Syllabus Edition

First teaching 2020

Last exams 2024

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Electrochemistry Calculations & Applications (A Level Only) (CIE A Level Chemistry)

Exam Questions

2 hours13 questions
1a
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1 mark

Table 1.1 shows the standard electrode potentials for the Mg2+ / Mg and I2 / I half-cells.

Table 1.1

Electrode reaction Eθ / V
I2 + 2e ⇌ 2I + 0.54
Mg2+ + 2e ⇌ Mg – 2.38

 

Calculate the standard cell potential for the electrochemical cell using the Mg2+ / Mg and I2 / I half-cells.

 
Eθ = .......... V
1b1 mark

Write the equation for the reaction that occurs in the electrochemical cell described in part (a).

1c1 mark

Explain why the I2 / I half cell is the positive pole for the electrochemical cell described in part (a).

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2a2 marks

The electrochemical series was established by measuring the potential of various electrodes against the standard hydrogen electrode.

The electrochemical series can be used to identify oxidising and reducing agents.

 

In terms of electrons, explain what it meant by the terms oxidising agent and reducing agent.

2b4 marks

Table 2.1 shows a section of the electrochemical series.

 
Table 2.1
 
Electrode reaction Eθ / V
Ba2+ + 2e ⇌ Ba – 2.90
Cu2+ + 2e ⇌ Cu  + 0.34
F2 + 2e ⇌ 2F + 2.87
MnO4 + 8H+ + 5e ⇌ Mn2+ + 4H2O + 1.52
O2 + H2O + 2e ⇌ HO2 + OH – 0.08
Zn2+ + 2e ⇌ Zn – 0.76
 

Using Table 2.1, identify the strongest and weakest reducing agents.

2c2 marks

Using Table 2.1, identify the strongest and weakest oxidising agents.

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3a1 mark

An electrochemical cell is set up to measure the electrode potential, E, for the Cu2+ / Cu half-cell with a standard hydrogen electrode.

The standard electrode potential for the Cu2+ / Cu half-cell is

Cu2+ + e ⇌ Cu    Eθ = + 0.34 V

State two factors that can affect the equilibrium position of the electrode reaction.

3b1 mark

A 0.750 mol dm-3 solution of copper(II) chloride to make the Cu2+ / Cu half-cell.

State the effects that this solution of copper(II) chloride has on the equilibrium position and electrode potential, E.

Equilibrium position ...............................................................................

Electrode potential ................................................................................

3c1 mark

Identify the oxidised and reduced species for the Cu2+ / Cu half-cell.

 

Oxidised species ....................

 

Reduced species ....................

3d
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1 mark

The Nernst equation is:

 
E = Eθbegin mathsize 14px style fraction numerator 0.059 over denominator z end fraction end style log10 fraction numerator open square brackets oxidised space species close square brackets over denominator open square brackets reduced space species close square brackets end fraction
 
i)
To calculate the electrode potential at 298 K of the Cu2+ / Cu half-cell, the Nernst equation can be revised to:
 
E = Eθfraction numerator 0.059 over denominator z end fraction log10 [oxidised species]
 
Using your answer to part (c), suggest why the concentration of the reduced species does not feature in the revised Nernst equation calculation for the Cu2+ / Cu half-cell.
 
[1] 
 
ii)
Use the information in parts (a) and (b) and the revised Nernst equation to calculate the electrode potential at 298K of the Cu2+ / Cu half-cell.
 
E = .......... V
 
[2]

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4a3 marks

The standard free energy change can be calculated using the standard electrode potentials of the half-cells in an electrochemical cell.

ΔGθ = – n x Eθcell x

Complete Table 4.1 to show what each term in the equation represents.

 
Table 4.1 
 
Term  
ΔGθ  Standard free energy change
n  
Eθcell   
F  
 
4b2 marks

The following electrochemical cell is set up.

 
Fe3+ (aq) + Cr (s) → Fe (s) + Cr3+ (aq)
 

Use Table 4.2 to determine the two half-cell electrode reactions and their Eθ values. 

 
Table 4.2
 
Electrode reaction Eθ / V
Cr2+ + 2e ⇌ Cr – 0.91
Cr3+ + 3e ⇌ Cr – 0.74
Cr3+ + e ⇌ Cr2+ – 0.41
Fe2+ + 2e ⇌ Fe – 0.44 
Fe3+ + 3e ⇌ Fe  – 0.04
Fe3+ + e ⇌ Fe2+ + 0.77
 
4c
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4 marks

Use the information in part (a) and your answer to part (b), to calculate the standard Gibbs free energy change, in kJ mol-1, for the electrochemical cell. 

F = 96 500 C mol-1 
 

ΔGθ = .................... kJ mol-1

4d1 mark

Use your answer to part (c) to suggest whether the reaction is feasible. Explain your answer.

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5a1 mark

Faraday’s constant can be used to calculate the mass of a substance deposited at an electrode or the volume of gas liberated at an electrode during electrolysis.

A current of 1.65 A flows through molten lead(II) chloride for 10 minutes.

 

Write the equation for the electrolysis of lead(II) chloride, PbCl2

5b
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6 marks
i)
Write the half-equation for the reduction of the lead(II) ion, Pb2+.
 
[1]
 
ii)
The equation to calculate the charge transferred is QIt. The Faraday constant, = 96 500 C mol-1.
 
Use the information in part (a) to calculate:
 
The number of coulombs required to deposit one mole of lead at the cathode =  .................... C
 
The charge transferred during the electrolysis of the molten lead(II) bromide = .................... C
 
The mass of lead deposited at the cathode = .................... (g)
 
[5]
5c
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3 marks
i)
Write the half-equation for the oxidation of the chloride ion, Cl to form chlorine, Cl2.
 
[1]
 
ii)
The equation to calculate the charge transferred is QIt. The Faraday constant, = 96 500 C mol-1.
 
Use the information in part (a) and your answers to part (b) to calculate:
 
The number of coulombs required to form one mole of chlorine at the anode =  .................... C
 
The volume of chlorine released at the anode = .................... (dm3)
 
[2]

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1a1 mark

This question is about electrolysis cells.

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

1b4 marks

If the charge on the electron, the Ar and the charge of copper ions are known, the Avogardo constant can be determined experimentally by passing a known current for a known time through a copper electrolysis cell and weighing the mass of copper deposited onto the cathode.

Draw a diagram of apparatus which is suitable for carrying out this practical.
Label the following:

  • power supply with the + and - terminals identified
  • anode
  • cathode
  • ammeter

State the composition of the electrolyte.

1c
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5 marks

The following are the results obtained from one such experiment.

current passed through the cell = 0.375 A
time current was passed through cell = 40.0 min
initial mass of copper cathode = 49.637 g
final mass of copper cathode = 49.936 g

Use the experimental results to calculate a value of the Avogardo constant, L, to 3 significant figures.

[electronic charge, e = -1.60 x 10-19 C]

1d
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5 marks

Table 2.1 shows a list of standard electrode potentials at 298 K.

Table 2.1

Electrode reaction Eθ / V
Ag+ + e–  rightwards harpoon over leftwards harpoon   Ag +0.80
Br2 + 2e–  rightwards harpoon over leftwards harpoon  2Br +1.07
F2 + 2e–  rightwards harpoon over leftwards harpoon  2F +2.87
Fe2+ + 2e–  rightwards harpoon over leftwards harpoon  Fe –0.44
2H+ + 2e rightwards harpoon over leftwards harpoon H2 0.00
Mg2+ + 2e–  rightwards harpoon over leftwards harpoon  Mg –2.38
O2 + 2H2O + 4e– rightwards harpoon over leftwards harpoon 4OH +0.40
SO42– + 4H+ + 2e–  rightwards harpoon over leftwards harpoon  SO2 + 2H2O +0.17



Use the information in Table 2.1 to determine the substances formed at the anode and cathode when the following substances are electrolysed.

compound product at anode product at cathode
AgF    
FeSO4    
MgBr2    

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2a1 mark

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

5-4-2a-m-iron-and-silver-cell


Fig. 2.1

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


Table 2.1

Electrode reaction Eθ / V
Ag+ + e rightwards harpoon over leftwards harpoon Ag +0.80
Fe2+ + 2e rightwards harpoon over leftwards harpoon Fe -0.44
Fe3+ + e rightwards harpoon over leftwards harpoonFe2+ +0.77
SO42– + 4H+ + 2e rightwards harpoon over leftwards harpoon SO2 + 2H2 +0.17

 Eθcell  = ..............................................

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

2b2 marks

How would the actual Ecell of the above cell compare to the Eθcell under standard conditions?


Explain your answer.

2c4 marks

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

i)
The beaker containing Fe3+ (aq) and Fe2+ (aq).

[1]

ii)
The beaker containing Ag2SO4 (aq).

[1]

iii)
Explain any changes in Ecell you have stated in (i) and (ii).

[2]

2d1 mark

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

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3a
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2 marks

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 rightwards harpoon over leftwards harpoon H2 (g)   0.00
2 Fe3+ (aq) + e rightwards harpoon over leftwards harpoon Fe2+ (aq) +0.77
3 Cu2+ (aq) + 2erightwards harpoon over leftwards harpoon Cu (s) +0.34
4 Cl2 (aq) + 2erightwards harpoon over leftwards harpoon 2Cl (aq) +1.36
5 O2 (g) + 4H+ (aq) + 4erightwards harpoon over leftwards harpoon 2H2O (l) +1.23
6 Al3+ (aq) + 3erightwards harpoon over leftwards harpoon Al (s) -1.66
7 I2 (aq) + 2erightwards harpoon over leftwards harpoon 2I (aq) +0.54

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

i)
Write the overall cell reaction.

[1]

ii)
Calculate the voltage of this cell.

[1]

3b5 marks
i)
Using redox systems 56 and 7 only in Table 3.1, write the overall equations for three reactions that might be feasible.

[3]

ii)
Give two reasons why these reactions might not take place, even if they are feasible.

[2]

3c2 marks

Select from Table 3.1,

i)
an oxidising agent that oxidises Fe2+ (aq) to Fe3+ (aq),

[1]

ii)
a species that reduces Fe3+ (aq) to Fe2+ but does not reduce Cu2+ (aq) to Cu (s).

[1]

3d
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3 marks

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

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4a2 marks

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 rightwards harpoon over leftwards harpoon Cu (s) +0.34
Ni2+ (aq) + 2e rightwards harpoon over leftwards harpoon Ni (s) -0.25
Fe3+ (aq) + e rightwards harpoon over leftwards harpoon Fe2+ (aq) +0.77
Sn2+ (aq) + 2e rightwards harpoon over leftwards harpoon Sn (s) -0.14
Fe2+ (aq) + 2e rightwards harpoon over leftwards harpoon Fe (s) -0.45
 

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

4b
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1 mark

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).

4c
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3 marks

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]

4d2 marks

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

i)
Determine which two half equations are connected using the data from Table 4.1.

[1]

ii)
Suggest the half-equation for the reaction that occurs at the cathode.

[1]

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5a3 marks

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 rightwards harpoon over leftwards harpoon Ag (s) +0.80
2 Cr3+ (aq) + 3e rightwards harpoon over leftwards harpoon Cr (s) -0.74
3 Mg2+ (aq) + 2erightwards harpoon over leftwards harpoon Mg (s) -2.37


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

5b
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2 marks

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

i)
Calculate the standard cell potential.

[1]

ii)
State the sign of the electrode in redox system 2 of the cell.

[1]

5c3 marks

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

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

5d
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2 marks

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 = Eθfraction numerator 0.059 over denominator z end fraction log subscript 10 fraction numerator open square brackets oxidised space species close square brackets over denominator open square brackets reduced space species close square brackets end fraction

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

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1a4 marks

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.

i)
State the conditions of temperature and pressure under which standard electrode potentials are measured.
 
[1]
 
ii)
Fig. 1.1 shows the diagram drawn by the student.
 
q2aii-9cho-al-3-june-2019-qp-edexcel-a-level-chem
 
Fig. 1.1
  
Complete Table 1.1 to identify three mistakes in this diagram and the modifications that should be made to correct them.
 
Table 1.1
 
Mistake in diagram Modification needed to correct mistake
   
   
   
 
[3]

1b
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3 marks

The standard electrode potential, Eθ, 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 = Eθbegin mathsize 14px style fraction numerator R T over denominator 96500 end fraction end styleln[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

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2a1 mark

Table 2.1 lists electrode potentials for the Cr2O72– (aq) / Cr3+ (aq) and Br2 (l) / Br (aq) half-cells.

Table 2.1

Electrode reaction  Eθ / V
½Br2 (l) + erightwards harpoon over leftwards harpoonBr (aq)  + 1.09
Cr2O72– (aq) + 14H+ (aq) + 6e begin mathsize 14px style rightwards harpoon over leftwards harpoon end style 2Cr3+ (aq) + 7H2O (l) + 1.36

 

Deduce the full equation for the Cr2O72- (aq) / Cr3+ (aq) and Br2 (l) / Br (aq) cell.

2b
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1 mark

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

2c
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1 mark

The electrochemical equation for standard free energy change is given.

ΔGθ = -nFEθ

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.

2d
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2 marks

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

Electrode reaction Eθ / V
Ag+ (aq) + e- Ag (s) +0.80
Li+ (aq) + e- Li (s) -3.04
ClO2 (aq) + e- ClO2- (aq) +0.95
H2O (l) + e- ⇌ ½H2 (g) + OH- (aq) -0.83
Fe3+ (aq) + e- ⇌ Fe2+ (aq) +0.77

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

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3a1 mark

Table 3.1 lists electrode potentials for some electrode reactions.

Table 3.1 

Electrode reaction  Eθ / V
Br2 + 2e ⇌ 2Br + 1.07
Cl2 + 2e ⇌ 2Cl   + 1.36
[Co(H2O)6]3++ e ⇌  [Co(H2O)6]2+ + 1.81
[Co(NH3)6]3+ + e ⇌ [Co(NH3)6]2+ 
+ 0.11
Cu2+ + 2e ⇌ Cu + 0.34
Fe2+ + 2e ⇌ Fe  – 0.44
Fe3+ + eFe2+ + 0.77
2H+ + 2e ⇌ H2 0.00
I2 + 2e ⇌ 2I + 0.54
NO3 + 2H+ + e ⇌ NO2 + H2O + 0.81
SO42– + 4H+ + 2e ⇌ SO2 + 2H2O + 0.17
VO2+ + 2H+ + e ⇌ VO2+ + H2O +1.00

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

3b1 mark

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

3c1 mark

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

3d
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3 marks
i)
Use Table 3.1 to identify an acid that will oxidise copper. Explain your answer.
 
[1]
 
ii)
Suggest a possible equation for the reaction.
 
[1]
 
iii)
Calculate the Eθcell for the same overall reaction.
 
[1]
3e1 mark

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

3f2 marks

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

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