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Context for question 6

A scientist constructs a galvanic cell as shown in the diagram. As the cell operates, the Zn(s) electrode increases in mass and the Al(s) electrode decreases in mass.

Galvanic cell with Zn(s) and Al(s) electrodes in 1.0 M solutions, a salt bridge and a 0.90 V voltmeter

A data table with the standard reduction potentials for the substances follows the diagram.

Half-Reaction	E $^{\circ}$ (V)
$Z n^{2 +} (aq) + 2 e^{-} \to Zn (s)$	−0.76
$A l^{3 +} (aq) + 3 e^{-} \to Al (s)$	−1.66

6(A)Oxidation Numbers, Redox Equations and Redox TitrationsEasy1 mark

Write the half-reaction for the oxidation that occurs at the anode.

6(B)Oxidation Numbers, Redox Equations and Redox TitrationsEasy1 mark

Write the balanced net ionic equation for the overall reaction that occurs in the galvanic cell.

6(C)Electrochemical Cells, Nernst Equation, Batteries and Fuel CellsMedium1 mark

Initially, each electrode has a mass of 50.0 g. The cell is allowed to run for a period of time and is then stopped. Which electrode's mass changed the most? Justify your answer with a calculation.

6(D)Electrochemical Cells, Nernst Equation, Batteries and Fuel CellsEasy1 mark

Reduction Half-Reaction	E $^{\circ}$ (V)
$A u^{3 +} (aq) + 3 e^{-} \to Au (s)$	+1.50
$Z n^{2 +} (aq) + 2 e^{-} \to Zn (s)$	−0.76
$M n^{2 +} (aq) + 2 e^{-} \to Mn (s)$	−1.19
$A l^{3 +} (aq) + 3 e^{-} \to Al (s)$	−1.66
$B e^{2 +} (aq) + 2 e^{-} \to Be (s)$	−1.85

The standard Zn/Al cell has a value of

E_{ce l l}^{\circ}

= 0.90 V. The scientist needs a galvanic cell that produces a greater voltage. The scientist has access to the chemical systems in the table. If the scientist uses the Zn half-cell and one of the other options from the table, what is the MAXIMUM voltage that could be generated at standard conditions?