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Lesson: The Nernst Equation

Worksheet • 10 Questions

Q1:

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 H g S ( ) + 2 e H g ( ) + S ( ) s l a q – 2 – A g ( ) + e A g ( ) + – a q s
Standard electrode potential, 𝐸 ⦵ (V) – 0 . 7 0 + 0 . 7 9 9 6
  • A 1.43 V
  • B 1.46 V
  • C 1.37 V
  • D 1.53 V
  • E 1.50 V

Q2:

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

Half-equation A l ( ) + 3 e A l ( ) 3 + – a q s C u ( ) + 2 e C u ( ) 2 + – a q s
Standard electrode potential, 𝐸 ⦵ (V) – 1 . 6 6 2 + 0 . 3 4 0

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 C d ( ) + 2 e C d ( ) 2 + – a q s C d S ( ) + 2 e C d ( ) + S ( ) s s a q – 2 –
Standard electrode potential, 𝐸 ⦵ (V) – 0 . 4 0 3 0 – 1 . 1 7
  • A 3 . 1 × 1 0   
  • B 5 . 8 × 1 0   
  • C 5 . 2 × 1 0   
  • D 2 . 8 × 1 0   
  • E 1 . 4 × 1 0   

Q4:

Calculate to 2 significant figures the equilibrium constant at 2 5 ∘ C for the reaction:

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

Half-equation 2 H ( ) + 2 e H ( ) + – 2 a q g 2 H O ( ) + 2 e H + 2 O H ( ) 2 – 2 – l a q
Standard electrode potential, 𝐸 ⦵ (V) 0.000 − 0 . 8 2 7 7
  • A 1 . 0 × 1 0   
  • B 1 . 0 × 1 0   
  • C 1 . 0 × 1 0   
  • D 1 . 0 × 1 0   
  • E 1 . 0 × 1 0   

Q5:

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 H g ( ) + 2 e H g ( ) 2 + – a q l [ H g B r ] ( ) + 2 e H g ( ) + 4 B r ( ) 4 2 – – – a q l a q
Standard electrode potential, 𝐸 ⦵ (V) + 0 . 8 5 1 + 0 . 2 1
  • A 4 . 7 × 1 0  
  • B 1 . 6 × 1 0   
  • C 6 . 8 × 1 0  
  • D 2 . 1 × 1 0   
  • E 7 . 4 × 1 0  

Q6:

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 A l ( ) + 3 e A l ( ) 3 + – a q s N i ( ) + 2 e N i ( ) 2 + – a q s
Standard Electrode Potential, 𝐸 ⦵ (V) – 1 . 6 6 2 – 0 . 2 5 7
  • A 1.42 V
  • B 1.41 V
  • C 1.43 V
  • D 1.40 V
  • E 1.39 V

Q7:

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 B r ( ) + 2 e 2 B r ( ) 2 – – a q a q A l ( ) + 3 e A l ( ) 3 + – a q s
Standard Electrode Potential, 𝐸 ⦵ (V) + 1 . 0 8 7 3 − 1 . 6 6 2

Q8:

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 C u ( ) + 2 e C u ( ) 2 + – a q s A g ( ) + e A g ( ) + – a q s
Standard electrode potential, 𝐸 ⦵ ( V )

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

  • A 3 . 4 × 1 0  
  • B 6 . 8 × 1 0  
  • C 5 . 9 × 1 0 
  • D 4 . 0 × 1 0  
  • E 3 . 4 × 1 0  

If a particular dead battery is found to have [ C u ] 2 + = 0.11 M, what is the concentration of silver ions?

  • A 5 . 6 × 1 0   M
  • B 3 . 8 × 1 0   M
  • C 3 . 2 × 1 0    M
  • D 0.22 M
  • E 0.11 M

Q9:

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

Half-equation C o ( ) + 2 e C o ( ) 2 + – a q s F e ( ) + 2 e F e ( ) 2 + – a q s
Standard electrode potential, 𝐸 ⦵ (V) − 0 . 2 8 − 0 . 4 4 7

Q10:

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 A g ( ) + e A g ( ) + – a q s A g C l ( ) + e A g ( ) + C l ( ) s s a q – –
Standard electrode potential, 𝐸 ⦵ (V) + 0 . 7 9 9 6 + 0 . 2 2 2 3 3
  • A 1 . 7 × 1 0   
  • B 8 . 8 × 1 0   
  • C 3 . 6 × 1 0   
  • D 1 . 6 × 1 0   
  • E 3 . 8 × 1 0   
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