In this worksheet, we will practice using the Nernst equation to calculate reduction potentials under nonstandard concentrations.

**Q1: **

A battery is dead when it has no cell potential. Consider a battery with the overall reaction: The standard electrode potentials for the half-cells in this battery are given in the table.

Half-equation | ||
---|---|---|

Standard electrode potential, |

To 2 significant figures, what is the value of when this battery is dead at 298.15 K?

- A
- B
- C
- D
- E

If a particular dead battery is found to have = 0.11 M, what is the concentration of silver ions?

- A M
- B 0.22 M
- C 0.11 M
- D M
- E M

**Q2: **

Calculate to 2 significant figures the equilibrium constant at for the reaction:

Note that each standard electrode potential is expressed per mole of the half-reaction shown in the table.

Half-equation | ||
---|---|---|

Standard electrode potential, (V) | 0.000 |

- A
- B
- C
- D
- E

**Q3: **

Using the standard electrode potentials shown in the table, calculate to 2 significant figures the equilibrium constant at 373 K for the reaction:

Half-equation | ||
---|---|---|

Standard electrode potential, (V) |

- A
- B
- C
- D
- E

**Q4: **

Using the standard electrode potentials shown in the table, calculate to 2 significant figures the equilibrium constant at 373 K for the reaction:

Half-equation | ||
---|---|---|

Standard electrode potential, (V) |

- A
- B
- C
- D
- E

**Q5: **

Using the standard electrode potentials shown in the table, calculate to 2 significant figures the equilibrium constant at 298.15 K for the reaction:

Half-equation | ||
---|---|---|

Standard electrode potential, (V) |

- A
- B
- C
- D
- E

**Q6: **

In the half-cells of an electrochemical cell, 1.00 M aqueous bromide ions are oxidized to 0.110 M bromine and 0.0230 M aluminum ions are reduced to aluminum metal. Using the standard electrode potentials shown in the table, calculate to 3 decimal places the cell potential for the cell at 298.15 K. Note that standard electrode potentials are measured using 1.00 M solutions of the reacting ions.

Half-Equation | ||
---|---|---|

Standard Electrode Potential, (V) |

**Q7: **

The half-cells of a galvanic cell consist of an aluminum electrode in a 0.0150 M aluminum nitrate solution and a nickel electrode in a 0.250 M nickel(II) nitrate solution. Using the standard electrode potentials shown in the table, calculate to 2 decimal places the cell potential for the galvanic cell at 298.15 K. Note that standard electrode potentials are measured using 1.00 M solutions of the reacting ions.

Half-Equation | ||
---|---|---|

Standard Electrode Potential, (V) |

- A 1.39 V
- B 1.43 V
- C 1.40 V
- D 1.42 V
- E 1.41 V

**Q8: **

Using the standard electrode potentials shown in the table, calculate to 2 decimal places the cell potential at 298.15 K for the cell with the overall reaction:

Half-equation | ||
---|---|---|

Standard electrode potential, (V) |

- A 1.50 V
- B 1.37 V
- C 1.53 V
- D 1.43 V
- E 1.46 V

**Q9: **

Using the standard electrode potential data in the table, calculate the standard cell potential for the following reaction at 298 K.

Half-equation | ||
---|---|---|

Standard electrode potential, (V) |

**Q10: **

Calculate to 3 significant figures the cell potential for the following reaction at 298 K?

Half-equation | ||
---|---|---|

Standard electrode potential, (V) |