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;

[Maximum number: 1]

An aqueous solution X reacts with a solid Y, to produce a flammable gas. Which of the following suggestions could substances X and Y be?

nitric acid, $\mathrm{HNO}_{3}$

calcium carbonate, $\mathrm{CaCO}_{3}$

sulfuric acid, $\mathrm{H}_{2} \mathrm{SO}_{4}$

zinc, Zn

hydrochloric acid, HCl

copper, Cu

sodium hydroxide solution, NaOH

aluminum oxide, $\mathrm{Al}_{2} \mathrm{O}_{3}$

[Maximum number: 4]

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

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

[Maximum number: 2]

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

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

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

[Maximum number: 1]

Magnesium is a group 2 metal which exists as a number of isotopes and forms many compounds.

Magnesium chloride can be electrolysed.

Identify the type of reaction occurring at the cathode (negative electrode).

[Maximum number: 1]

As well as being burnt, methanol can also be used to provide electricity through a fuel cell. A schematic diagram of such a fuel cell, that depends on the transfer of hydrogen ions between the electrodes, is shown below.

Even though fuel cells, primary cells and rechargeable cells have similar fundamental characteristics, there are important differences between them.

Suggest a way in which they are similar.

[Maximum number: 1]

Even though fuel cells, primary cells and rechargeable cells have similar fundamental characteristics, there are important differences between them.

Outline the difference between primary and rechargeable cells.

[Maximum number: 2]

Magnesium is a group 2 metal which exists as a number of isotopes and forms many compounds.

Magnesium chloride can be electrolysed.

Deduce the half-equations for the reactions at each electrode when molten magnesium chloride is electrolysed, showing the state symbols of the products. The melting points of magnesium and magnesium chloride are 922 K and 987 K respectively.

Anode (positive electrode):

Cathode (negative electrode):

[Maximum number: 2]

Phenylethene (styrene) is produced from ethylbenzene in a gas-phase equilibrium.

Figure for Question 10 — IB Chemistry HL

A compound used to make polymers, benzene-1,4-dicarboxylic acid, can be produced by a series of reactions.

Suggest the conditions required and the role of the reagent $\mathrm{KMnO}_{4}$ used to convert intermediate B into intermediate C.

Conditions:

Role of $\mathrm{KMnO}_{4}$ :

[Maximum number: 2]

Ethanol is obtained by the hydration of ethene, $\mathrm{C}_{2} \mathrm{H}_{4}$ .

Alternative synthetic routes exist to produce alcohols.

Propan-2-ol can also be formed in one step from a compound containing a carbonyl group.

State the name of this compound and the type of reaction that occurs.

Name of carbonyl compound:

Type of reaction:

[Maximum number: 1]

Alkenes, such as A (shown below), are important intermediates in the petrochemical industry because they undergo addition reactions to produce a wide variety of products, such as the conversion shown below.

Figure for Question 12 — IB Chemistry HL

In the gas phase, A reacts with hydrogen to form D :

State the conditions required for this reaction to occur.

$\quad \mathrm{Ni} / \mathrm{Pt} / \mathrm{Pd}$ catalyst;

[Maximum number: 2]

Magnesium is a group 2 metal which exists as a number of isotopes and forms many compounds.

Standard electrode potentials are measured relative to the standard hydrogen electrode. Describe a standard hydrogen electrode.

Any two of:

«inert» Pt electrode

OR

platinum black conductor

$1 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{H}^{+}(\mathrm{aq})$ $\mathrm{H}_{2}(\mathrm{~g})$ at 100 kPa

[Maximum number: 1]

Magnesium is a group 2 metal which exists as a number of isotopes and forms many compounds.

A magnesium half-cell, $\mathrm{Mg}(\mathrm{s}) / \mathrm{Mg}^{2+}(\mathrm{aq})$ , can be connected to a copper half-cell, $\mathrm{Cu}(\mathrm{s}) / \mathrm{Cu}^{2+}(\mathrm{aq})$ .

Determine the standard cell potential, in V, for the cell. Refer to section 24 of the data booklet.

$<+0.34 \mathrm{~V}-(-2.37 \mathrm{~V})=+>2.71$ «V»

[Maximum number: 2]

A standard hydrogen electrode is connected to a copper-copper(II) nitrate half-cell.

Calculate the standard Gibbs free energy of the cell, in $\mathrm{kJ} \mathrm{mol}^{-1}$ . Use sections 1, 2 and 24 of the data booklet.

$\begin{aligned} \mathrm{n}=2 \text { « }-2(96500)(0.34)=»-65620 \text { «J mol }{ }^{-1} » /-65.6 \text { « } \mathrm{kJ} \mathrm{~mol}^{-1} » \vee \end{aligned}$

Answer must be negative

[Maximum number: 2]

Magnesium is a group 2 metal which exists as a number of isotopes and forms many compounds.

Magnesium chloride can be electrolysed.

State the products when a very dilute aqueous solution of magnesium chloride is electrolysed.

Anode (positive electrode):

Cathode (negative electrode):

Anode (positive electrode):

oxygen/ $\ce{O2}$

OR

hydogen ion/proton/ H^+AND oxygen/ O_2

Cathode (negative electrode):

hydrogen/ H_2

OR

hydroxide «ion»/ $\ce{OH}$ ^-AND hydrogen/ H_2

Award [1 max] if correct products given at wrong electrodes.

[Maximum number: 3]

Iron rusts in the presence of oxygen and water. Rusting is a redox process involving several steps that produces hydrated iron(III) oxide, $\mathrm{Fe}_{2} \mathrm{O}_{3} \bullet \mathrm{nH}_{2} \mathrm{O}$ , as the final product. The half-equations involved for the first step of rusting are given below.

Half-equation 1: $\quad \mathrm{Fe}(\mathrm{s}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-}$

Half-equation 2: $\quad \mathrm{O}_{2}(\mathrm{aq})+4 \mathrm{e}^{-}+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 4 \mathrm{OH}^{-}(\mathrm{aq})$

Describe how electrolysis can be used to electroplate a bracelet with a layer of silver metal. Include the choice of electrodes and electrolyte needed in your description.

bracelet/object to be electroplated is the cathode/negative electrode;
silver anode/positive electrode;
electrolyte: liquid $\mathrm{Na}[\mathrm{Ag}(\mathrm{CN})_{2}]$ / sodium dicyanoargentate / $\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}$ / solution of an appropriate silver salt / $\mathrm{AgNO}_{3}$ / silver nitrate;

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)

Question 3

Question 4

Question 5

Question 6

Question 7

Question 8

Question 9

Question 10

Question 11

Question 12

Question 13

Question 14

Question 15

Question 16

Question 17