5.4 Electrochemistry Calculations & Applications (A Level Only)

Table 2.1 lists electrode potentials for the Cr₂O₇^2– (aq) / Cr³⁺ (aq) and Br₂ (l) / Br^– (aq) half-cells.

Table 2.1

Electrode reaction	Eθ / V
½Br2 (l) + e−Br− (aq)	+ 1.09
Cr2O72– (aq) + 14H+ (aq) + 6e–  2Cr3+ (aq) + 7H2O (l)	+ 1.36

Deduce the full equation for the Cr₂O₇^2- (aq) / Cr³⁺ (aq) and Br₂ (l) / Br^– (aq) cell.

Using Table 2.1, calculate the E^θ_cell for the electrochemical cell outlined in part (a).

The electrochemical equation for standard free energy change is given.

ΔG^θ = -nFE^θ

The Faraday constant, F = 9.65 × 10⁴ C mol^–1.

Use your answer to parts (a) and (b) to determine whether the reaction of the electrochemical cell is feasible.

An electrochemical cell has a free energy change of -14.475 kJ mol^-1.

Use the information in Table 2.1 to determine the reactions taking place at each electrode of the electrochemical cell. 

Table 2.1

• reaction at anode ........................................................................................
   
• reaction at cathode ........................................................................................

Table 3.1 lists electrode potentials for some electrode reactions.

Table 3.1 

Explain how Table 3.1 could be adapted to show an electrochemical series.

Use Table 3.1 to identify the halide ion that is the weakest reducing agent.

Use Table 3.1 to justify why sulfate ions should not be capable of oxidising iodide ions.

Suggest why the two cobalt(III) complex ions in Table 3.1 have different electrode potentials.

Use Table 3.1 to explain why [Co(H₂O)₆]³⁺(aq) will undergo a redox reaction with [Fe(H₂O)₆]²⁺(aq).

Give an equation for this reaction.

explanation ................................................................................

equation ................................................................................

An electrochemical cell is set up as shown in Fig. 2.1.

Fig. 2.1

Use the electrode potential list in Table 2.1 to calculate the value of E^θ_cell under standard conditions, stating which electrode is the negative one.

Table 2.1

Electrode reaction	Eθ / V
Ag+ + e– Ag	+0.80
Fe2+ + 2e–  Fe	-0.44
Fe3+ + e– Fe2+	+0.77
SO42– + 4H+ + 2e–  SO2 + 2H2O	+0.17

 E^θ_cell  = ..............................................

negative electrode = ..............................................

How would the actual E_cell of the above cell compare to the E^θ_cell under standard conditions?

Explain your answer.

How would the E_cell of the cell in Fig. 2.1 change, if at all, if a few cm³ of concentrated Na₂SO₄ (aq) were added to the following?

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[1]

[2]

Write an equation to show the reaction taking place in the electrochemical cell in Fig. 2.1.

The standard electrode potentials for seven different redox systems are shown in Table 3.1.

Table 3.1

redox system	equation	Eθ / V
1	2H+ (aq) + 2e– H2 (g)	0.00
2	Fe3+ (aq) + e– Fe2+ (aq)	+0.77
3	Cu2+ (aq) + 2e– Cu (s)	+0.34
4	Cl2 (aq) + 2e– 2Cl– (aq)	+1.36
5	O2 (g) + 4H+ (aq) + 4e– 2H2O (l)	+1.23
6	Al3+ (aq) + 3e– Al (s)	-1.66
7	I2 (aq) + 2e– 2I– (aq)	+0.54

An electrochemical cell can be made based on redox systems 2 and 3.

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[1]

[3]

[2]

Select from Table 3.1,

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Calculate the electrode potential at 298 K of redox system 3 if the concentration of Cu²⁺ (aq) ions was 0.0002 mol dm^-3.

Table 4.1 below contains some standard electrode potential data which you will need to answer the following questions.

Table 4.1

electrode reaction	Eθ / V
Cu2+ (aq) + 2e−  Cu (s)	+0.34
Ni2+ (aq) + 2e−  Ni (s)	-0.25
Fe3+ (aq) + e−  Fe2+ (aq)	+0.77
Sn2+ (aq) + 2e−  Sn (s)	-0.14
Fe2+ (aq) + 2e−  Fe (s)	-0.45

Deduce the species from Table 1 that is the weakest oxidising agent. Explain your choice.

A cell is made by connecting two half-cells with a salt bridge.  

A student produced a cell by using nickel in a solution of nickel chloride solution and another consisted of copper in a solution of copper sulfate solution.

Calculate the standard cell potential of this cell using the values given in Table 4.1 in part (a).

Calculate the standard Gibbs free energy change, and give the units, for the electrochemical cell in part (b) and state whether the reaction is feasible.

Show your working.

[Faraday constant, F = 9.65 × 10⁴ C mol^–1]

Two half-cells, involving species in Table 4.1, are connected together to give a cell with a standard cell potential = +0.31 V.

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The standard electrode potentials in Table 5.1 can be used to predict redox reactions.

Table 5.1

redox system	equation	Eθ / V
1	Ag+ (aq) + e– Ag (s)	+0.80
2	Cr3+ (aq) + 3e– Cr (s)	-0.74
3	Mg2+ (aq) + 2e– Mg (s)	-2.37

Using the information in Table 5.1, write equations for the reactions that are feasible.

A student sets up a standard cell, using half-cells based on redox systems 2 and 3 at 298 K.

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The student diluted the solution in redox system 3 with distilled water.

Predict what would happen to the cell potential.
Explain your reasoning.

The student set up the cell between redox systems 2 and 3 again. This time, they did not dilute the solution of redox system 3 but they used a solution of 0.01 mol dm^-3 of Cr³⁺ ions for redox system 2. The temperature remained at 298 K.

Use the Nernst equation below to predict what the new electrode potential, E, would be for redox system 2.

E = E^θ + 

where z = the number of electrons transferred in the reaction.