Redox & Electrochemistry (OCR A Level Chemistry)

Exam Questions

3 hours33 questions
1a2 marks

Vanadium is a transition metal that can form a variety of ions with multiple oxidation states possible. This question is about the reactions of vanadium ions.

Deduce the oxidation state of vanadium in each of the following ions:

  • V2+
  • VO2+
  • VO2+
  • V3+
1b
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2 marks

The standard electrode potentials for two half equations are shown below:

Half equations Eθ / V
VO2+ (aq) + 2H+ (aq) + e- rightwards harpoon over leftwards harpoonH2O (l) + VO2+ (aq) +1.00
S4O62- (aq) + 2e- rightwards harpoon over leftwards harpoon2S2O32- (aq) +0.47

i)
Identify the species acting as the oxidising agent in the reaction between these two half equations.

[1]

ii)
Write an overall equation for the reaction that would occur between these species.

[1]

1c
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1 mark

Calculate the standard cell potential between half cells containing VO2+ (aq)/VO2+ (aq) and containing S4O62- (aq)/ S2O32- (aq), using information from the table below.

 
Half equations Eθ / V
VO2+ (aq) + 2H+ (aq) + e- rightwards harpoon over leftwards harpoonH2O (l) + VO2+ (aq) +1.00
S4O62- (aq) + 2e- rightwards harpoon over leftwards harpoon2S2O32- (aq) +0.47

 

1d2 marks

Some other reactions of vanadium ions are shown below: 

Half equations Eθ / V
Zn2+ (aq) + 2e- rightwards harpoon over leftwards harpoon Zn(s) -0.76
V3+ (aq) + e- rightwards harpoon over leftwards harpoon V2+ (aq) -0.26
Sn2+ (aq) + 2e- rightwards harpoon over leftwards harpoon Sn (s) -0.14
VO2+ (aq) + 2H+ (aq) + e- rightwards harpoon over leftwards harpoon H2O (l) + V3+ (aq) +0.34
S4O62- (aq) + 2e- rightwards harpoon over leftwards harpoon2S2O32- (aq) +0.47
VO2+ (aq) + 2H+ (aq) + e- rightwards harpoon over leftwards harpoonH2O (l) + VO2+ (aq) +1.00

i)
Which vanadium ion would be produced from a reaction of VO2+ (aq) with tin, Sn?

[1]

ii)
Which vanadium ion would be produced from a reaction of VO2+ (aq) with zinc, Zn?

[1]

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2a4 marks

This question is about the techniques involved in some common redox titrations.

Two common redox titration methods involve titrating Fe2+ with MnO4- and titrating I2 with S2O32-. The reactions that occur are shown below:

5Fe2+ + MnO4 + 8H+ → 5Fe3+ + Mn2+ + 4H2O

2S2O32– (aq) + I2 (aq) → S4O62– (aq) + 2I (aq)

i)
Explain why no indicator is used for the Fe2+/ MnO4- titration
[2]
ii)
Explain why a starch indicator is needed for the I2/ S2O32- titration
[2]
2b2 marks

Even at very low concentrations, the solution of MnO4- used in the burette for the Fe2+/MnO4- titration is still opaque rather than translucent.

Describe and explain one adjustment to the traditional titration technique that is needed to compensate for this.

2c
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3 marks

A titration was carried out to determine the unknown concentration of a solution of 25.00 cm3 FeCl2.

An average titre of 22.65 cm3 of 0.0500 mol dm-3 of potassium permanganate was used.

Calculate the concentration of FeCl2 to 3 significant figures.

2d2 marks

To get a reliable titre value the end-point needs to be obtained accurately.

For this titration, a white tile is used under the conical flask and the sides of the flask are rinsed down with distilled water.

Explain how each of these helps to accurately identify the end-point of the titration.

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3a2 marks

This question is about setting up and using electrochemical cells.

A cell is set up with a mixture of acidified Cr2O72- and Cr3+ in one half-cell and methanoic acid, HCOOH, with CO2 bubbled through it in the other.

 
Half equations Eθ / V
CO2 (g) + 2H+ (aq) + 2e-rightwards arrowHCOOH (aq) -0.22
Cr2O72- (aq) + 14H+ (aq) +  6e- rightwards harpoon over leftwards harpoon2Cr3+ (aq) + 7H2O (l) +1.33

Write an overall equation for the reaction that could occur.

3b3 marks

What should be used as the electrode for each of the half cells described in part a) and why?

3c
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1 mark

Calculate the standard cell potential for the reaction using the values from the table.

 
Half equations Eθ / V
CO2 (g) + 2H+ (aq) + 2e-rightwards arrowHCOOH (aq) -0.22
Cr2O72- (aq) + 14H+ (aq) +  6e- rightwards harpoon over leftwards harpoon2Cr3+ (aq) + 7H2O (l) +1.33

3d2 marks

Describe what would be seen at each of the electrodes if the reaction was allowed to occur.

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4a1 mark

This question is about storage cells.

Describe the difference between primary and secondary cells.

4b
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1 mark

One example of a secondary cell is the lead-acid battery found in cars.

Half equations Standard electrode potential / V
PbSO4 (s)  +  2e-  rightwards harpoon over leftwards harpoon   Pb (s) +  SO42- (aq)   -0.36 
PbO2 (s) +  4H(aq) +  SO42- (aq) +  2e- rightwards harpoon over leftwards harpoon  PbSO4 (s)  + 2H2O (l)  +1.70

                                               

Calculate the standard cell potential of a lead-acid cel

4c2 marks

Identify the substances used to form the positive and negative electrodes in the lead-acid cell.

4d2 marks

Lead acid batteries have a finite lifespan and become less effective over time. 

Outline two difficulties with their disposal.

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5a
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1 mark

This question is about using standard electrode potentials.

A chromium half cell containing Cr/Cr3+ is connected to a copper half cell containing Cu/Cu2+.

The standard electrode potential for copper is shown as: 

Cu2+ (aq) + 2e-  rightwards harpoon over leftwards harpoon  Cu (s)            Eθ = +0.34 V

The standard cell potential of this combination was found to be + 1.08 V, where the copper half cell would undergo reduction.

Calculate the standard electrode potential of the Cr/Cr3+ half cell.

5b4 marks

Describe what the standard cell and cell conditions are that the standard electrode potential measurements are measured against.

5c2 marks

Write out two half equations to show what would occur if the chromium half cell was attached to the standard hydrogen half cell.

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1a2 marks

Some standard electrode potential data are shown below.

Half-equation

  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)

−0.44

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

1b
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1 mark

A cell is constructed from nickel in a solution of nickel(II) chloride and copper in a solution of copper(II) sulfate. Calculate the standard cell potential of this cell using the values given in the table.

1c
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3 marks

Two half-cells, involving species in the table, are connected together to give a cell with a standard cell potential of +0.30 V.

i)

Determine which two half equations produce this cell potential using the data from the table and write the overall equation for the reaction.

ii)

Suggest the half-equation for the reaction that occurs at the positive electrode (cathode).

1d
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4 marks

An electrochemical cell is set up using the following two half-cells.

Cl2 (g) + 2e⇌ 2Cl- (aq)        E = +1.36 V

Sn2+ (aq) + 2e⇌ Sn (s)        E = -0.14 V

i)

Predict whether the formation or dissociation of tin(II) chloride, SnCl2, is spontaneous.

ii)

Suggest one reason why the spontaneous reaction may not occur.

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2a6 marks

The diagram shown below represents the standard hydrogen electrode.

standard-hydrogen-electrode

i)

Name the substance used as the electrode in the diagram.

ii)
Suggest why this substance is used as an electrode.

iii)

Give the standard conditions used in a standard hydrogen electrode (SHE).

2b6 marks

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

zinc-half-cell

i)

Complete the diagram to show the components and reagents, including their concentrations and label any apparatus required to complete the electrochemical cell.

ii)

Use the IUPAC convention to give the half equations occurring at each electrode

2c3 marks

A student set up another electrochemical cell consisting of the Cu2+ / Cu (Eθ = + 0.34 V) and Ag+ / Ag (Eθ = + 0.80 V) half cells.

 
i)
Write a half-equation for the reaction that occurs at the positive electrode.
 
ii)
Write a half-equation for the reaction that occurs at the negative electrode
 
iii)
Use the half-equations to deduce an overall equation for the cell. Include all state symbols.
2d2 marks

An electrochemical cell is shown below:

magnesium-silver-electrochemical-cell

i)

Explain how the salt bridge provides an electrical connection between the two solutions.

ii)

Suggest why potassium chloride would not be suitable for use in the salt bridge of this cell.

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3a
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1 mark

Lithium ion cells are a type of rechargeable cell. In the cell lithium ions flow, enter and exit the solid electrodes.

The half-equations for the reaction at the electrodes can be represented as follows.

   Positive electrode:   Li+ + CoO2 + e Li+[CoO2]   Eθ = +0.36 V

   Negative electrode:   Li+ + e → Li    Eθ = -3.04 V

Calculate the standard cell potential of the cell. 

3b2 marks

During discharge the lithium ions move from the negative electrode to the positive electrode; whilst recharging a lithium ion cell, the lithium ions move from the positive electrode to the negative electrode.

Give the half equations for the electrodes during recharging.

3c2 marks

The aircraft Pathfinder is an example of a solar rechargeable aircraft (SRA). Solar cells are powered during the day by capturing energy from sunlight and converting it into electricity. 

Explain why rechargeable cells are often attached to the solar cells.

3d3 marks

A rechargeable nickel–cadmium cell is an alternative to a lithium ion cell. The half-equations for this cell are given below:

   Cd(OH)2 (s) + 2e → Cd (s) + 2OH (aq)   Eθ = -0.88 V

   NiO(OH) (s) + H2O (I) + e → Ni(OH)2 (s) + OH (aq)   Eθ = +0.52 V

i)

Deduce the oxidation state of the cadmium in this cell after recharging is complete.

ii)

Write an equation for the overall reaction that occurs when the cell is recharged.

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4a3 marks

Fuel cells are used to generate an electric current and do not need to be electrically recharged. The Gemini and Apollo moon probes use hydrogen-oxygen fuel cells. The product of this reaction can be used to supplement the drinking water for astronauts.

Deduce the half equations for the reactions at each electrode in a hydrogen oxygen fuel cell, and then the overall equation for the hydrogen-oxygen fuel cell, to show the product used to supplement drinking water.

4b3 marks

A fuel cell is an electrochemical device which converts chemical energy into electrical energy. A continuous supply of fuel is supplied to one electrode and an oxidant to the other. 

i)

Explain how an electric current is generated in the fuel cell.

ii)

Suggest why a fuel cell does not need to be recharged.

4c1 mark

The electrodes used in hydrogen fuel cells are often made of a porous mixture of carbon-supported platinum or a porous ceramic material coated in platinum.

State why the electrodes must be porous.

4d3 marks

The General Motors Electrovan was built in 1966. It was the first vehicle powered by a hydrogen fuel cell and could travel at up to 70 mph for 30 seconds.

i)

Suggest a main advantage of using hydrogen in a fuel cell rather than an internal combustion engine. 

ii)

State why using a fuel cell to power a vehicle has an environmental advantage over the internal combustion

iii)

Hydrogen fuel cells are commonly considered to be carbon neutral by the general public although this is not technically correct. Suggest why hydrogen fuel cells cannot be classed as carbon neutral.

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5a
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8 marks

A 500 cm3 standard solution was made using 13.0 g of hydrated iron(II) sulfate, FeSO4.xH2O in acidic solution

A 25.0 cm3 sample of this solution was titrated against 0.0200 mol dm-3 potassium manganate(VII) solution. 

26.75 cm3 of the potassium manganate(VII) solution was required.

i)

Write the ionic equation for the reaction of manganate and iron(II) ions.

ii)

State and explain the colour change that occurs at the end-point.

iii)

Calculate the value of x in FeSO4.xH2O.

5b2 marks

The percentage of iron in a steel wire can be determined from a redox titration using acidified potassium manganate(VII) solution.

The steel wool is placed in an excess of sulfuric acid to dissolved the iron.

Explain why the iron wire is dissolved in an excess of sulfuric acid and not ethanoic acid.

5c
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5 marks

1.60 g of steel wire was dissolved in an excess of dilute sulfuric acid and the solution made up to 250 cm3.

A 25.0 cm3 portion of this solution required 26.0 cm3 of 0.0200 mol dm–3 potassium manganate(VII) solution for complete reaction. 

Calculate the percentage of iron in the steel wire.

5d
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5 marks

In acidic conditions, potassium manganate(VII) solution can be used to analyse sodium sulfite, Na2SO3, samples according to the following procedure:

Procedure

  1. Dissolve approximately 2.00 g of sodium sulfite in 250.0 cm3 of deionised water
  2. Acidify 25.0 cm3 portions of the sodium sulfite solution
  3. Titrate the sodium sulfite solution with 0.0150 mol dm-3 potassium mangante(VII) solution

The equation for the oxidation of sodium sulfite is given below.

2MnO4 (aq) + 6H+ (aq) + 5SO32– (aq) → 2Mn2+ (aq) + 5SO42– (aq) + 3H2O (l)

A student used 1.97 g of impure sodium sulfite to make their 250.0 cm3 solution. The student recorded an average titre of 27.80 cm3 for their titration experiments.

Determine the percentage of sodium sulfite, Na2SO3, in the impure sample. Show all your working.

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6a
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2 marks

Hydrated copper(II) sulfate can have different amounts of water of crystallisation, x. Compound A has the formula CuSO4•xH2O.

An iodine–thiosulfate titration is performed to find the value of x and the formula of Compound A, using the following procedure.

Procedure

   Step 1 A 5.60 g sample of A is dissolved in water to make a 250.0 cm3 standard    solution.

   Step 2 A 25.00 cm3 sample of this solution is pipetted into a conical flask.

   Step 3 An an excess of KI (aq) is added.

2Cu2+ (aq) + 4I (aq) → 2CuI (s) + I2 (aq)

   Step 4 The resulting mixture is titrated with 0.105 mol dm–3 Na2S2O3 (aq).

2S2O32– (aq) + I2 (aq) → S4O62– (aq) + 2I (aq)

Titration readings

Titration Trial  1 2 3
Final burette reading / cm3  23.20 46.15 22.95 45.95
Initial burette reading / cm3  0.00 23.20 0.05 22.95
Titre / cm3         

Complete the table and calculate the mean titre. 

6b1 mark

In Step 3, suggest why an excess of KI (aq) is added.

6c2 marks

Near the end point of the titration is approached, a solution that can accurately detect the end point is added.

State the solution and explain the colour change observed at the end point.

6d
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4 marks

Determine the formula of compound A. Show your working. 

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1a
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2 marks

Silver(I) ions and iron(II) ions can react according to the following equation:

Ag+ (aq) + Fe2+ (aq) rightwards harpoon over leftwards harpoon  Ag (s) + Fe3+ (aq)


This is an example of a redox equilibrium. The reaction is relatively fast and can reach the equilibrium point within a day.

By titrating the mixture with potassium or sodium thiocyanate solution, the moles of silver ions present at equilibrium can be found. As the thiocyanate solution is added, a white precipitate of silver thiocyanate is formed.


The equation for the reaction that occurs during the titration is:

Ag+ (aq) + SCN (aq) → AgSCN (s)

When all of the silver ions have been used up, the thiocyanate ions then react with iron(III) ions in the mixture to form a red precipitate of iron(III) thiocyanate, Fe(SCN)2+.

 

Explain why no indicator is needed for the titration.

1b
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3 marks

25.0 cm3 of 0.100 mol dm-3 AgNO3 and 25.0 cm3 of 0.100 mol dm-3 FeSO4 were mixed and left to reach equilibrium. 

A solution of 0.0200 mol dm-3 sodium thiocyanate, NaSCN, was then used to titrate against 10.0 cm3 of the equilibrium mixture. An average titre of 21.90 cm3 of sodium thiocyanate was used. 

Calculate the number of moles of AgNO3 remaining in the equilibrium mixture.

1c
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5 marks

Calculate the equilibrium concentrations and therefore the Kc value for this reaction.

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2a
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3 marks

Cobalt, like many transition elements, can form multiple oxidation states. Some states are more stable than others.

The table below shows a range of half-equations possible for cobalt and its ions:

half-equation Eθ / V
Co2+ (aq) + 2e- rightwards harpoon over leftwards harpoon Co (s) -0.28
[Co(NH3)6]3+ (aq) + e- rightwards harpoon over leftwards harpoon[Co(NH3)6]2+ (aq) +0.10
Co3+ (aq) + e- rightwards harpoon over leftwards harpoonCo2+ (aq) +1.80
O2 (g) + 4H+ (aq) + 4e- rightwards harpoon over leftwards harpoon2H2O (l) +1.20
2H+ (aq) + 2e- rightwards harpoon over leftwards harpoonH2 (g) 0.00

Use the values to explain why the Co3+ ion and cobalt metal are unstable in an acidic solution, whilst the Co2+ ion is stable.

2b2 marks

Explain how the Co3+ ion can be stabilised, using information from part a).

2c
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4 marks

Iron can also form two ions, Fe3+ and Fe2+ with the following half-equation: 

 
Fe3+ (aq) + e- rightwards harpoon over leftwards harpoonFe2+ (aq)       Eθ = +0.77 V
 
i)
Predict which of Fe3+ (aq) and Fe2+ (aq) will be stable in an acidic solution.
[2]
ii)
Predict which of Fe3+ (aq) and Fe2+ (aq) will be stable in an acidic solution in the absence of oxygen.

[2]

2d3 marks

Suggest why decreasing the pH increases the electrode potential of the reaction:

 

O2 (g) + 4H+ (aq) + 4e- rightwards harpoon over leftwards harpoon 2H2O (l)

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3a
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3 marks

Electrochemical redox reactions can occur whenever suitable ions are present in an electrolyte. This can be within a synthetic system or within a biological system.

Below are some half-reactions of metals:

half-reaction Eθ / V
Cu2+ (aq) + 2e- rightwards harpoon over leftwards harpoon Cu (s) +0.34
Ag+ (aq) + e- rightwards harpoon over leftwards harpoonAg (s) +0.80
Fe2+ (aq) + 2e- rightwards harpoon over leftwards harpoonFe (s) -0.41
Al3+ (aq) + 3e- rightwards harpoon over leftwards harpoonAl (s) -1.60
Hg2+ (aq) + 2e- rightwards harpoon over leftwards harpoonHg (s) +0.79
Sn2+ (aq) + 2e- rightwards harpoon over leftwards harpoonSn (s) -0.14

Amalgam is an alloy commonly containing mercury, silver, tin and copper and is commonly used for fillings in damaged teeth.  

 
i)
If a person accidentally bites on a piece of aluminium foil they can experience a sharp pain. Suggest a reason for this pain.

[2]

ii)
If a person has multiple fillings they can get a sharp pain without biting on a piece of foil. Suggest why this is still possible.

[1]

3b
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2 marks

When a variety of metal/ metal ion systems a present there are many different potential reactions. 

i)
Calculate the maximum voltage which could be generated in this system of metal ions

[1]

ii)
Explain why the actual voltage is likely to be much lower in a biological system

[1]

3c
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4 marks

Two of the metal/ metal ion half-reactions from the table in part a) were set up in separate cells and were tested against an unknown metal ion/ metal cell, X2+ (aq)/ X(s).

 When the X2+ (aq)/ X(s) was paired with one of the metal ion/ metal cells from the table, it formed the negative electrode and an Eθ cell value of 0.62 V was recorded. 

When the X2+ (aq)/ X(s) was paired with the second of the metal ion/ metal cells from the table, it formed the negative electrode and an Eθ cell value of 1.10 V was recorded.

 
i)
Identify the two different metal ion/ metals used from the table in part a).
[3]
ii)
Hence, deduce the standard cell potential of the X2+ (aq)/ X (s) cell.
[1]
3d
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3 marks

Suggest whether using mercury in amalgam fillings is dangerous, using data from the table.

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4a3 marks

Manganese is one of the metals that can be added to iron in small quantities to make different types of steel.

It is possible to calculate the % of manganese used in a particular steel by first oxidising the manganese to manganate(VII) and then performing a titration.

Nitric acid is added initially, but another chemical is needed to complete the oxidation. The following half-equations and standard cell potentials can help to identify how to complete the oxidation:

half-equation Eθ / V
Mn2+ (aq) + 2e- rightwards harpoon over leftwards harpoonMn (s) -1.17
MnO4- (aq) + 8H+ (aq) + 5erightwards harpoon over leftwards harpoonMn2+ (aq) + 4H2O (l) +1.51
NO3- (aq) + 2H+ (aq) + e- rightwards harpoon over leftwards harpoonNO2 (g) + H2O (l) +0.80
IO3- (aq) + 6H+ (aq) + 5e- rightwards harpoon over leftwards harpoonbevelled 1 halfI2 (s) + 3H2O (l)  +2.00
IO3- (aq) + 3H2O (l) + 6e- rightwards harpoon over leftwards harpoonI- + 6OH- (aq) +0.26

i)
Write an equation to show the reaction between manganese and nitric acid.

[1]

ii)
Identify the oxidising agent needed to complete the oxidation to MnO4- (aq) and write an equation for the reaction that occurs.

[2]

4b6 marks

 Phosphoric acid is used to complete the dissolving of the piece of steel and to prevent the precipitation of iron compounds. The solution containing MnO4- ions is then ready for titration, with a sample of the solution being placed in a conical flask.

i)
Suggest a reactant that would be suitable to place in the burette.
[1]
ii)
Outline the steps needed to perform the titration accurately.

[5]

4c
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5 marks

A 0.800 g sample of high-manganese steel was dissolved in acid and the manganese oxidised to manganate(VII).

The acidified solution was made up to 250 cm3 and 25.0 cm3 portions were titrated against a 0.050 mol dm-3 solution of FeSO4.

An average titre of 17.50 cm3 was recorded.

Calculate the % mass of manganese in the steel.

4d2 marks

Manganate(VII) ions cause a solution, even when very dilute, to take on a pink hue.

Identify a method, other than titration, that could be used to identify the concentration of a solution containing manganate(VII) ions.

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5a
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2 marks

Electrochemical cells can be non-rechargeable or chargeable. A common type of non-rechargeable cell is a small zinc-carbon cell used to power electronic devices.

An example is shown in Figure 1.

zinc-carbon-cell-ocr-a2-sq-h-5-6-q5a

Add the missing label to Figure 1 and state its function.

5b2 marks

At the positive electrode, manganese(IV) oxide reacts with the ammonium chloride in water to form manganese(III) oxide and one other product.

Write the balanced symbol equation for this reaction.

5c1 mark

The zinc is oxidised by the chloride ion at the negative electrode according to the following equation, including spectator ions and assuming that no side reactions occur.

Zn + 2Cl- rightwards arrowZnCl2 + 2e-

Using your answer from part b), write an equation for the overall reaction that occurs when the zinc-carbon cell discharges.

5d
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3 marks

Lead-acid cells are used in cars and other vehicles. One way that the lead-acid cell can discharge is by the following overall equation:

PbO2 (s) + 2H+ (aq) + 2HSO4- (aq) + Pb (s) rightwards arrow2PbSO4 (s) + 2H2O (l)     (Eθ cell = +2.15 V) 

The half-equation at the positive electrode is:

PbO2 (s) + 3H+ (aq) + HSO4- (aq) + 2e- rightwards arrowPbSO4 (s) + 2H2O (l)     (Eθ cell = +1.69 V) 

Deduce the half equation at the negative electrode and calculate the standard electrode potential, Eθ. Write your answer according to the chemical convention for displaying half-equations.

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