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Question 1
a) | Define the term standard electrode potential.
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b) | The table below shows the standard electrode potentials, E⦵. of two half-equations. |
b) | Write the conventional representation of the cell used to measure the standard electrode potential for the Ag+/Ag electrode.
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c) | A standard Cu2+(aq)/Cu(s) half-cell is connected to a standard Ag+(aq)/Ag(s) half-cell. The potential of the cell is measured.
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Write down the equation for the overall cell reaction. Include state symbols.
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d) | Give the value of E⦵ cell for the cell in part c).
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e) | Water is then added to the half-cell. Explain, in terms of equilibrium, how the cell potential changes when water is added.
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Question 2
Chlorine dioxide, ClO2, has several uses including in the purification of water. ClO2 is made by reacting ClO3- with concentrated hydrochloric acid. 2ClO3- + 4H+ + 2Cl- ⇌ 2ClO2 + 2H2O + Cl2 The table below shows some electrode potential data. |
a) | Use data in the table above to explain why the formation of ClO2 via the equation above does not occur under standard conditions.
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b) | Suggest why the forward reaction for the formation of ClO2 does occur in the presence of concentrated hydrochloric acid.
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c) | A Cu2+(aq)/Cu(s) half-cell is used to confirm the E⦵ of a Cl2(aq)/Cl-(aq) half-cell. |
Complete and label the diagram of the apparatus the student would set up. Include state symbols.
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d) | Give the value of E⦵ cell for the cell in part c).
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e) | State and explain which way the electrons move when the cell in part c) delivers a current.
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Question 3
a) | State the meaning of the term electochemical series.
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b) | Standard electrode potentials for three redox systems are shown in the table below. |
Use information in the table below to predict three reactions that might be feasible. Write the overall equation for each predicted reaction.
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c) | Give two reasons why it is uncertain whether reactions predicted from E⦵ values may actually take place.
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d) | In the table above, state the strongest reducing agent and the strongest oxidising agent. Explain your reasoning.
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Question 4
The table below shows two standard electrode potentials, E⦵. |
a) | Calculate E⦵cell.
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b) | Write the cell equation that produced the E⦵cell value calculated in part a).
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Question 5
The diagram below shows a non-rechargeable cell that can be used to power electronic devices.
Standard electrode potentials of the relevant half-equations for this cell are shown in the table below. |
a) | Use data in the table to calculate the e.m.f of this cell.
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b) | Write an equation for the overall reaction that occurs when the cell discharges.
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c) | Identify which of A or B in the diagram is the positive electrode. Explain your reasoning.
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d) | Suggest one reason, other than cost, why the Zn-MnO2 cell is not recharged.
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Question 6
The e.m.f of a rechargeable lithium cell is 2.90 V. The simplified reactions taking place at each electrode are: Electrode A: Li+ + CoO2 + e- ➔ LiCoO2 E⦵ = -0.15 V Electrode B: Li+ + e- ➔ Li Electrode B is the negative electrode. |
a) | Calculate a value for the electrode potential, E⦵, of electrode B.
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b) | Write an equation for the overall reaction that occurs when this lithium cell is being recharged.
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c) | Deduce the oxidation state of the cobalt in this cell before recharging begins.
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d) | State one environmental advantage of this rechargeable lithium cell compared with non-rechargeable cells.
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Question 7
Hydrogen–oxygen fuel cells are used to provide electrical energy for electric motors in vehicles. In an alkaline hydrogen–oxygen fuel cell, a current is generated that can be used to drive an electric motor. The reactions that occur are: 2H2(g) + 4OH-(aq) ➔ 4H2O(l) + 4e- O2(g) + 2H2O(l) + 4e- ➔ 4OH-(aq)
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a) | Use these half-equations to explain how an electric current can be generated.
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b) | Explain why a fuel cell does not need to be recharged.
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c) | To provide energy for a vehicle, hydrogen can be used either in a fuel cell or in an internal combustion engine. Suggest one advantage and one disadvantage of using hydrogen in a fuel cell rather than in an internal combustion.
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d) | Direct-ethanol fuel cells (DEFCs) are being developed in which the fuel is ethanol rather than hydrogen.
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Suggest two advantages of using ethanol as the fuel in a fuel cell rather than hydrogen.
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Question 8
a) | Draw a fully labelled diagram of the experimental set-up you could use to measure the standard electrode potential of the V3+(aq)/V2+(aq) electrode. Include the necessary chemicals.
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b) | State the three standard conditions needed for this measurement.
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c) | The salt bridge typically contains a solution of potassium nitrate. Give two reasons why potassium nitrate can be used as a salt bridge.
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d) | The table below shows some standard electrode potentials. |
Suggest how the E for the V3+(aq)/V2+(aq) electrode would differ from its E⦵ value if the concentration of V2+(aq) ions is reduced. Explain your answer.
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e) | A solution containing iodide ions, I-, was added to an acidified solution containing vanadium(V) ions, VO2+. Predict the oxidation state of the vanadium ions left at the end of the reaction. Justify your prediction by calculating the E⦵cell for any relevant reactions. Write the ionic equation for any reaction(s) occurring.
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Question 9
Rechargeable lead-acid cells are used in car batteries. Each cell consists of a negative electrode made of Pb metal and a positive electrode made of PbO2. The electrolyte is H2SO4(aq). When a lead-acid cell is in use, Pb2+ ions are precipitated out as PbSO4(s) at the negative electrode. Pb(s) + SO42-(aq) ➔ PbSO4(s) + 2e-
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a) | Complete the half-equation for the reaction taking place at the positive electrode. PbO2(s) + SO42-(aq) + .............. + .............. ➔ PbSO4(s) + ..............
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b) | Complete the graphs below to show how the voltage across a lead-acid cell and a H2/O2 fuel cell changes with time when each cell is used to provide an electric current.
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c) | Suggest a reason why;
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