Standard Electrode Potentials E⦵, Standard Cell Potentials E⦵cell & the Nernst Equation (CIE A Level Chemistry)

This question is about electrochemical cells.

Fig. 1.1 shows the apparatus used to measure the standard electrode potential, E^θ, of a cell composed of a Cu(II)/Cu half-cell and an Fe(II)/Fe half-cell. 

Fig. 1.1

Finish the diagram by adding components to show the complete circuit. Label the components you add.

In the spaces below, identify or describe what the four letters A-D in Fig. 1.1 represent.

The standard electrode potentials of the two half-cells are as follows.

Fe2+ + 2e–		Fe	Eθ = -0.44 V
Cu2+ + 2e–		Cu	Eθ = +0.34 V

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E^θ _cell of this cell is +0.78 V.

State whether this reaction is feasible. Explain your answer.

This question is about electrochemical cells.

The diagram shown in Fig. 3.1 represents the standard hydrogen electrode.

Fig 3.1

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A student set up an electrochemical cell consisting of copper and zinc.

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Fig. 3.2

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A student set up another electrochemical cell consisting of copper and silver as shown in the following cell representation where the half cell with the greatest negative standard electrode potential is written on the left: 

Cu / Cu²⁺ ; Ag⁺ / Ag

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A diagram of a cell is shown below in Fig. 3.3.

Fig. 3.3

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The apparatus shown was used to measure the standard electrode potential for the reduction of Cr₂O₇^2– ions to Cr³⁺ ions in acid solution:

Cr₂O₇^2– (aq) + 14H⁺ (aq) + 6e^–→ 2Cr³⁺ (aq) + 7H₂O (l)

Which material should be used for each electrode?

Suggest a suitable solution that could be used as solution 1.

Solution 2 contains 14.71g of K₂Cr₂O₇ .

What mass of Cr₂(SO₄)₃⋅18H₂O should also be used?

[M_r values: K₂Cr₂O₇ = 294.2 Cr₂(SO₄)₃⋅18H₂O = 716.3]

Solution 2 is best acidified with H₂SO₄ instead of HCl or HBr.

Suggest why.

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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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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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.

Explain why Fig. 3.1 does not represent the standard hydrogen electrode.

Fig. 3.1

The standard electrode potential for Zn²⁺ (aq) + 2e^– → Zn (s) is –0.76 V.

State the meaning of the minus sign in the value of –0.76 V.

Zinc coating on metals serves as physical protection which prevents rust from affecting the underlying metal surface. This is achieved by electroplating as shown in Fig. 3.2.

Fig. 3.2

Calculate the length of time, in hours, required to deposit 1.0 g of zinc on the item to be electroplated. Assume the current is a constant 0.1 A throughout this period.  

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

State your answer to 2 significant figures and show your working.

time = .................... hours

This question is about the Ag⁺(aq) / Ag(s) half-cell.

A student was asked to plan an experiment to measure the standard electrode potential of the Ag⁺ (aq) / Ag(s) half-cell.

The standard electrode potential, E^θ, of the Ag⁺ (aq) / Ag (s) half-cell is +0.80 V.

The effect of changing the concentration of the ions on the value of the electrode potential, E, in this half-cell is calculated using the equation

E = E^θ + ln[Ag⁺ (aq)]

where T is the temperature in kelvin and R is the gas constant, 8.31 J K^–1 mol^–1.

The electrode potential of an Ag⁺ (aq) / Ag (s) half-cell was measured at 20 °C and found to be +0.72 V.

Calculate the concentration of silver ions, in mol dm^-3, in this half-cell. Show your working.

[Ag⁺] = .................. mol dm^-3

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 ................................................................................