A2 LEVEL CHEMISTRY

ELECTRODE POTENTIALS AND FUEL CELLS 

TEST

Answer all questions
Max 43 marks

	
Name		…………………………………………………………..

Mark		……../43	……....%	Grade ………

1.	Concentrated nitric acid is a powerful oxidising agent. Concentrated nitric acid oxidises
sulphur to sulphuric acid. Nitrogen dioxide and another product are also formed.
(i)	Suggest a balanced equation for this reaction.
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[1]
(ii)	Deduce the change in oxidation number of nitrogen in this reaction.
from ............... to ...............
[1]
[Total 2 marks]

2.	NO2 reacts with oxygen and water to form nitric acid, HNO3. In the atmosphere, this contributes to acid rain. Construct a balanced equation for this formation of nitric acid and use oxidation numbers to show that this is a redox reaction.
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[Total 2 marks]

3.	The standard electrode potential of the Cl2/ Cl– half-cell may be measured using the following apparatus.

(a)	Suggest suitable labels for A, B, C and D.
A ....................................................................................................................
B ....................................................................................................................
C ....................................................................................................................
D ....................................................................................................................
[2]

(b)	The half cell reactions involved are shown below.
		Cl2 + e–  Cl–	Eο = +1.36 V
		H+ + e–  H2	Eο = 0.00V
(i)	Use an arrow to show the direction of flow of electrons in the diagram of the apparatus. Explain your answer.
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[2]
(ii)	The values of Eο are measured under standard conditions. What are the standard conditions?
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[2]
(c)	The half cell reaction for ClO3–/Cl2 is shown below.
		ClO3– + 6H+ + 5e–  Cl2 + 3H2O	Eο = +1.47 V
	What does this tell you about the oxidising ability of ClO3– compared with Cl2?
	Explain your answer.
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[2]
[Total 8 marks]

4.	Chlorine gas may be prepared in the laboratory by reacting hydrochloric acid with potassium manganate(VII). The following standard electrode potentials relate to this reaction.
		Cl2 + e–		Cl–	Eο = +1.36 V
		MnO4– + 8H+ + 5e–		Mn2+ + 4H2O	Eο = +1.52 V
(a)	Define the term standard electrode potential.
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[3]
(b)	Determine the standard cell potential for a cell constructed from these two redox systems.

[1]
[Total 4 marks]

5.	Chromium is an important metallic element. Its compounds have a number of different oxidation states.
The following equations relate to half-cells involving iron and chromium ions.
		Fe3+ + e–  Fe2+	Eο = +0.77 V
		Cr2O72– + 14H+ + 6e–  2Cr3+ + 7H2O	Eο = +1.33 V
A cell was set up by combining these two half-cells.
(i)	Derive a balanced equation for the reaction that would occur when the cell is in use.
Explain your reasoning in terms of oxidation and reduction.
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[3]
(ii)	Determine the emf of the cell under standard conditions.

emf = ...................................................... V
[1]
[Total 6 marks]

6.	Some standard electrode potentials are shown below.
			Eο/V
	Ag+ + e–  Ag	+ 0.80
	Cl2 + e–  Cl–	+ 1.36
	Cu2+ + 2e–  Cu	+ 0.34
	Fe3+ + e–  Fe2+	+ 0.77
	I2 + e–  I–	+ 0.54
(a)	The diagram below shows an incomplete cell consisting of Cu/Cu2+ and Ag/Ag+ half-cells.

(i)	Complete and label the diagram to show how the cell potential of this cell could be measured.
[2]
(ii)	On the diagram, show the direction of electron flow in the circuit if a current was allowed.
[1]
(iii)	Calculate the standard cell potential.

standard cell potential = ……………………V
[1]
 (iv)	Write the overall cell reaction.
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[1]
 
(b)	Chlorine will oxidise Fe2+ to Fe3+ but iodine will not. Explain why, using the electrode potential data.
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[2]
[Total 7 marks]

7.	Chlorine gas may be prepared in the laboratory by reacting hydrochloric acid with potassium manganate(VII). The following standard electrode potentials relate to this reaction.
		Cl2 + e–		Cl–	Eο = +1.36 V
		MnO4– + 8H+ + 5e–		Mn2+ + 4H2O	Eο = +1.52 V
(a)	Use the half-equations above to:
(i)	construct an ionic equation for the reaction between hydrochloric acid and potassium manganate(VII);
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(ii)	determine the oxidation numbers of chlorine and manganese before and after the reaction has taken place;
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[2]
(iii)	state what is oxidised and what is reduced in this reaction.
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[2]

(b)	If potassium manganate(VII) and very dilute hydrochloric acid are mixed, there is no visible reaction. Suggest why there is no visible reaction in this case.
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[1]
[Total 7 marks]
 
8.	Use the standard electrode potentials in the table below to answer the questions that follow.
I	Fe2+(aq) + 2e–  Fe(s)	Eο = –0.44 V	
II	V3+(aq) + e–  V2+(aq)	Eο = –0.26 V	
III	2H+(aq) + 2e–  H2(g)	Eο = 0.00 V	
IV	O2(g) + 4H+(aq) + 4e–  2H2O(l)	Eο = +0.40 V	
	An electrochemical cell was set up based on systems I and II.
(i)	Write half-equations to show what has been oxidised and what has been reduced in this cell.
oxidation:

reduction:

[2]
(ii)	Determine the cell potential of this cell.

Ecell = ......................................................... V
[1]
[Total 3 marks]

9.	Use the standard electrode potentials in the table below to answer the questions that follow.
I	Fe2+(aq) + 2e–  Fe(s)	Eο = –0.44 V	
II	V3+(aq) + e–  V2+(aq)	Eο = –0.26 V	
III	2H+(aq) + 2e–  H2(g)	Eο = 0.00 V	
IV	O2(g) + 4H+(aq) + 4e–  2H2O(l)	Eο = +0.40 V	
	An electrochemical fuel cell was set up based on systems III and IV.
 
(i)	Construct an equation for the spontaneous cell reaction. Show your working.

[2]
(ii)	Fuels cells based on systems such as III and IV are increasingly being used to generate energy.
	Discuss two advantages and two disadvantages of using fuels cells for energy rather than using fossil fuels.
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[4]
[Total 6 marks]