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 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)

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)

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)

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)

Which metal cation is the best oxidizing agent? a. Pb²⁺ b. Cr³⁺ c. Fe²⁺ d. Sn²⁺

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)

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)

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

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

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

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.

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.

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?

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?

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?

A voltaic cell consists of a Pb/Pb²⁺ half-cell and a Cu/Cu²⁺ half-cell at 25°C. The initial concentrations of Pb²⁺ and Cu²⁺ are 0.0500 M and 1.50 M, respectively. a. What is the initial cell potential?

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

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

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 1.0×10^–3 M. Label the anode and the cathode and indicate the half-reaction occuring at each electrode. Also indicate the direction of electron flow.

Consider the concentration cell: b. Indicate the direction of electron flow.

Consider the concentration cell:

c. Indicate what happens to the concentration of Pb²⁺ in each half-cell.