IB Chemistry HL 3.2 Electron Transfer Reactions Questions

Use IB Chemistry HL questions to assign oxidation states, balance redox equations and interpret cells, electrolysis or organic redox products.

Exam points

Assign oxidation states, identify oxidation and reduction, and identify the oxidizing and reducing agents in a stated reaction.
Write and balance oxidation and reduction half-equations, then combine them into a balanced redox equation with conserved atoms, charge and electrons.
Use displacement observations and metal reactivity to deduce an activity order, predict reactions with acids or salts and identify products such as hydrogen.
Assign anode and cathode processes, electron flow and polarity in voltaic or electrolytic cells, keeping oxidation at the anode and reduction at the cathode.
Draw or interpret voltaic and secondary-cell arrangements, including electrodes, electrolytes, salt bridges, ion movement, discharge and recharge.
Deduce molten-electrolyte products, half-equations and current carriers in electrolysis, including the positive anode and negative cathode.
Predict alcohol oxidation, carbonyl reduction and catalytic hydrogenation products from functional groups, reagents, conditions and structural changes.
Use the standard hydrogen electrode and tabulated reduction potentials to compare oxidizing or reducing strength and calculate or interpret half-cell potentials.
Calculate E°cell = E°cathode − E°anode, use its sign to predict standard-state spontaneity and relate it to the feasible overall redox equation.
Calculate ΔG° = −nFE°cell with the balanced electron number and consistent units, and connect negative ΔG° with spontaneous redox change.
Predict competing aqueous electrolysis products from reduction potentials, concentration and electrode material, then justify the electrode half-equations.
Explain electroplating by placing the object at the cathode and maintaining metal ions with a suitable anode and electrolyte.

Covered syllabus points

3.2.1—Oxidation and reduction 3.2.2—Half-equations 3.2.3—Periodic trends in redox 3.2.4—Acids with reactive metals 3.2.5—Electrodes in electrochemical cells 3.2.6—Primary (voltaic) cells 3.2.7—Secondary (rechargeable) cells 3.2.8—Electrolytic cells 3.2.9—Oxidation of organic functional groups 3.2.10—Reduction of organic functional groups 3.2.11—Reduction of unsaturated compounds 3.2.12 (HL)—Standard electrode potential (E⦵)3.2.13 (HL)—Standard cell potential (E⦵cell)3.2.14 (HL)—Gibbs energy and cell potential 3.2.15 (HL)—Electrolysis of aqueous solutions 3.2.16 (HL)—Electroplating

[Maximum number: 3]

Phosphine (IUPAC name phosphane) is a hydride of phosphorus, with the formula $\mathrm{PH}_{3}$ .

Phosphine is usually prepared by heating white phosphorus, one of the allotropes of phosphorus, with concentrated aqueous sodium hydroxide. The equation for the reaction is:

\mathrm{P}_{4}(\mathrm{~s})+3 \mathrm{OH}^{-}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{PH}_{3}(\mathrm{~g})+3 \mathrm{H}_{2} \mathrm{PO}_{2}^{-}(\mathrm{aq})

Oxidation is now defined in terms of change of oxidation number. Explore how earlier definitions of oxidation and reduction may have led to conflicting answers for the conversion of $\mathrm{P}_{4}$ to $\mathrm{H}_{2} \mathrm{PO}_{2}^{-}$ and the way in which the use of oxidation numbers has resolved this.

[Maximum number: 11]

Bleaches in which chlorine is the active ingredient are the most common, although some environmental groups have concerns about their use.

(a)

Outline, with the help of a chemical equation, why this reaction occurs.

[ 2 ]

chlorine more reactive/more powerful oxidizing agent (than bromine);
$\mathrm{Cl}_{2}(aq)+2\mathrm{Br}^{-}(aq)\rightarrow \mathrm{Br}_{2}(aq)+2\mathrm{Cl}^{-}(aq)$ ;

(b)

Aqueous sodium chlorate(I), NaOCl, the most common active ingredient in chlorine based bleaches, oxidizes coloured materials to colourless products while being reduced to the chloride ion. It will also oxidize sulfur dioxide to the sulfate ion.

[ 6 ]

(i)

Deduce a balanced equation for the reaction between the chlorate(I) ion and sulfur dioxide from the appropriate half-equations.

[ 4 ]

$\mathrm{ClO}^{-}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{H}_{2}\mathrm{O}(\mathrm{l})+\mathrm{Cl}^{-}(\mathrm{aq})$ ;
$\mathrm{SO}_{4}^{2-}(\mathrm{aq})+4 \mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{SO}_{2}(\mathrm{aq})+2 \mathrm{H}_{2}\mathrm{O}(\mathrm{l})$ ;
final equation: $\mathrm{ClO}^{-}(\mathrm{aq})+\mathrm{SO}_{2}(\mathrm{aq})+\mathrm{H}_{2}\mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{SO}_{4}^{2-}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})$ ;

(ii)

State the initial and final oxidation numbers of both chlorine and sulfur in the final equation.

Table for Question (ii) — IB Chemistry HL

[ 2 ]

Element	Initial oxidation number	Final oxidation number
Chlorine	+I / +1	-I / -1
Sulfur	+IV / +4	+VI / +6

(c)

The standard electrode potential for the reduction of the chlorate(V) ion to the chloride ion is +1.49 V .

[ 3 ]

(i)

Define the term standard electrode potential.

[ 1 ]

(ii)

Referring to Table 14 of the Data Booklet, deduce, giving a reason, whether the oxidation of the chromium(III) ion to the dichromate(VI) ion by the chlorate(V) ion is energetically feasible.

[ 2 ]

yes / energetically feasible;
would have a positive $E_{\text{cell}}$ / chlorate(V) ion stronger oxidizing agent than dichromate(VI) ion;

All question bank results loaded

Covered syllabus points

Question 1

Question 2

Question (a)

Question (b)

Question (i)

Question (ii)

Question (c)

Question (i)

Question (ii)