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Problem 47d
Consider the voltaic cell:
d. Indicate the direction of anion and cation flow in the salt bridge
Problem 49
Use line notation to represent each electrochemical cell in Problem 43.
Problem 51
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) | Sn2+(aq) || NO(g) | NO3–(aq), H+(aq) | Pt(s)
- Is the question formulating correctly? If yes, return the question without changes. If not, please fix it and return the output as a JSON of the form: {'question': 'question text'}. Here is the question: 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) | Mn2+(aq) || ClO2-(aq) | ClO2(g) | Pt(s)
Problem 52
Problem 53
Determine whether or not each redox reaction occurs spontaneously in the forward direction.
a. Ni(s) + Zn2+(aq) → Ni2+(aq) + Zn(s)
b. Ni(s) + Pb2+(aq) → Ni2+(aq) + Pb(s)
c. Al(s) + 3 Ag+(aq) → Al3+(aq) + 3 Ag(s)
d. Pb(s) + Mn2+(aq) → Pb2+(aq) + Mn(s)
Problem 54
Determine whether or not each redox reaction occurs spontaneously in the forward direction.
a. Ca2+(aq) + Zn(s) → Ca(s) + Zn2+(aq)
b. 2 Ag+(aq) + Ni(s) → 2 Ag(s) + Ni2+(aq)
c. Fe(s) + Mn2+(aq) → Fe2+(aq) + Mn(s)
d. 2 Al(s) + 3 Pb2+(aq) → 2 Al3+(aq) + 3 Pb(s)
Problem 55
Which metal could you use to reduce Mn2+ ions but not Mg2+ ions?
Problem 56
Which metal can be oxidized with an Sn2+ solution but not with an Fe2+ solution?
Problem 57
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 59
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
Problem 60
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
Problem 61a
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. a. 2 Cu(s) + Mn2+(aq) → 2 Cu+(aq) + Mn(s)
Problem 61b,c
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. b. MnO2(aq) + 4 H+(aq) + Zn(s) → Mn2+(aq) + 2H2O(l) + Zn2+(aq) c. Cl2(g) + 2 F–(aq) → F2(g) + 2 Cl–(aq)
Problem 62a,b
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. a. O2(g) + 2 H2O(l) + 4 Ag(s) → 4 OH–(aq) + 4 Ag+(aq) b. Br2(l) + 2 I–(aq) → 2 Br–(aq) + I2(s)
Problem 62c
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. c. PbO2(s) + 4 H+(aq) + Sn(s) → Pb2+(aq) + 2 H2O(l) + Sn2+(aq)
Problem 63
Which metal cation is the best oxidizing agent? a. Pb2+ b. Cr3+ c. Fe2+ d. Sn2+
Problem 65b,c
Use tabulated electrode potentials to calculate ∆G°rxn for each reaction at 25 °C. b. Br2(l) + 2 Cl–(aq) → 2 Br–(aq) + Cl2(g) c. MnO2(s) + 4 H+(aq) + Cu(s) → Mn2+(aq) + 2 H2O(l) + Cu2+(aq)
Problem 66
Use tabulated electrode potentials to calculate ∆G°rxn for each reaction at 25 °C. a. 2 Fe3+(aq) + 3 Sn(s) → 2 Fe(s) + 3 Sn2+(aq) b. O2(g) + 2 H2O(l) + 2 Cu(s) → 4 OH–(aq) + 2 Cu2+(aq) c. Br2(l) + 2 I–(aq) → 2 Br–(aq) + I2(s)
Problem 67
Calculate the equilibrium constant for each of the reactions in Problem 65.
- Calculate the equilibrium constant for the reaction between Ni2+(aq) and Cd(s) at 25 °C.
Problem 69
Problem 70
Calculate the equilibrium constant for the reaction between Fe2+(aq) and Zn(s) (at 25 °C).
- Calculate ΔG°_rxn and E°_cell for a redox reaction with n = 2 that has an equilibrium constant of K = 25 (at 25 °C).
Problem 71
- Calculate ΔG°rxn and E°cell for a redox reaction with n = 3 that has an equilibrium constant of K = 0.050 (at 25 °C).
Problem 72
Problem 73
A voltaic cell employs the following redox reaction: Sn2+(aq) + Mn(s) → Sn(s) + Mn2+(aq) Calculate the cell potential at 25 °C under each set of conditions. c. [Sn2+] = 2.00 M; [Mn2+] = 0.0100 M
- A voltaic cell employs the redox reaction: 2 Fe3+(aq) + 3 Mg(s) → 2 Fe(s) + 3 Mg2+(aq). Calculate the cell potential at 25 °C under each set of conditions. a. standard conditions. b. [Fe3+] = 1.0 × 10^-3 M; [Mg2+] = 2.50 M. c. [Fe3+] = 2.00 M; [Mg2+] = 1.5 × 10^-3 M.
Problem 74
Problem 75
An electrochemical cell is based on these two half-reactions:
Ox: Pb(s) → Pb2+(aq, 0.10 M) + 2 e–
Red: MnO4–(aq, 1.50 M) + 4 H+(aq, 2.0 M) + 3 e– → MnO2(s) + 2 H2O(l)
Calculate the cell potential at 25 °C.
Problem 76
An electrochemical cell is based on these two half-reactions:
Ox: Sn(s) → Sn2+(aq, 2.00 M) + 2 e–
Red: ClO2(g, 0.100 atm) + e– → ClO2–(aq, 2.00 M)
Calculate the cell potential at 25 °C.
Problem 77a
A voltaic cell consists of a Zn/Zn2+ half-cell and a Ni/Ni2+ half-cell at 25 °C. The initial concentrations of Ni2+ and Zn2+ are 1.50 M and 0.100 M, respectively. a. What is the initial cell potential?
Problem 77b
A voltaic cell consists of a Zn/Zn2+ half-cell and a Ni/Ni2+ half-cell at 25 °C. The initial concentrations of Ni2+ and Zn2+ are 1.50 M and 0.100 M, respectively. b. What is the cell potential when the concentration of Ni2+ has fallen to 0.500 M?
Problem 77c
A voltaic cell consists of a Zn/Zn2+ half-cell and a Ni/Ni2+ half-cell at 25 °C. The initial concentrations of Ni2+ and Zn2+ are 1.50 M and 0.100 M, respectively. c. What are the concentrations of Ni2+ and Zn2+ when the cell potential falls to 0.45 V?