IB Chemistry R3.2: Electron transfer reactions worked solutions

Which change is an oxidation? (A) \ce{MnO4- -> Mn^2+} (B) \ce{SO3^2- -> SO4^2-} (C) \ce{Cl2 -> 2Cl-} (D) \ce{Fe^3+ -> Fe^2+}

A cell is written \ce{Fe(s)}\,|\,\ce{Fe^2+(aq)}\,\|\,\ce{Cu^2+(aq)}\,|\,\ce{Cu(s)}, with a potassium nitrate salt bridge. Which statement is correct while it supplies current? (A) Electrons pass through the salt bridge. (B) Nitrate ions move towards the iron half-cell. (C) Copper is the anode. (D) The iron electrode gains mass.

One mole of each compound is completely hydrogenated with a nickel catalyst. Which uses the most hydrogen? (A) hex-1-ene (B) hex-1-yne (C) benzene (D) cyclohexene

The standard electrode potentials are \ce{Al^3+ + 3e- <=> Al}, -1.66\,\mathrm{V}, and \ce{Cu^2+ + 2e- <=> Cu}, +0.34\,\mathrm{V}. What is E^\ominus_\mathrm{cell} for an aluminium-copper cell? (A) +1.32\,\mathrm{V} (B) +2.00\,\mathrm{V} (C) +4.34\,\mathrm{V} (D) -2.00\,\mathrm{V}

Find the oxidation state of V in \ce{VO2+}. The ion contains one V and two O atoms. (a) State the value. [1]

In \ce{2Al + 3Ni^2+ -> 2Al^3+ + 3Ni}, which species is the oxidizing agent? (A) \ce{Al} (B) \ce{Al^3+} (C) \ce{Ni^2+} (D) \ce{Ni}

A little bromine water is mixed separately with aqueous potassium iodide and potassium chloride. (a) Identify the solution in which a displacement occurs and name the halogen formed. [2]

A clean iron nail reacts with dilute sulfuric acid. (a) Write the ionic equation, including the formula and state of the gas. [2]

Acidified permanganate oxidizes ethanedioate ions, \ce{C2O4^2-}, to carbon dioxide. (a) Write the half-equation for the oxidation of ethanedioate. [1] (b) Combine it with \ce{MnO4- + 8H+ + 5e- -> Mn^2+ + 4H2O} to give the overall ionic equation. [2] (c) A 25.0\,\mathrm{cm^3} sample needs 20.00\,\mathrm{cm^3} of 0.0200\,\mathrm{mol\,dm^{-3}} permanganate. Calculate the ethanedioate concentration. [2]

A zinc half-cell and a nickel half-cell are joined by a potassium nitrate salt bridge. Electrons flow in the wire from zinc to nickel. (a) State which electrode is the anode and give its sign. [1] (b) Write the half-equation at each electrode. [2] (c) State which way nitrate ions move, and how the mass of the nickel electrode changes. [1]

Pentan-1-ol is warmed with acidified dichromate(VI). (a) Name the product obtained if it is distilled out as it forms. [1] (b) Heated under reflux with excess oxidant instead, it gives a different product. Name it and write the equation using [\ce{O}]. [2] (c) State the colour change of the oxidizing agent. [1]

Metals P, Q and R form 2+ ions. P deposits Q from \ce{Q^2+} solution; R deposits P from \ce{P^2+} solution. (a) Arrange the metals in increasing ease of oxidation. [1] (b) Write the ionic equation when R is added to \ce{Q^2+} solution and identify the reducing agent. [2]

Molten magnesium chloride is electrolysed using inert electrodes. (a) Write both electrode half-equations, naming the electrodes. [2] (b) State the energy conversion and the charge carrier in the melt. [1]

On discharge, a simplified secondary cell has the overall reaction \ce{A + B^2+ -> A^2+ + B}. (a) Write the half-equation at the A electrode on discharge and again on recharge. [2] (b) Explain why the A electrode changes from anode to cathode. [1]

A remote sensor can use either a primary cell with 80 h of service or a secondary cell with 50 h per charge. A visit to replace or recharge the cell is expensive. (a) Give one operational advantage of the primary cell in this setting. [1] (b) Give one resource advantage of the secondary cell, and one condition needed for that advantage to be realized. [2]

A sample of pentan-2-one is treated with sodium borohydride and then aqueous acid. (a) Name the organic product. [1] (b) Write an equation using [\ce{H}] and state the change in oxidation state of the carbonyl carbon. [2]

Supplied reduction potentials: For \ce{2Fe^3+ + Sn^2+ -> 2Fe^2+ + Sn^4+}, use E^\ominus(\ce{Fe^3+/Fe^2+})=+0.77\,\mathrm V and E^\ominus(\ce{Sn^4+/Sn^2+})=+0.15\,\mathrm V. (a) Calculate E^\ominus_\mathrm{cell}. [1] (b) Calculate \Delta G^\ominus in \mathrm{kJ\,mol^{-1}}. [2] (c) A student writes E^\ominus_\mathrm{cell}=2(0.77)-0.15. State the mistake. [1] (d) Explain why potentials are measured against the standard hydrogen electrode, and state the hydrogen pressure. [1]

Aqueous solutions are electrolysed; electrodes are inert unless stated. (a) Concentrated sodium chloride solution: name the product at each electrode. [2] (b) Copper(II) sulfate solution, copper electrodes: state what happens at the anode and to the blue colour. [2] (c) A steel key is plated with copper here. Which terminal of the supply is the key joined to? [1] (d) In the copper(II) sulfate cell, replace the copper anode with graphite while maintaining the current. Deduce the changes in blue colour and acidity, with an anode equation. [2]

A proposed recovery process passes \ce{Fe^2+} solution over silver metal. Use E^\ominus(\ce{Fe^2+/Fe})=-0.45\,\mathrm V and E^\ominus(\ce{Ag+/Ag})=+0.80\,\mathrm V. (a) Write the proposed overall equation and calculate its standard cell potential. [2] (b) Assess whether this process deposits iron under standard conditions, and explain what reversing the reaction would achieve. [2] (c) Explain why a favorable potential alone does not specify a reaction rate. [1]

A copper coating is deposited on a steel part using aqueous copper(II) sulfate and a copper anode. The current is 0.800\,\mathrm A for 2412.5\,\mathrm s; the measured coating mass is 0.572\,\mathrm g. Use F=96500\,\mathrm{C\,mol^{-1}} and M(\ce{Cu})=63.55\,\mathrm{g\,mol^{-1}}. (a) Write both electrode equations and state the terminal connected to the steel. [2] (b) Calculate the theoretical coating mass. [2] (c) Calculate the current efficiency and suggest one process that could make it less than 100 percent. [2]

R3.2 + R2.1 + S1.4. A 1.20\,\mathrm g iron-containing solid is dissolved to make 100.0\,\mathrm{cm^3} of solution containing only iron(II) as the reducing solute. A 20.00\,\mathrm{cm^3} portion requires 18.40\,\mathrm{cm^3} of 0.0100\,\mathrm{mol\,dm^{-3}} acidified permanganate. (a) Combine \ce{Fe^2+ -> Fe^3+ + e-} with the permanganate reduction half-equation to obtain the ionic equation. [2] (b) Calculate the mass percentage of iron in the solid. Use M(\ce{Fe})=55.85\,\mathrm{g\,mol^{-1}}. [3] (c) Explain the effect on the calculated percentage if some iron(II) is oxidized by air before titration. [1]

R3.2 + R1.4. For \ce{Zn + Cu^2+ -> Zn^2+ + Cu}, take E^\ominus(\ce{Cu^2+/Cu})=+0.34\,\mathrm V, E^\ominus(\ce{Zn^2+/Zn})=-0.76\,\mathrm V, F=96500\,\mathrm{C\,mol^{-1}}, and supplied \Delta H^\ominus=-218\,\mathrm{kJ\,mol^{-1}}. (a) Calculate E^\ominus_\mathrm{cell} and \Delta G^\ominus at 298\,\mathrm K. [2] (b) Calculate \Delta S^\ominus in \mathrm{J\,K^{-1}\,mol^{-1}}. [2] (c) Assuming constant \Delta H^\ominus and \Delta S^\ominus, predict the effect of a temperature rise on \Delta G^\ominus and on E^\ominus_\mathrm{cell}. [2]

R3.2 + S3.2 + R3.1. An unknown alcohol has formula \ce{C4H10O}. Oxidation without C–C bond cleavage gives an organic acid; a separate sample of that acid reacts with aqueous sodium carbonate. (a) State the class of alcohol and give two possible alcohol names. [2] (b) Choose one candidate. Name its aldehyde and acid, and state how the apparatus is changed to favor the acid. [2] (c) Write the acid–carbonate equation using RCOOH and explain why this final reaction is not redox. [2]