Electrode Potentials & Electrochemical Cells (AQA A Level Chemistry)

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

Table 1

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

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).

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).

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).

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

Figure 1

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).

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

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

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.

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

Figure 3

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.

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 = 0.36V

Negative electrode                Li⁺ + e^–   Li                                            Eᶿ = -3.04V

i) Write the representation for this cell.

ii) Calculate the standard electrode potential of the cell.

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.

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.

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)                                     E^ϴ = -0.88

NiO(OH) (s)  + H₂O (I) + e^–   Ni(OH)₂ (s) + OH^– (aq)                     E^ϴ = +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.

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.

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

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

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.

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.

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.

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.

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.

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

Table 1 shows some standard electrode potential data.

Table 1

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

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.

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

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

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

Figure 1

Complete Figure 1 by adding the missing label and stating its function.

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.

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.

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)            E_cell^Ө = +2.15 V

The half equation at the positive electrode is

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

Deduce the half equation at the negative electrode, including the electrode potential, E^Ө.

Your answer should be written according to electrochemical convention.

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

Figure 1

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.

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

Figure 2

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.

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

Figure 3

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

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

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.

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.

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

The standard cell potential, E_cell, 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

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

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.

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.

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

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

Table 1

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)

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

Table 2

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

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

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.