Standard Electrode Potentials E⦵, Standard Cell Potentials E⦵cell & the Nernst Equation (Cambridge (CIE) A Level Chemistry): Exam Questions

This question is about electrochemical cells.

Fig. 1.1 shows the apparatus used to measure the standard electrode potential, E^θ, of a cell composed of a Cu(II)/Cu half-cell and an Fe(II)/Fe half-cell.

Fig. 1.1

Complete the diagram by adding components to show the complete circuit. Label the components you add.

In the spaces below, identify what the four letters A–D in Fig. 1.1 represent.

The standard electrode potentials of the two half-cells are as follows.

Fe²⁺ + 2e^– ⇌ Fe E^θ = −0.44 V

Cu²⁺ + 2e^– ⇌ Cu E^θ = +0.34 V

i) Identify which electrode acts as the negative electrode.

[1]

ii) Construct the overall equation for this electrochemical cell.

[1]

E^θ_cell of this cell is +0.78 V.

State whether this reaction is feasible. Explain your answer.

This question is about electrochemical cells.

The diagram shown in Fig. 3.1 represents the standard hydrogen electrode.

Fig 3.1

i) Name the substance used as the electrode in Fig 3.1.

[1]

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

[2]

iii) Give the standard conditions used in a standard hydrogen electrode.

[3]

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

i) Complete Fig. 3.2 to show the components and reagents, including their concentrations and label any apparatus required to complete the electrochemical cell.

[4]

Fig. 3.2

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

[2]

A student set up another electrochemical cell consisting of copper and silver as shown in the following cell representation where the half cell with the greatest negative standard electrode potential is written on the left:

Cu / Cu²⁺ ; Ag⁺ / Ag

i) Construct a half-equation for the reaction that occurs at the positive electrode.

[1]

ii) Construct a half-equation for the reaction that occurs at the negative electrode.

[1]

iii) Use the half-equations to deduce an overall equation for the cell. Include all state symbols.

[2]

A diagram of a cell is shown below in Fig. 3.3.

Fig. 3.3

i) Explain how the salt bridge, in Fig. 3.3, provides an electrical connection between the two solutions.

[1]

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

[1]

The apparatus shown was used to measure the standard electrode potential for the reduction of Cr₂O₇^2– ions to Cr³⁺ ions in acid solution:

Cr₂O₇^2–(aq) + 14H⁺(aq) + 6e^– ⇌ 2Cr³⁺(aq) + 7H₂O(l)

Which material should be used for each electrode?

Suggest a suitable solution that could be used as solution 1.

Solution 2 contains 14.71 g of K₂Cr₂O₇.

Calculate the mass of Cr₂(SO₄)₃⋅18H₂O that should also be used. Show your working.

[M_r values: K₂Cr₂O₇ = 294.2, Cr₂(SO₄)₃⋅18H₂O = 716.3]

Solution 2 is best acidified with H₂SO₄ instead of HCl or HBr.

Suggest why.

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

Fig. 2.1

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

Table 2.1

[2]

How would the actual E_cell of the above cell compare to the E^θ_cell under standard conditions?

Explain your answer.

How would the E_cell 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 E_cell you have stated in (i) and (ii).

[2]

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

The standard electrode potentials for seven different redox systems are shown in Table 3.1.

Table 3.1

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

i) Construct the overall equation for the cell reaction.

[1]

ii) Calculate the voltage of this cell.

[1]

i) Using redox systems 5, 6 and 7 only in Table 3.1, construct 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]

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]

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

Table 4.1 below contains some standard electrode potential data which you will need to answer the following questions.

Table 4.1

Deduce the species from Table 4.1 that is the weakest oxidising agent. Explain your choice.

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).

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.

[4]

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]

The standard electrode potentials in Table 5.1 can be used to predict redox reactions.

Table 5.1

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

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]

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

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

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.

Calculate the electrode potential, E, for redox system 2 under these conditions. Show your working.

Define the term standard electrode potential.

Draw a labelled diagram to show how the standard electrode potential of a Cu²⁺/Cu half-cell can be measured relative to the Standard Hydrogen Electrode (SHE).

Identify the chemicals, conditions, and relevant pieces of apparatus.

A current of 1.50 A is passed through aqueous CuSO₄ for 25.0 minutes.

Calculate the mass of copper deposited at the cathode. Show your working.

Use the Nernst equation from the Data Booklet to calculate the electrode potential (E) for the Fe³⁺/Fe²⁺ half-cell where [Fe³⁺] = 0.50 mol dm^-3 and [Fe²⁺] = 0.01 mol dm^-3 (E^θ = +0.77 V). Show your working.

A cell is constructed using the Fe³⁺/Fe²⁺ and Cu²⁺/Cu half-cells under standard conditions. The equation for the overall cell reaction is shown.

2Fe³⁺ (aq) + Cu (s) → 2Fe²⁺ (aq) + Cu²⁺ (aq)

The standard cell potential, E^θ_cell, for this reaction is +0.43 V.

Calculate the standard Gibbs free energy change, ΔG^θ, for this reaction. Give your answer in kJ mol^-1 to three significant figures. Show your working.

Explain why Fig. 3.1 does not represent the standard hydrogen electrode.

Fig. 3.1

The standard electrode potential for Zn²⁺ (aq) + 2e^– → Zn (s) is –0.76 V.

State the meaning of the minus sign in the value of –0.76 V.

Zinc coating on metals serves as physical protection which prevents rust from affecting the underlying metal surface. This is achieved by electroplating as shown in Fig. 3.2.

Fig. 3.2

i) Suggest a suitable solution to act as the electrolyte during zinc electroplating. 

[1]

 ii) Complete the diagram by labelling the polarity of the power source by using a + and - sign.  

[1]

Calculate the length of time, in hours, required to deposit 1.0 g of zinc on the item to be electroplated. Assume the current is a constant 0.1 A throughout this period.

State your answer to 3 significant figures and show your working.

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.

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

[3]

The standard electrode potential, E^θ, of the Ag⁺ (aq) / Ag (s) half-cell is +0.80 V.

The electrode potential of an Ag⁺ (aq) / Ag (s) half-cell was measured at 25 °C and found to be +0.72 V.

Use the Nernst equation from the Data Booklet to calculate the concentration of silver ions, in mol dm^-3, in this half-cell. Show your working.

Table 3.1 lists electrode potentials for some electrode reactions.

Table 3.1 

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

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

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

i) Use Table 3.1 to identify an acid that will oxidise copper. Explain your answer.

[2]

ii) Construct an equation for the reaction.

[1]

iii) Calculate the E^θ_cell for the same overall reaction.

[1]

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

i) Use Table 3.1 to explain why [Co(H₂O)₆]³⁺ (aq) will undergo a redox reaction with [Fe(H₂O)₆]²⁺ (aq).

[1]

ii) Give an equation for this reaction.

[2]