the anode and the cathode and indicate the half-reaction occurring at each electrode. Also indicate the direction of electron flow. Text Transcription: Zn/Zn^2+ Zn^2+ 1.0 x 10^-3

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 \(Pb^{2+}\) in each half-cell. Text Transcription: Pb^2+

A concentration cell consists of two \(Sn/Sn^{2+}\) half-cells. The cell has a potential of 0.10 V at 25 °C. What is the ratio of the \(Sn^{2+}\) concentrations in the two half-cells? Text Transcription: Sn/Sn^2+ Sn^2+

A \(Cu/Cu^{2+}\) concentration cell has a voltage of 0.22 V at 25 °C. The concentration of \(Cu^{2+}\) in one of the half-cells is \(1.5 \times 10^{-3} M\). What is the concentration of \(Cu^{2+}\) in the other half-cell? (Assume the concentration in the unknown cell is the lower of the two concentrations.) Text Transcription: Cu/Cu^2 Cu^2+ 1.5 x 10^-3 M

Determine the optimum mass ratio of Zn to \(MnO_{2}\) in an alkaline battery. Text Transcription: MnO_2

Refer to the tabulated values of \(?G°_{f}\) in Appendix IIB to calculate \(E°_{cell}\) for a fuel cell that employs the reaction between methane gas \((CH_{4})\) and oxygen to form carbon dioxide and gaseous water. Text Transcription: ?G°_f E°_cell

Refer to the tabulated values of \(?G°_{f}\) in Appendix IIB to calculate \(E°_{cell}\) for the fuel-cell breathalyzer, which employs the following reaction. (\(?G°_{f}\) for \(HC_{2}H_{3}O_{2}\)( g) = -374.2 kJ/mol.) Text Transcription: ?G°_f E°_cell HC_2H_3O_2

Consider the electrolytic cell: a. Label the anode and the cathode and indicate the halfreactions 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 \(Mn^{2+}\) is reduced to Mn and Sn is oxidized to \(Sn^{2+}\). 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? Text Transcription: Mn^2+ Sn^2+

Write equations for the half-reactions that occur at the anode and cathode for the electrolysis of each aqueous solution. a. NaBr(aq) b. \(PbI_{2}(aq)\) c. \(Na_{2}SO_{4}(aq)\) Text Transcription: PbI_2(aq) Na_2SO_4(aq)

Write equations for the half-reactions that occur at the anode and cathode for the electrolysis of each aqueous solution. a. \(Ni(NO_{3})_{2}(aq) b. KCI(aq) c. \(CuBr_{2}(aq)\) Text Transcription: Ni(NO_3)_2(aq) CuBr_2(aq)

Silver can be electroplated at the cathode of an electrolysis cell by the half-reaction: \(Ag^{+}(aq) + e^{-} \rightarrow Ag(s)\) What mass of silver would plate onto the cathode if a current of 6.8 A flowed through the cell for 72 min? Text Transcription: A^+ (aq) + e^- Ag(s)

Consider the reaction shown here occurring at 25 °C. \(A(s) + B^{2+}(aq) \rightarrow A^{2+}(aq) + B(s)\) \(?G°_{rxn} = -14.0 kJ\) Determine the value of \(E°_{cell}\) and K for the reaction and complete the table. Text Transcription: A(s) + B^2+ (aq) A^2+ (aq) + B(s) ?G°_rxn = -14.0 kJ

Consider the reaction shown here occurring at 25 °C. \(Cr(s) + Cd^{2+}(aq) \rightarrow Cr^{2+}(aq) + Cd(s)\) Determine \(E°_{cell}\), K, and \(?G°_{rxn}\) for the reaction and complete the table. Text Transcription: Cr(s) + Cd^2+(aq) Cr^2+(aq) + Cd(s) E°_cell ?G°_rxn

Consider the unbalanced redox reaction: \(MnO_{4}^{-}(aq) + Zn(s) \rightarrow Mn^{2+}(aq) + Zn^{2+}(aq)\) Balance the equation and determine the volume of a 0.500 M \(KMnO_{4}\) solution required to completely react with 2.85 g of Zn. Text Transcription: MnO_4^-(aq) + Zn(s) Mn^2+(aq) + Zn^2+ (aq) KMnO_4

Consider the unbalanced redox reaction: \(Cr_{2}O_{7}^{2-}(aq) + Cu(s) \rightarrow Cr^{3+}(aq) + Cu^{2+}(aq)\) Balance the equation and determine the volume of a 0.850 M \(K_{2}Cr_{2}O_{7}\) solution required to completely react with 5.25 g of Cu. Text Transcription: Cr_2O_7^2-(aq) + Cu(s) Cr^3+(aq) + Cu^2+(aq) K_2Cr_2O_7

Consider the molecular views of an Al strip and \(Cu^{2+}\) 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. Text Transcription: Cu^2+

Consider the molecular view of an electrochemical cell involving the overall reaction: \(Zn(s) + Ni^{2+}(aq) \rightarrow Zn^{2+}(aq) + Ni(s)\) Draw a similar sketch of the cell after it has generated a substantial amount of electrical current. Text Transcription: Zn(s) + Ni^2+ (aq) Zn^2+(aq) + Ni(s)

Determine if \(HNO_{3}\) 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 \(HNO_{3}\) required to completely dissolve the sample. a. 5.90 g Au b. 2.55 g Cu c. 4.83 g Sn Text Transcription: HNO_3

The cell potential of this electrochemical cell depends on the pH of the solution in the anode half-cell. \(Pt(s) | H_{2}( g, 1 atm) H^{+}(aq, ? M) || Cu^{2+}(aq, 1.0 M) | Cu(s)\) What is the pH of the solution if \(E_{cell}\) is 355 mV? Text Transcription: Pt(s) | H_2( g, 1 atm) H^+(aq, ? M) || Cu^2+(aq, 1.0 M) | Cu(s) E_cell

The cell potential of this electrochemical cell depends on the pH of the solution in the anode half-cell. \(Pt(s) | H_{2}( g, 1.0 atm) H^{+}(aq, 1.0 M) || Au^{3+}(aq, ? M) Au(s)\) What is the concentration of \(Au^{3+}\) in the solution if \(E_{cell}\) is 1.22 V? Text Transcription: Pt(s) | H_2( g, 1 atm) H^+(aq, 1.0M) || Au^3+ (aq, ? M) | Au(s) Au^3+ E_cell

A battery relies on the oxidation of magnesium and the reduction of \(Cu^{2+}\). The initial concentrations of \(Mg^{2+}\) and \(Cu^{2+}\) are \(1.0 \times 10^{-4} M\) 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? Text Transcription: Cu^2+ Mg^2+ 1.0 x 10^-4 M

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 \(1.0 \times 10^{-3} M\). 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 \(1.00 \times 10^{2} g\) of silver at a charging current of 10.0 amps? Text Transcription: Ag/Ag^+ Ag^+ 1.0 x 10^-3 M 1.00 x 10^2 g

The \(K_{sp}\) of CuI is \(1.1 \times 10^{-12}. Find \(E_{cell}\) for the cell: \(Cu(s) | CuI(s) | I^{-}(aq)(1.0 M) || Cu^{+}(aq)(1.0 M) | Cu(s)\) Text Transcription: K_sp 1.1 x 10^12 E_cell Cu(s) | CuI(s) | I^-(aq)(1.00M) || Cu^+(aq)(1.0M) | Cu(s)

The \(K_{sp}\) of \(Zn(OH)_{2}\) is \(1.8 \times 10^{-14}\). Find \(E_{cell}\) for the half-reaction: \(Zn(OH)_{2}(s) + 2 e^{-} ? Zn(s) + 2 OH^{-}(aq)\) Text Transcription: K_sp Zn(OH)_2 1.8 x 10^-14 E_cell Zn(OH)_{2}(s) + 2 e^{-} ? Zn(s) + 2 OH^{-}(aq)

Calculate \(?G°_{rxn}\) and K for each reaction. a. The disproportionation of \(Mn^{2+}(aq)\) to Mn(s) and \(MnO_{2}(s)\) in acid solution at 25 °C. b. The disproportionation of \(MnO_{2}(s)\) to \(Mn^{2+}(aq)\) and \(MnO_{4}^{-}(aq)\) in acid solution at 25 °C. Text Transcription: ?G°_rxn Mn^2+(aq) MnO_2(s) MnO_4^-(aq)

Calculate \(?G°_{rxn}\) and K for each reaction. a. The reaction of \(Cr^{2+}(aq)\) with \(Cr_{2}O_{7}^{2-}\)(aq) in acid solution to form \(Cr^{3+}(aq)\). b. The reaction of \(Cr^{3+}(aq)\) and Cr(s) to form \(Cr^{2+}(aq)\). [The electrode potential of \(Cr^{2+}(aq)\) to Cr(s) is -0.91 V.] Text Transcription: ?G°_rxn Cr^2+(aq) Cr_2O_7^2- Cr^3+(aq) Cr^2+(aq)

A metal forms the fluoride \(MF_{3}\). 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. Text Transcription: MF_3

A sample of impure tin of mass 0.535 g is dissolved in strong acid to give a solution of \(Sn^{2+}\). The solution is then titrated with a 0.0448 M solution of \(NO^{3-}\), which is reduced to NO( g). The equivalence point is reached upon the addition of 0.0344 L of the \(NO^{3-}\) solution. Find the percent by mass of tin in the original sample, assuming that it contains no other reducing agents. Text Transcription: Sn^2+ NO^3-

A 0.0251-L sample of a solution of \(Cu^{+}\) requires 0.0322 L of 0.129 M \(KMnO_{4}\) solution to reach the equivalence point. The products of the reaction are \(Cu^{2+}\) and \(Mn^{2+}\). What is the concentration of the \(Cu^{2+}\) solution? Text Transcription: Cu^+ Mn^2+ Cu^2+

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 \(H_{2}\)( g), which bubbles out of solution. What is the pH of the solution after 73 minutes? Text Transcription: H^+ H_2

An \(MnO_{2}(s)/Mn^{2+}(aq)\) electrode in which the pH is 10.24 is prepared. Find the \([Mn^{2+}]\) necessary to lower the potential of the half-cell to 0.00 V (at 25 °C). Text Transcription: MnO_2(s)/Mn^2+(aq) [Mn^2+]

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 \(1.2 \times 10^{3} kWh\) of electricity per month. Text Transcription: 1.2 x 10^3 kWh

A voltaic cell designed to measure [\(Cu^{2+}\)] is constructed of a standard hydrogen electrode and a copper metal electrode in the \(Cu^{2+}\) 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? Text Transcription: Cu^2+

The surface area of an object to be gold plated is \(49.8 cm^{2}\), and the density of gold is \(19.3 g/cm^{3}\). 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 \(1.00 \times 10^{-3} cm\) thick on the object. Text Transcription: 49.8 cm^2 19.3 g/cm^3 1.00 x 10^-3 cm

To electrodeposit all the Cu and Cd from a solution of \(CuSO_{4}\) and \(CdSO_{4}\) 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 \(CuSO_{4}\) was present in the original mixture? Text Transcription: CuSO_4 CdSO_4 1 F = 1 mol e^-

Sodium oxalate, \(Na_{2}C_{2}O_{4}\), in solution is oxidized to \(CO_{2}( g)\) by \(MnO_{4}^{-}\), which is reduced to \(Mn^{2+}\). A 50.1-mL volume of a solution of \(MnO_{4}^{-}\) is required to titrate a 0.339-g sample of sodium oxalate. This solution of \(MnO_{4}^{-}\) 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 \(UO^{2+} \rightarrow UO_{2}^{2+}\). Calculate the percentage of uranium in the sample. Text Transcription: Na_2C_2O_4 CO_2(g) MnO_4^- UO^2+ UO_2^2+

The cell \(Pt(s) | Cu^{+}(1 M), Cu^{2+}(1 M) || Cu^{+}(1 M) Cu(s)\) has E° = 0.364 V. The cell \(Cu(s) | Cu^{2+}(1 M) || Cu^{+}(1 M) | Cu(s)\) has E° = 0.182 V. Write the cell reaction for each cell and explain the differences in E°. Calculate ?G° for each cell reaction to help explain these differences. Text Transcription: Pt(s) | Cu^+(1 M), Cu^2+(1 M) || Cu^+(1 M) Cu(s) Cu(s) | Cu^2+(1 M) || Cu^+(1 M) | Cu(s)

Which oxidizing agent will oxidize \(Br^{-}\) but not \(Cl^{-}\)? a. \(K_{2}Cr_{2}O_{7}\) (in acid) b. \(KMnO_{4}\) (in acid) c. \(HNO_{3}\) Text Transcription: K_2Cr_2O_7 KMnO_4 HNO_3

A redox reaction employed in an electrochemical cell has a negative \(?G°_{rxn}\). Which statement is true? a. \(E°_{cell}\) is positive; K < 1 b. \(E°_{cell}\) is positive; K > 1 c. \(E°_{cell}\) is negative; K > 1 d. \(E°_{cell}\) is negative; K < 1 Text Transcription: ?G°_rxn E°_cell

A redox reaction has an equilibrium constant of K = 0.055. What is true of \(?G°_{rxn}\) and \(E°_{cell}\) for this reaction? Text Transcription: ?G°_rxn E°_cell

Balance the redox reactions by following the steps in the text. Rotate through the group, having each group member do the next step in the process and explain that step to the rest of the group. a. \(I_{2}(s) + Fe(s) \rightarrow FeI_{2}(s)\) b. \(Cl_{2}( g) + H_{2}O_{2}(aq) \rightarrow Cl^{-}(aq) + O_{2}( g) (acidic)\) c. \(Hg^{2+}(aq) + H_{2}( g) \rightarrow Hg(l) + H_{2}O(l) (basic)\) d. \(CH_{3}OH(l) + O_{2}( g) \rightarrow CO_{2}( g) + H_{2}O(l) (acidic)\) Text Transcription: I_2(s) + Fe(s) FeI_2(s) Cl_2( g) + H_2O_2(aq) Cl^-(aq) + O_2( g) (acidic) Hg^{2+}(aq) + H_{2}( g) \rightarrow Hg(l) + H_{2}O(l) (basic) CH_3OH(l) + O_2( g) CO_2( g) + H_2O(l) (acidic)

Calculate ?G° and K for each reaction the group created in Question 143. For one of the reactions, explain how the sign or magnitude of each quantity (\(E°_{cell}\), ?G°, and K) is consistent with the fact that the reaction is spontaneous in the direction written. Text Transcription: E°_cell

Design a device that uses an electrochemical cell to determine the amount of \(Cu_{2+}\) in a sample of water. Describe, in detail, the construction and the theory of operation of your device. If you are able to measure voltage with one-millivolt accuracy, what will the uncertainty in your measured concentration be? Text Transcription: Cu^2+

In this chapter, you have seen that the voltage of an electrochemical cell is sensitive to the concentrations of the reactants and products in the cell. As a result, electrochemical cells can be used to measure the concentrations of certain species in solution. For example, the voltage of an electrochemical cell based on the reaction \(H_{2}( g) + Cu^{2+}(aq) \rightarrow 2 H^{+}(aq) + Cu(s)\) is sensitive to both the \(Cu^{2+}\) concentration and the \(H^{+}\) concentration in solution. If the H+ concentration is held constant, then the voltage only depends on the \(Cu^{2+}\) concentration, and we can use the cell to measure the \(Cu^{2+}\) concentration in an unknown solution. The tabulated data shows the measured voltage in the hydrogen/copper electrochemical cell just discussed for several different \(Cu^{2+}\) concentrations. Examine the data and answer the questions that follow. a. Construct a graph of the measured voltage versus the copper concentration. Is the graph linear? b. Determine how you might manipulate the data to produce a linear graph. (Hint: See the Nernst equation.) c. Reconstruct a graph of the data using the method to produce a linear graph from part b. Determine the slope and y-intercept of the best-fitting line to the points in your graph. Could you have predicted the slope and intercept from the Nernst equation? d. The voltage of two unknown solutions are measured and recorded. Use the slope and intercept from part c to determine the \(Cu^{2+}\) concentrations of the unknown solutions. Text Transcription: H_2( g) + Cu^2+(aq) 2 H^+(aq) + Cu(s) Cu^2+ H^+

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

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

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

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

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?

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?

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

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.

What is the definition of the standard cell potential (E°cell)? 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?

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

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

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

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

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

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 20.1?

What is a concentration electrochemical cell?

What are the anode and cathode reactions in a common drycell battery? In an alkaline battery?

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

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

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

Explain how a fuel-cell breathalyzer works.

List some applications of electrolysis.

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

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

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?

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

What is overvoltage in an electrochemical cell? Why is it important?

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

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

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

How can the corrosion of iron be prevented?

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

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

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

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

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

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

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

What products are obtained in the electrolysis of molten NaI?

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

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

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.

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.

https://docs.google.com/document/d/1bSPpfyZ6KESGjimisZOMRSrX3jQZ_vfIsirapNtBqkU/edit

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?

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

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 g Al b. 4.85 g Cu c. 2.42 g Ag

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?

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?

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 °C?

When a suspected drunk driver blows 188 mL of his breath through the fuel-cell breathalyzer described in Section 20.7, 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 °C, what percent (by volume) of the driver’s breath is ethanol?

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 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?

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.)

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 was deposited?

An electrochemical cell has a positive standard cell potential but a negative cell potential. Which statement is true for the cell? a. K 7 1; Q 7 K b. K 6 1; Q 7 K c. K 7 1; Q 6 K d. K 6 1; Q 6 K

Have each group member select a half-reaction from Table 20.1. Each member should calculate the standard cell potential of an electrochemical cell formed between each member’s halfreaction and the half-reaction of each of the other group members. For each pair of half-reactions, write the overall balanced chemical reaction that will be spontaneous.

Using a library or the Internet, research a fuel cell that uses methanol for fuel. What is the reaction at the anode? What is the reaction at the cathode? What is the overall reaction? What is the standard cell potential? How many kWh can it generate from 1 L (0.792 kg) of methanol?