Principles of Electrochemistry (A Level Only) (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.

A student decided to determine the value of the Faraday constant by an electrolysis experiment. Fig. 2.1 is an incomplete diagram showing the apparatus that was used.

Fig. 2.1

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List the measurements the student would need to make in order to use the results to calculate a value for the Faraday constant.

Using an equation, state the relationship between the Faraday constant, F, the Avogadro constant, L, and the charge on the electron, e.

The value the student obtained was: 1 faraday = 9.65 × 10⁴ coulombs

Use this value and your equation in (c) to calculate the Avogadro constant (take the charge on the electron to be 1.60 × 10^–19 coulombs).

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.

The electrolytic purification of copper can be carried out in an apparatus similar to the one shown in Fig. 5.1.

Fig. 5.1

The impure copper anode contains small quantities of metallic nickel, zinc and silver, together with inert oxides and carbon resulting from the initial reduction of the copper ore with coke. The copper goes into solution at the anode, but the silver remains as the metal and falls to the bottom as part of the anode ‘sludge’. The zinc also dissolves.

Table 5.1 shows a list of standard electrode potentials at 298 K.

Table 5.1

Electrode reaction	Eθ / V
Ag+ + e– Ag	+0.80
Cu2+ + 2e–  Cu	+0.34
Fe2+ + 2e–  Fe	-0.44
Ni2+ + 2e– Ni	-0.25
SO42– + 4H+ + 2e–  SO2 + 2H2O	+0.17
Zn2+ + 2e– Zn	-0.76

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As the electrolysis proceeds, the blue colour of the electrolyte slowly fades. 

Suggest why the blue colour fades.

Most of the current passed through the cell is used to dissolve the copper at the anode and precipitate pure copper onto the cathode. However, a small proportion of it is ‘wasted’ in dissolving the impurities at the anode which then remain in solution. When a current of 20.0 A was passed through the cell for 10.0 hours, it was found that 225 g of pure copper was deposited on the cathode.

[Faraday constant, F = 9.65 x 10⁴ C mol^-1]

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Nickel often occurs in ores along with iron. After the initial reduction of the ore with coke, a nickel-iron alloy is formed.

Use data from Table 5.1 to explain why nickel can be purified by a similar electrolysis technique to that used for copper, using an impure nickel anode, a pure nickel cathode, and nickel sulfate as the electrolyte.

Explain what would happen to the iron during this process.

Fig. 1.1 shows a lithium–iodine electrochemical cell. These cells are often used in pacemakers as they are reliable and have a life span in the region of 10 years.

 Fig. 1.1

It consists of a lithium electrode and an inert electrode immersed in body fluids that are separated by a nickel mesh and collect charge from the anode. It has a high internal resistance which means that only a low current can be drawn.

Explain why the lithium-iodine electrochemical cell is a dry cell.

Write the overall equation for the reaction taking place at the electrodes of the lithium-iodine electrochemical cell when a current flows. 

• overall equation ................................................................................

Table 1.1 lists electrode potentials for some electrode reactions.

Table 1.1

Electrode reaction	Eθ / V
I2 + 2e–2I–	+ 0.54
Li+ + e–Li	– 3.04
Ni2+ + 2e–Ni	– 0.25

Calculate the E^θ_cell for the lithium-iodine electrochemical cell.

E^θ_cell = .................... V

A current of 1.5 × 10^–5 A is drawn from this cell.

Calculate the number of days for 0.08 g of the lithium electrode to be used up. Assume the current remains constant throughout this period. Show your working.

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

time = .............................. days

A student set up an electrochemical cell using a concentrated sodium chloride solution using a current of 6 A. 

State the half-equations occurring at the electrodes during the electrolysis of the concentrated aqueous solution of sodium chloride.

   Cathode ..................................................................................................
 

   Anode ....................................................................................................

Calculate the time, in minutes, to produce 2.00 dm³ of gas at the anode at standard temperature and pressure.

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

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

time = ..................... minutes

The student changed the electrolyte to a very dilute sodium chloride solution.

State what change would occur at the anode and give the half equation for the process.

In a different electrolysis experiment, copper(II) sulfate solution was electrolysed using graphite electrodes.

Table 2.1 lists electrode potentials for some electrode reactions.

Table 2.1

Electrode reaction	Eθ / V
Cu+ + e–Cu	+ 0.52
Cu2+ + 2e–Cu	+ 0.34
Cu2+ + e–Cu+	+ 0.15
2H2O + 2e–H2 + 4OH–	– 0.83
O2 + 2H2O + 4e–4OH–	+ 0.40
½O2 + 2H+ + 2e–H2O	+ 1.23
SO42– + 4H+ + 2e–SO2 + 2H2O	+ 0.17
S2O82– + 2e–2SO42–	+ 2.01

Explain how the products at the anode and cathode are produced. 

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