AQA A Level Chemistry

Topic Questions

5.4 Electrode Potentials & Electrochemical Cells (A Level only)

1a3 marks

State the standard conditions for measurements in electrochemical cells.

1b1 mark

Figure 1 shows a half cell that can be used to calculate the standard electrode potential of the Fe2+ / Fe reaction.

Figure 1

gKx3Bjz2_1

Write the half equation, including state symbols, that represents this half cell.

1c5 marks

A standard hydrogen electrode can be connected to the Fe2+ / Fe electrode to measure the standard electrode potential, EӨ.

Using the information in part (b), draw the labelled, electrochemical cell that is used to measure the standard electrode potential of the Fe2+ / Fe electrode.

1d3 marks

Write the cell representation for the electrochemical cell set up in parts (b) and (c), using the standard hydrogen electrode and the Fe2+ / Fe electrode.

Did this page help you?

2a2 marks

Explain why platinum is used as the electrode for the standard hydrogen electrode.

2b2 marks

There are three different types of half-cells that can be connected to a standard hydrogen electrode:

  • A metal / metal ion half-cell, e.g. Cu / Cu2+
  • A non-metal / non-metal ion half-cell, e.g. Cl2 / Cl-
  • An ion / ion half-cell, e.g. Fe2+ / Fe3+

State whether each half cell requires a platinum electrode or not.

2c3 marks

Figure 1 shows the electrochemical cell used to measure the standard electrode potential, EӨ, for the Cl2 / Cl- half cell.

Figure 1

3xtFbz3k_3

Write the conventional cell representation for this electrochemical cell.

2d1 mark

Figure 1  is repeated below to help you answer this question.

Figure 1

3xtFbz3k_3

Calculate the standard electrode potential, EӨ, for the Cl2 / Cl- half cell.

Show your workings.

Did this page help you?

3a1 mark

Table 1 contains some standard electrode potential data.

Table 1

Electrode half-equation

EӨ / V

Fe3+ + eFe2+

+0.77

2H+ + 2eH2

0.00


Define the term oxidising agent, in terms of electrons.

3b2 marks

Using the information in Table 1, from part (a), identify the oxidising agent and the reducing agent when the two half cells are connected.

3c3 marks

The standard hydrogen electrode and the Fe2+ / Fe3+ electrode are connected together to make an electrochemical cell as shown in Figure 1.

Figure 1

6LkNdxrX_4

Using the information in Table 1, from part (a), and Figure 1, write the conventional cell representation for this electrochemical cell.

3d2 marks

Potassium nitrate solution is one of the common chemical choices for use in the salt bridge of an electrochemical cell.

State one key feature of the salt bridge in an electrochemical cell.

Did this page help you?

4a3 marks

Name the three types of cell that are used commercially.

4b3 marks

The Daniell cell was one of the first electrochemical cells and was invented in 1836.

Table 1 contains some standard electrode potential data.

Table 1

Electrode half-equation

EӨ / V

Fe3+ + e ⇋ Fe2+

+0.77

2NH4+ + 2e ⇋ 2NH3 + H2

+0.74

NiO(OH) + H2O + e ⇋ Ni(OH)2 + OH

+0.38

Cu2+ + 2e- ⇋ Cu

+0.34

2H+ + 2e ⇋ H2

0.00

Ni2+ + 2e- ⇋ Ni

-0.25

V2+ + 2e- ⇋ V

-0.26

Zn2+ + 2e- ⇋ Zn

-0.76

Cd +  2OH Cd(OH)2 + 2e

-0.82


The Daniell cell is non-rechargeable and has an electromotive force of +1.10 V.

 

Identify the half cells involved in the Daniell cell and demonstrate that they produce an
            electromotive force of 1.10 V.

4c3 marks

Zinc-carbon cells are the most common type of non-rechargeable cells.

One example of a zinc-carbon cell has a standard cell potential of 1.50 V and uses the following reaction at one of the electrodes.

Zn2+ + 2e- ⇋ Zn           EӨ = -0.76 V

i)

Using the information in Table 1, identify the reaction that forms the other half cell. Show your working.

ii)

Suggest why this cell is called a zinc-carbon cell, even though neither reaction involves carbon.

4d2 marks

Car batteries often use a lead-acid system, where lead is the negative electrode and lead (IV) oxide is the positive electrode.

i)

Table 2 contains some standard electrode potential data.

Table 2

Electrode half-equation

EӨ / V

PbO2 + 4H+ + SO42- + 2e ⇋ PbSO4 + 2H2O

+1.70

2NH4+ + 2e ⇋ 2NH3 + H2

+0.74

NiO(OH) + H2O + e ⇋ Ni(OH)2 + OH

+0.38

2H+ + 2e ⇋ H2

0.00

Pb + SO42- ⇋ PbSO4 + 2e

-0.36

Zn2+ + 2e- ⇋ Zn

-0.76


Using the information in Table 2, calculate the standard cell potential of the lead-acid cell.

ii)
Suggest how a car battery can be 12 V.
4e2 marks

State two disadvantages of the use of a lead-acid battery.

Did this page help you?

5a1 mark

Nickel-cadmium, or NiCad, cells are becoming more common. They are available in a variety of sizes and voltages. Their increased cost is balanced out by the fact that they are rechargeable.

The positive electrode of a NiCad cell uses the reaction of cadmium, while the negative electrode uses the reaction of a nickel (II) hydroxide-oxide system.

Cd(OH)2 (s) + 2e ⇋ Cd (s) + 2OH (aq)                                EӨ = -0.82 V 

NiO(OH) (s) + H2O (l) + e ⇋ Ni(OH)2 (s) + OH (aq)            EӨ = +0.38 V

Identify the strongest reducing agent.

5b2 marks

Explain what the standard electrode potential values tell us about the position of the equilibrium for each half reaction.

Cd(OH)2 (s) + 2e ⇋ Cd (s) + 2OH (aq)                                EӨ = -0.82 V

NiO(OH) (s) + H2O (l) + e ⇋ Ni(OH)2 (s) + OH (aq)            EӨ = +0.38 V

5c3 marks
i)

Identify the reduction and oxidation reactions for the nickel-cadmium cell.

Cd(OH)2 (s) + 2e ⇋ Cd (s) + 2OH (aq)                                EӨ = -0.82 V 

NiO(OH) (s) + H2O (l) + e ⇋ Ni(OH)2 (s) + OH (aq)            EӨ = +0.38 V

ii)
Write the overall balanced symbol equation for the nickel-cadmium cell.
5d2 marks

The overall cell potential for the nickel-cadmium cell is 1.20 V.

Explain how this feasible reaction can be reversed to recharge the cell.

Did this page help you?

1a2 marks

Table 1 shows some standard electrode potential data.

Table 1

Electrode half-equation

EӨ / V

Ag+ (aq) + e- Ag (s)

+0.80

Al3+ (aq) + 3e- Al (s)

-1.66

Cl2 (g) + 2e- 2Cl - (aq)

+1.36

ClO2 (aq) + e- ClO2- (aq)

+0.95

Cu2+ (aq) + 2e- Cu (s)

+0.34

Pb4+ (aq) + 2e- Pb2+ (aq)

+1.67


Use data from Table 1 to deduce the species that most readily loses electrons. Justify your answer.

1b4 marks
i)
Use the information in Table 1, part (a), to draw the electrochemical cell for the feasible reaction of Ag / Ag+ and Al / Al 3+.

ii)
Write the conventional representation, including state symbols, for this cell.
1c2 marks

Explain why the salt bridge connecting the silver and aluminum electrodes cannot be made with potassium chloride solution.

1d2 marks

Use the information in Table 1, to write the conventional representation, including state symbols, for the electrochemical cell with the lowest electromotive force.

Table 1

Electrode half-equation

EӨ / V

Ag+ (aq) + e- Ag (s)

+0.80

Al3+ (aq) + 3e- Al (s)

-1.66

Cu2+ (aq) + 2e- Cu (s)

+0.34

Cl2 (g) + 2e- 2Cl - (aq)

+1.36

ClO2 (aq) + e- ClO2- (aq)

+0.95

Pb4+ (aq) + 2e- Pb2+ (aq)

+1.67

Did this page help you?

2a2 marks

Figure 1 shows a non-rechargeable, zinc-carbon dry cell that can be used to power electronic devices.

Figure 1

3a

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

2b2 marks

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

Write the balanced symbol equation for this reaction.

2c1 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 - → ZnCl2 + 2e

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

2d3 marks

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

PbO2 (s) + 2H+ (aq) + 2HSO4- (aq) + Pb (s) → 2PbSO4 (s) + 2H2O (l)            EcellӨ = +2.15 V

The half equation at the positive electrode is

PbO2 (s) + 3H+ (aq) + HSO4- (aq) + 2e- → PbSO4 (s) + 2H2O (l)                     EӨ = +1.69 V

Deduce the half equation at the negative electrode, including the electrode potential, EӨ.

Your answer should be written according to electrochemical convention.

Did this page help you?

3a2 marks

Explain why Figure 1 does not represent the standard hydrogen electrode.

Figure 1

wya3iWDO_4

3b2 marks

The approximate cost of platinum in the UK is £22 500 per kg.

Explain how the design of the platinum electrode minimises the cost while maximising its efficiency.

3c3 marks

A student set up an electrochemical cell as shown in Figure 2.

Figure 2

KgNL~I2y_5

The student made the silver electrode by adding 7.78 g of silver sulfate to 25.0 cm3 of water at 298 K.

The solubility of silver sulfate in water at 298 K is 4.73 g dm-3.

Explain if the setup in Figure 2 can be used to measure the standard electrode potential of the Ag / Ag+ half cell. Show your working.

3d2 marks

The cell represented in Figure 3 was set up under standard conditions.

Figure 3

lvyYNaxa_6

Explain why this electrochemical cell cannot measure the standard electrode potential of the Fe2+ / Fe3+ half cell.

Did this page help you?

4a3 marks

The 2019 Nobel Prize for Chemistry was awarded to three scientists for their work on the development of lithium-ion cells.

One of the first lithium-ion batteries developed was the lithium titanium (III) sulfide cells shown in Figure 1, which had an electromotive force of around 2.0 V.

Figure 1

ND-94P8k_7

The cell works according to the following overall reaction.

Li (s) + TiS2 (s) → LiTiS2 (s)

Use the information to deduce the ionic half equations at the positive and negative electrodes and write the conventional representation for this cell.

4b1 mark

Suggest one reason, other than the reactivity of lithium, why the lithium titanium (III) sulfide cell, in part (a), was not commercially produced on a large scale.

4c2 marks

One major development in lithium-ion technology was replacing the titanium disulfide cathode with a cobalt (IV) oxide cathode, as shown in Figure 2.

Figure 2

WoZbuwwu_8

The standard cell potential, Ecell, of the lithium cobalt (III) oxide cell was increased by 95% compared to 2.0 V of the lithium titanium (III) sulfide cell in part (a).

Table 1 shows the electrode half-equations and some standard electrode potential data for the lithium cobalt (IV) oxide cell.

Table 1

Electrode half-equation

Eϴ / V

Li+ + e  Li

-3.00

Li+ + CoO2 + e  LiCoO2

To be calculated


Use the information to calculate the standard electrode potential for the lithium cobalt (III) oxide half cell.

4d2 marks

Commercial lithium ion batteries often have a graphite lattice filled with lithium ions as the negative terminal. In industry, the overall reaction for the discharge of a rechargeable lithium ion cell can be shown as

LiC6 + CoO2 → LiCoO2 + C6

i)
Write an equation for the overall reaction that occurs when this lithium cell is being recharged.

ii)

Suggest why recharging the lithium cobalt (IV) oxide cell may lead to increased carbon dioxide emissions.

4e2 marks

In certain circumstances a fuel cell, such as the hydrogen-oxygen fuel cell, can be used in preference to a rechargeable cell.

The cell representation of an alkaline hydrogen-oxygen fuel cell is:           

Pt (s) | H2 (g) | H2O (l) || O2 (g) | OH- (aq) | Pt (s)

Write the half equations for the reaction that occurs at each electrode.

4f6 marks

Describe the benefits and drawbacks of the commercial and large-scale use in the transportation industry of hydrogen-oxygen fuel cells.

Did this page help you?

5a2 marks

Table 1 shows some redox half-equations and standard electrode potentials.

Table 1

Electrode half-equation

Eϴ / V

MnO4- (aq) + 8H+ (aq) + 5e- ⇋ Mn2+ (aq) + 4H2O (l)

+1.52

O2 (g) + 2H+ (aq) + 2e- ⇋ H2O2 (aq)

+0.68


Use the data given to determine if the following reaction is feasible at 298 K.

2KMnO4 (aq) +5H2O2 (aq) + 6HCl (aq) → 2MnCl2 (aq)+ 8H2O (l)+ 5O2 (g) + 2KCl (aq)

5b4 marks

Some redox half-equations and standard electrode potentials are shown in Table 2.

Table 2

Electrode half-equation

EӨ / V

Ag+ (aq) + e- ⇌  Ag (s)

+0.80

Al 3+ (aq) + 3e- ⇌  Al (s)

-1.66

Cu2+ (aq) + 2e- ⇌  Cu (s)

+0.34

Cl2 (g) + 2e- ⇌  2Cl- (aq)

+1.36

ClO2 (aq) + e- ⇌  ClO2- (aq)

+0.95

Pb4+ (aq) + 2e- ⇌  Pb2+ (aq)

+1.67


Write the conventional representation for the cell that has an electromotive force of 
-0.61 V.

5c3 marks

The standard electrode potentials for the following reaction of copper sulfate with hydrogen is shown in Table 3.

CuSO4 (aq) + H2 (g) → Cu (s) + H2SO4 (aq)

Table 3

Electrode half-equation

EӨ / V

Cu2+ (aq) + 2e-   Cu (s)

+0.34

 2H+. (aq) + 2e H2 (g)

0.00

i)
Use the information to explain why the reaction is thermodynamically feasible.

ii)

Suggest a reason why the reaction does not occur despite being thermodynamically feasible.

Did this page help you?

1a2 marks

Table 1 below contains some standard electrode potential data which you will need to answer the following questions.

Table 1

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

1b2 marks

Give the conventional representation of the cell that is used to measure the standard electrode potential of chlorine to chloride ions as shown in Table 1 in part (a).

1c1 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 EMF of this cell using the values given in Table 1 in part (a).

1d3 marks

Two half-cells, involving species in Table 1, are connected together to give a cell with an e.m.f. = +0.30 V.

i)
Determine which two half equations are connected using the data from Table 1 giving the conventional representation for the cell.

ii)

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

Did this page help you?

2a6 marks

The diagram shown in Figure 1 represents the standard hydrogen electrode.

Figure 1

YgsBUorK_5

i)
Name the substance used as the electrode in Figure 1

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.

i)

Complete the diagram (Figure 2) to show the components and reagents, including their concentrations and label any apparatus required to complete the electrochemical cell.

Figure 2

Igw11YTw_6

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 copper and silver as shown in the following cell representation: 

Figure 3

Cu | Cu2+ || Ag+ | Ag

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

A diagram of a cell is shown below in Figure 3.

Figure 3

hIqb-jOV_7

i)
Explain how the salt bridge, in Figure 3, 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.

Did this page help you?

3a3 marks

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  rightwards harpoon over leftwards harpoonLi+[CoO2]                 Eᶿ = 0.36V

Negative electrode                Li+ + e  rightwards harpoon over leftwards harpoon Li                                            Eᶿ = -3.04V

i)
Write the representation for this cell.

ii)
Calculate the standard electrode 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      charging.

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  rightwards arrow  Cd (s)  +  2OH(aq)                                     Eϴ = -0.88

NiO(OH) (s)  + H2O (I) + e  rightwards arrow Ni(OH)2 (s) + OH(aq)                     Eϴ = +0.52

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.

Did this page help you?

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.

Figure 1 shows a hydrogen-oxygen fuel cell.

i)

Using Figure 1 below, explain how an electric current is generated in the fuel cell.

Figure 1

iU-TCsVk_9

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 has an environmental advantage over the internal combustion engine. 

iii)

Suggest why hydrogen fuel cells cannot be classed as carbon neutral.

Did this page help you?

5a3 marks

Electrochemical cells involve redox reactions, these allow the electrons to flow from the more reactive metal to the less reactive metal. The voltage, or electrode potential, can then be measured using a voltmeter.

i)
Define the terms redox reaction and standard electrode potential. 

ii)
State the IUPAC convention for writing half equations.
5b3 marks

When measuring electrode potentials standard conditions must be used. This is because changing conditions including the temperature, pressure and concentration will affect the value of the electrode potential.

State the standard conditions used to measure electrode potentials and explain why changing these conditions will affect the electrode potential.

5c2 marks

Electrode potentials can be determined by constructing a cell using a standard hydrogen electrode and the electrode being investigated.

Electrochemical cells are made by dipping each metal into a salt solution containing the metal’s ions.

i)
Suggest why the electrode in a standard hydrogen cell is made of platinum.

ii)
State why the standard hydrogen electrode is shown on the left.
5d2 marks

If a cell reaction is to occur spontaneously, the overall cell potential must be positive.

Give the equation to show the spontaneous reaction which takes place and calculate the EMF for this cell reaction.

Fe3+ (aq) + e  → Fe2+ (aq) + 0.77 V

Ce4+ (aq) + e  → Ce3+ (aq) + 1.45 V

Did this page help you?