Electrochemistry - exam practice
Electrochemistry
Question 1 of 11
Complete the table, identifying the substances liberated at each electrode during electrolysis with inert electrodes.
Molten lead(II) bromide is electrolysed in an inert atmosphere with inert electrodes.
Write ionic equations for the reactions occuring at each electrode.
The electrolysis of molten lead(II) bromide is a redox process.
Identify the ion that is oxidised and the ion that is reduced.
Explain your answer in terms of oxidation numbers.
Describe two visual observations that would be made during the electrolysis of molten lead(II) bromide.
An oxide of iron dissolved in an inert solvent is electrolysed for 1.50 hours using a current of 1.20 A. The electrolysis products are iron and oxygen. The mass of iron produced is 1.25 g.
Calculate the oxidation number of Fe in the oxide of iron.
Write ionic equations for the equations occuring at each electrode.
Calculate the volume of oxygen produced, in dm3, at room temperature and pressure, during the electrolysis of the oxide of iron.
Give your answer to 2 significant figures.
Calculate the value of the electrode potential at 298 K of a Co2+(aq)/Co(s) electrode that has a concentration of aqueous Co2+ ions of 5.00 x 10-2 mol dm-3.
Give your answer to 2 decimal places.
E⦵ = -0.28 V.
Calculate the value of the electrode potential at 298 K of a Fe3+(aq)/Fe2+(aq) half-cell.
Give your answer to 2 decimal places.
[Fe3+] = 3.45 x 10-3 mol dm-3
[Fe2+] = 5.60 x 10-2 mol dm-3
E⦵ = +0.77 V.
Compound Y is an iodate(V) salt of a group 1 metal.
The iodate(V) ion has the formula IO3-.
A scientist carries out a titration to find the formula of compound Y.
Step 1: the scientist dissolves 1.39 g of Y in water and makes up the solution to 250.0 cm3 in a volumetric flask.
Step 2: the scientist pipettes 20.00 cm3 of the solution of Y into a conical flask, followed by 10 cm3 of dilute sulfuric acid and an excess of KI(aq).
The iodate(V) ions are reduced to iodine, as shown below.
IO3-(aq) + 6H+(aq) + 5I-(aq) ➔ 3I2(aq) + 3H2O(l)
Step 3: the resulting mixture is titrated with 0.200 mol dm-3 Na2S2O3(aq).
2S2O32-(aq) + I2(aq) ➔ S4O62-(aq) + 2I-(aq)
The scientist's titration readings are shown in the table below.
Complete the table.
Calculate the mean titre, in cm3, that the scientist should use for analysing the results.
Describe and explain how the student should determine the end point of this titration accurately.
Determine the relative formula mass and formula of the group 1 iodate(V), Y.
Define the term standard electrode potential.
The table below shows the standard electrode potentials, E⦵. of two half-equations.
Write the conventional representation of the cell used to measure the standard electrode potential for the Ag+/Ag electrode.
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.
Write down the equation for the overall cell reaction.
Include state symbols.
Give the value of E⦵ cell for the cell in part c).
Water is then added to the half-cell.
Explain, in terms of equilibrium, how the cell potential changes when water is added.
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.
Use data in the table above to explain why the formation of ClO2 via the equation above does not occur under standard conditions.
Suggest why the forward reaction for the formation of ClO2 does occur in the presence of concentrated hydrochloric acid.
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.
Give the value of E⦵ cell for the cell in part c).
State and explain which way the electrons move when the cell in part c) delivers a current.
State the meaning of the term electochemical series.
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.
Give two reasons why it is uncertain whether reactions predicted from E⦵ values may actually take place.
In the table above, state the strongest reducing agent and the strongest oxidising agent.
Explain your reasoning.
The table below shows two standard electrode potentials, E⦵.
Calculate E⦵cell.
Write the cell equation that produced the E⦵cell value calculated in part a).
The electrochemical equation for standard free-energy change is:
ΔG = nFE⦵
Calculate the standard free-energy change, in kJ mol-1, for the Cd2+/Cd and Fe/Fe2+ electrochemical cell.
Give your answer to 3 significant figures.
The Faraday constant, F = 96,500 C mol-1
Use your answer to part c), to determine whether the cell reaction is feasible.
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.
State the three standard conditions needed for this measurement.
The salt bridge typically contains a solution of potassium nitrate.
Give two reasons why potassium nitrate can be used as a salt bridge.
The table below shows some standard electrode potentials.
| Redox system | Electrode half-equation | E^⦵^ (V) |
|---|---|---|
| 1 | V^2+^(aq) + 2e^-^ ⇌ V_(s)_ | -1.18 |
| 2 | V^3+^(aq) + e^-^ ⇌ V^2+^(aq) | -0.26 |
| 3 | VO^2+^(aq) + 2H^+^(aq) + e^-^ ⇌ V^3+^(aq) + H_2_O_(l)_ | +0.34 |
| 4 | I_2(aq)_ + 2e^-^ ⇌ 2I^-^(aq) | +0.54 |
| 5 | VO_2_^+^(aq) + 2H^+^(aq) + e^-^ ⇌ VO^2+^(aq) + H_2_O_(l)_ | +1.00 |
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.
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⦵cellfor any relevant reactions.
Write the ionic equation for any reaction(s) occurring.
Explain, using information in the table, whether the dispropotionation of V2+(aq) into V3+(aq) and V(s) is feasible under standard conditions.
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-
Calculate the time taken to convert 2.50 g of Pb to PbSO4 when a current of 0.600 A is delivered.
Give your answer to the nearest minute.
Complete the half-equation for the reaction taking place at the positive electrode.
PbO2(s) + SO42-(aq) + .............. + .............. ➔ PbSO4(s) + ..............
Potassium sulfite, K2SO3, is used as a preservative in some foods and drinks.
Food safety laws permit a maximum of 0.500 g of K2SO3 per kg of sausage meat.
A scientist uses a manganate(VII) titration to determine how much K2SO3 a sample of sausage meat contains.
Step 1: the K2SO3 from 1.00 kg of sausage meat is extracted to form a solution containing aqueous SO32- ions.
Step 2: the solution from step 1 is made up to 250.0 cm3 in a volumetric flask with water. 25.0 cm3 of this diluted solution is pipetted into a conical flask.
Step 3: the pipetted solution from step 2 is acidified with dilute sulfuric acid and then titrated with 1.20 x 10-2 mol dm-3 solution of potassium manganate(VII), KMnO4.
The equation for this titration is:
2MnO4-(aq) + 6H+(aq) + 5SO32-(aq) ➔ 2Mn2+(aq) + 3H2O(l) + 5SO42-(aq)
10.40 cm3 of KMnO4(aq) is required to reach the endpoint.
Analyse the results to determine whether the amount of potassium sulfite in sausage meat is above or below the legal limit.