Balance each redox reaction occurring in acidic aqueous solution. a. K(s) + Cr³⁺(aq) → Cr(s) + K⁺(aq) b. Al(s) + Fe²⁺(aq) → Al³⁺(aq) + Fe(s)

Balance each redox reaction occurring in acidic aqueous solution. c. BrO₃^–(aq) + N₂H₄(g) → Br^–(aq) + N₂(g)

Balance each redox reaction occurring in acidic aqueous solution. a. Zn(s) + Sn²⁺(aq) → Zn²⁺(aq) + Sn(s)

Balance each redox reaction occurring in acidic aqueous solution. b. Mg(s) + Cr³⁺(aq) → Mg²⁺(aq) + Cr(s)

Balance each redox reaction occurring in acidic aqueous solution. c. MnO₄^–(aq) + Al(s) → Mn²⁺(aq) + Al³⁺(aq)

Balance each redox reaction occurring in acidic aqueous solution. a. PbO₂(s) + I^–(aq) → Pb²⁺(aq) + I₂(s) b. SO₃^2–(aq) + MnO₄^–(aq) → SO₄^2–(aq) + Mn²⁺(aq) c. S₂O₃^2–(aq) + Cl₂(g) → SO₄^2–(aq) + Cl^–(aq)

Balance each redox reaction occurring in acidic aqueous solution. a. I^–(aq) + NO₂^–(aq) → I₂(s) + NO(g) b. ClO₄^–(aq) + Cl^–(aq) → ClO₃^–(aq) + Cl₂(g)

Balance each redox reaction occurring in acidic aqueous solution. c. NO₃^–(aq) + Sn²⁺(aq) → Sn⁴⁺(aq) + NO(g)

Balance each redox reaction occurring in basic aqueous solution. a. H₂O₂(aq) + ClO₂(aq) → ClO₂^–(aq) + O₂(g)

Balance each redox reaction occurring in basic aqueous solution. b. Al(s) + MnO₄^–(aq) → MnO₂(s) + Al(OH)₄^–(aq)

Balance each redox reaction occurring in basic aqueous solution. c. Cl₂(g) → Cl^–(aq) + ClO^–(aq)

Balance each redox reaction occurring in basic aqueous solution. a. MnO₄^–(aq) + Br^–(aq) → MnO₂(s) + BrO₃^–(aq)

Balance each redox reaction occurring in basic aqueous solution. b. Ag(s) + CN^–(aq) + O₂(g) → Ag(CN)₂^–(aq)

Balance each redox reaction occurring in basic aqueous solution. c. NO₂^–(aq) + Al(s) → NH₃(g) + AlO₂^–(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) → 2 Ag(s) + Pb²⁺(aq)

b. 2 ClO₂(g) + 2 I^–(aq) → 2 ClO₂^–(aq) + I₂(s)

c. O₂(g) + 4 H⁺(aq) + 2 Zn(s) → 2 H₂O(l) + 2 Zn²⁺(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. Ni²⁺(aq) + Mg(s) → Ni(s) + Mg²⁺(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.

b. 2 H⁺(aq) + Fe(s) → H₂(g) + Fe²⁺(aq)

c. 2 NO₃^–(aq) + 8 H⁺(aq) + 3 Cu(s) → 2 NO(g) + 4 H₂O(l) + 3 Cu²⁺(aq)

Calculate the standard cell potential for each of the electro- chemical cells in Problem 43.

Consider the voltaic cell:

d. Indicate the direction of anion and cation flow in the salt bridge

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

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)

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)

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)

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?

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