R3.2 Electron transfer reactions - IB Chemistry Question Bank

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R3.2 Electron transfer reactions

Practice exam-style IB Chemistry questions for Electron transfer reactions, aligned with the syllabus and grouped by topic.

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Question 1

SL • Paper 1A

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SL • Paper 1A

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In the reaction

$CuO(s) + H_2(g) \to Cu(s) + H_2O(l)$

the reducing agent is:

A.

$H_2(g)$

B.

$CuO(s)$

C.

$Cu(s)$

D.

$H_2O(l)$

Question 2

SL • Paper 1A

Easy

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SL • Paper 1A

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Molten lead(II) bromide, $PbBr_2(l)$ , is electrolysed using inert electrodes.

The products at the cathode and anode are:

A.

Cathode: $Pb(s)$ ; anode: $Br_2(g)$

B.

Cathode: $Br_2(g)$ ; anode: $Pb(s)$

C.

Cathode: $H_2(g)$ ; anode: $Br_2(g)$

D.

Cathode: $Pb(s)$ ; anode: $O_2(g)$

Question 3

HL • Paper 1A

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HL • Paper 1A

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The standard hydrogen electrode uses hydrogen gas and hydrogen ions under standard conditions.

The correct description of the electrode is:

A.

$H_2(g)$ at $100\ \text{kPa}$ in contact with $1.0\ \text{mol dm}^{-3}$ $H^+(aq)$ on platinum

B.

$H^+(aq)$ at $100\ \text{kPa}$ in contact with $1.0\ \text{mol dm}^{-3}$ $H_2(g)$ on platinum

C.

$H_2(g)$ at $298\ \text{kPa}$ in contact with $0.10\ \text{mol dm}^{-3}$ $H^+(aq)$ on copper

D.

$H_2(g)$ at $1.0\ \text{mol dm}^{-3}$ in contact with $100\ \text{kPa}$ $H^+(aq)$ on carbon

Question 4

SL • Paper 1A

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SL • Paper 1A

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The table shows observations when three metals are placed in aqueous solutions of metal ions.

A reaction means that the solid metal is oxidized.

The order of increasing ease of oxidation is:

solid metal	P ions	Q ions	R ions
P(s)	no reaction	reaction	reaction
Q(s)	no reaction	no reaction	reaction
R(s)	no reaction	no reaction	no reaction

A.

$R < Q < P$

B.

$P < Q < R$

C.

$R < P < Q$

D.

$Q < R < P$

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Question 5

SL • Paper 1A

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SL • Paper 1A

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A zinc-copper voltaic cell is set up using $Zn(s)\mid Zn^{2+}(aq)$ and $Cu^{2+}(aq)\mid Cu(s)$ half-cells connected by a salt bridge.

The diagram that correctly shows electron flow and salt bridge ion movement is:

A.

B.

C.

D. Question 6

SL • Paper 1A

Easy

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SL • Paper 1A

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Butan-2-ol is heated under reflux with an oxidizing agent.

The organic product is:

A.

B.

C.

D. Question 7

HL • Paper 1A

Easy

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HL • Paper 1A

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The standard electrode potentials are:

$Ag^+(aq) + e^- \rightleftharpoons Ag(s) \quad E^\circ = +0.80\ \text{V}$

$Fe^{2+}(aq) + 2e^- \rightleftharpoons Fe(s) \quad E^\circ = -0.44\ \text{V}$

The standard cell potential for the spontaneous cell reaction is:

A.

$+1.24\ \text{V}$

B.

$-0.36\ \text{V}$

C.

$+0.36\ \text{V}$

D.

$-1.24\ \text{V}$

Question 8

HL • Paper 1A

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HL • Paper 1A

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Concentrated aqueous sodium chloride is electrolysed using inert electrodes.

The main products at the cathode and anode are:

A.

Cathode: $H_2(g)$ ; anode: $Cl_2(g)$

B.

Cathode: $H_2(g)$ ; anode: $O_2(g)$

C.

Cathode: $Na(s)$ ; anode: $Cl_2(g)$

D.

Cathode: $Na(s)$ ; anode: $O_2(g)$

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Question 9

HL • Paper 1A

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HL • Paper 1A

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A steel spoon is to be electroplated with silver using an aqueous solution containing $Ag^+(aq)$ .

The correct arrangement is:

A.

B.

C.

D. Question 10

SL • Paper 2

Easy

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SL • Paper 2

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Chlorine reacts with aqueous potassium iodide according to the equation:

$Cl_2(aq) + 2I^-(aq) \to 2Cl^-(aq) + I_2(aq)$

Define an oxidizing agent in terms of electron transfer.

[1]

State the change in oxidation state of iodine and of chlorine.

[2]

Identify the reducing agent.

[1]

Question 11

SL • Paper 2

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SL • Paper 2

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Zinc reacts with dilute sulfuric acid to release a gas.

Write a balanced equation for the reaction.

[1]

Write the half-equation for the reduction of hydrogen ions.

[1]

Explain why copper does not react with dilute hydrochloric acid in the same way.

[1]

Question 12

HL • Paper 2

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HL • Paper 2

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The standard hydrogen electrode is used as a reference half-cell for measuring standard electrode potentials.

State the role of the platinum electrode.

[1]

State three standard conditions used for the standard hydrogen electrode.

[3]

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Question 13

SL • Paper 1A

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SL • Paper 1A

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Acidified manganate(VII) ions are reduced to manganese(II) ions.

The balanced reduction half-equation is:

A.

$MnO_4^-(aq) + 8H^+(aq) \to Mn^{2+}(aq) + 4H_2O(l) + 5e^-$

B.

$MnO_4^-(aq) + 8H^+(aq) + 5e^- \to Mn^{2+}(aq) + 4H_2O(l)$

C.

$Mn^{2+}(aq) + 4H_2O(l) \to MnO_4^-(aq) + 8H^+(aq) + 5e^-$

D.

$MnO_4^-(aq) + 4H^+(aq) + 3e^- \to Mn^{2+}(aq) + 2H_2O(l)$

Question 14

HL • Paper 1A

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HL • Paper 1A

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For a redox reaction, $n = 2$ and $E^\circ_{\text{cell}} = +0.50\ \text{V}$ .

Using $F = 9.65 \times 10^4\ \text{C mol}^{-1}$ , the value of $\Delta G^\circ$ is:

A.

$+96.5\ \text{kJ mol}^{-1}$

B.

$-9.65\ \text{kJ mol}^{-1}$

C.

$-96.5\ \text{kJ mol}^{-1}$

D.

$+193\ \text{kJ mol}^{-1}$

Question 15

HL • Paper 1A

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HL • Paper 1A

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The standard electrode potentials are:

$Cl_2(aq) + 2e^- \rightleftharpoons 2Cl^-(aq) \quad E^\circ = +1.36\ \text{V}$

$Br_2(aq) + 2e^- \rightleftharpoons 2Br^-(aq) \quad E^\circ = +1.07\ \text{V}$

$I_2(aq) + 2e^- \rightleftharpoons 2I^-(aq) \quad E^\circ = +0.54\ \text{V}$

The reaction predicted to be spontaneous under standard conditions is:

A.

$Cl_2(aq) + 2I^-(aq) \to 2Cl^-(aq) + I_2(aq)$

B.

$2Cl^-(aq) + I_2(aq) \to Cl_2(aq) + 2I^-(aq)$

C.

$I_2(aq) + 2Br^-(aq) \to 2I^-(aq) + Br_2(aq)$

D.

$Br_2(aq) + 2Cl^-(aq) \to 2Br^-(aq) + Cl_2(aq)$

Question 16

SL • Paper 2

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SL • Paper 2

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A student mixes aqueous halogen solutions with aqueous halide ion solutions.

Halogen	KCl(aq)	KBr(aq)	KI(aq)
chlorine	pale green	orange-brown	brown
bromine	orange-brown	orange-brown	brown
iodine	brown	brown	brown

Using the table, state the halogen produced when chlorine solution is added to aqueous potassium bromide.

[1]

Deduce the order of the three halogens from greatest to least ease of reduction.

[1]

Explain why the ease of reduction of Group 17 elements changes down the group.

[2]

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Question 17

SL • Paper 2

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SL • Paper 2

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A voltaic cell is made from a zinc half-cell and a silver half-cell under standard laboratory conditions.

Identify the anode and the cathode.

[2]

State the direction of electron flow in the external circuit.

[1]

Explain the role of the salt bridge.

[1]

Question 18

HL • Paper 2

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HL • Paper 2

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The standard reduction potentials for two half-cells are:

$Ag^+(aq) + e^- \rightleftharpoons Ag(s) \quad E^\circ = +0.80\ \text{V}$

$Sn^{2+}(aq) + 2e^- \rightleftharpoons Sn(s) \quad E^\circ = -0.14\ \text{V}$

Identify the cathode in a spontaneous cell made from these half-cells.

[1]

Calculate the standard cell potential.

[1]

Deduce the overall equation for the spontaneous reaction.

[1]

State how the sign of the standard cell potential relates to spontaneity.

[1]

Question 19

HL • Paper 2

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HL • Paper 2

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A steel key is copper-plated using an aqueous copper(II) sulfate electrolyte and a copper electrode.

State whether the steel key is the anode or the cathode.

[1]

Write the half-equation at the copper electrode.

[1]

Write the half-equation at the steel key.

[1]

Explain why the concentration of copper(II) ions remains approximately constant during plating.

[1]

Question 20

SL • Paper 1B

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SL • Paper 1B

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The reaction between sulfite ions and iodine in aqueous solution can be followed by testing samples taken at intervals. The table shows the relevant species before and after reaction.

Species	Before reaction	After reaction
sulfur-containing species	$SO_3^{2-}$	$SO_4^{2-}$
iodine-containing species	$I_2$	$I^-$

Deduce the oxidation state of sulfur in the sulfite ion and in the sulfate ion.

[2]

Identify the species that is reduced.

[1]

Explain why the sulfite ion acts as the reducing agent.

[1]

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Question 21

SL • Paper 1B

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SL • Paper 1B

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Equal masses of magnesium, zinc and copper were added separately to excess dilute hydrochloric acid. The volume of gas collected was recorded over time.

Identify the metal that is oxidized most rapidly.

[1]

Write the ionic equation for the reaction of magnesium with dilute acid.

[1]

Explain why copper gives no measurable gas under these conditions.

[2]

Question 22

SL • Paper 2

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SL • Paper 2

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Nitrite ions, $NO_2^-(aq)$ , can be oxidized to nitrate ions, $NO_3^-(aq)$ , in acidic solution.

State whether electrons appear on the left or right in an oxidation half-equation.

[1]

Deduce the balanced half-equation for this oxidation in acidic solution.

[2]

Outline why acidified manganate(VII) titrations can be described as self-indicating.

[1]

Question 23

SL • Paper 2

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SL • Paper 2

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Organic functional groups can undergo oxidation or reduction.

Write an equation for the complete oxidation of pentan-1-ol to pentanoic acid using $[O]$ .

[1]

Explain why distillation is used when preparing an aldehyde from a primary alcohol.

[2]

State the organic product formed when pent-2-ene reacts with hydrogen in a hydrogenation reaction.

[1]

Question 24

HL • Paper 2

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HL • Paper 2

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A standard electrochemical cell has $E^\circ_{\text{cell}} = +1.02\ \text{V}$ . The balanced cell reaction transfers two moles of electrons per mole of reaction. Use $F = 9.65 \times 10^4\ \text{C mol}^{-1}$ .

State the value of $n$ for this cell reaction.

[1]

Calculate the standard change in Gibbs energy, $\Delta G^\circ$ , in $\text{kJ mol}^{-1}$ .

[2]

Explain whether this value of $\Delta G^\circ$ is consistent with a spontaneous cell reaction.

[1]

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Question 25

HL • Paper 2

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HL • Paper 2

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Selected standard reduction potentials are shown.

$Mg^{2+}(aq) + 2e^- \rightleftharpoons Mg(s) \quad E^\circ = -2.37\ \text{V}$

$Cu^{2+}(aq) + 2e^- \rightleftharpoons Cu(s) \quad E^\circ = +0.34\ \text{V}$

$I_2(aq) + 2e^- \rightleftharpoons 2I^-(aq) \quad E^\circ = +0.54\ \text{V}$

Identify the strongest oxidizing agent from the species in the reduction half-equations.

[1]

Identify the strongest reducing agent from the reduced forms shown.

[1]

Predict whether magnesium will react spontaneously with iodine to form magnesium ions and iodide ions. Support your answer with a calculation.

[2]

Question 26

HL • Paper 2

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HL • Paper 2

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Concentrated aqueous sodium chloride is electrolysed using inert electrodes.

Write the half-equation for the main reaction at the cathode.

[1]

Write the half-equation for the main reaction at the anode.

[1]

Explain why sodium metal is not produced at the cathode.

[1]

State the change in pH expected near the cathode.

[1]

Question 27

SL • Paper 1B

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SL • Paper 1B

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A student determined the concentration of iron(II) ions in a solution by titration with acidified potassium manganate(VII), $KMnO_4(aq)$ . The half-equation for manganate(VII) in acid is shown with the titration data.

Quantity	Value / units
$KMnO_4$ concentration	0.0200 mol dm $^{-3}$
$\text{Fe}^{2+}$ sample volume	25.00 cm $^3$
$MnO_4^-$ : $\text{Fe}^{2+}$ ratio	1 : 5
Titre 1	18.55 cm $^3$
Titre 2	18.65 cm $^3$
Average titre	18.60 cm $^3$

State the colour change at the endpoint in this self-indicating titration.

[1]

Calculate the concentration of $Fe^{2+}(aq)$ in the sample. The average titre is $18.60\ \text{cm}^3$ and the mole ratio $MnO_4^-:Fe^{2+}$ is $1:5$ .

[3]

Question 28

SL • Paper 1B

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SL • Paper 1B

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Small pieces of different metals were placed in aqueous solutions containing metal ions. The observations are summarized in the table.

metal added	$\text{Ag}^+(aq)$	$\text{Cu}^{2+}(aq)$	$\text{Mg}^{2+}(aq)$	$\text{Zn}^{2+}(aq)$
$\text{Ag}(s)$	no visible reaction	no visible reaction	no visible reaction	no visible reaction
$\text{Cu}(s)$	reaction	no visible reaction	no visible reaction	no visible reaction
$\text{Mg}(s)$	reaction	reaction	no visible reaction	reaction
$\text{Zn}(s)$	reaction	reaction	no visible reaction	no visible reaction

Deduce the order of decreasing ease of oxidation of the metals.

[2]

Write an ionic equation for the reaction observed when zinc is added to copper(II) sulfate solution.

[1]

Explain why no visible reaction occurs when copper is added to zinc sulfate solution.

[2]

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Question 29

SL • Paper 1B

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SL • Paper 1B

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A zinc-copper voltaic cell was assembled using $1.0\ \text{mol dm}^{-3}$ solutions of $ZnSO_4(aq)$ and $CuSO_4(aq)$ connected by a salt bridge.

Identify the anode and the cathode.

[2]

State the direction of electron flow in the external circuit.

[1]

Explain the direction of anion movement in the salt bridge.

[1]

Question 30

SL • Paper 1B

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SL • Paper 1B

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Three alcohols were warmed separately with an oxidizing agent under the same conditions. The observations and structures are shown.

Alcohol	Structural formula	Observation on warming with acidified potassium dichromate(VI)
A	CH $_3$ CH $_2$ CH $_2$ OH	Orange solution turns green
B	CH $_3$ CH(OH)CH $_3$	Orange solution turns green
C	(CH $_3$ ) $_3$ COH	No visible change; solution remains orange

Deduce the type of organic product formed when alcohol B is oxidized.

[1]

State the product type obtained from alcohol A when the reaction mixture is heated under reflux with excess oxidizing agent.

[1]

Explain why distillation is used instead of reflux when an aldehyde is required from alcohol A.

[1]

Explain the observation for alcohol C.

[1]

Question 31

HL • Paper 1B

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HL • Paper 1B

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Several metal/metal ion half-cells were connected separately to a standard hydrogen electrode under standard conditions. The measured standard electrode potentials are shown as reduction potentials.

Half-equation	E $^\circ$ / V
$\text{Ag}^+(aq) + e^- \rightleftharpoons \text{Ag}(s)$	0.80
$\text{Cu}^{2+}(aq) + 2e^- \rightleftharpoons \text{Cu}(s)$	0.34
$2\text{H}^+(aq) + 2e^- \rightleftharpoons \text{H}_2(g)$	0.00
$\text{Fe}^{2+}(aq) + 2e^- \rightleftharpoons \text{Fe}(s)$	-0.44
$\text{Zn}^{2+}(aq) + 2e^- \rightleftharpoons \text{Zn}(s)$	-0.76

State why platinum is used in the standard hydrogen electrode.

[1]

State two standard conditions used for measuring standard electrode potentials.

[2]

Identify the strongest oxidizing agent among the species listed in the table.

[1]

Question 32

HL • Paper 1B

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HL • Paper 1B

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A steel spoon was silver-plated using an electrolytic cell. The mass changes of the two electrodes during electrolysis are shown.

Identify which electrode is the cathode.

[1]

Write the half-equation for silver deposition on the spoon.

[1]

Write the half-equation occurring at the silver strip.

[1]

Explain why using a silver anode helps maintain the concentration of silver ions in the electrolyte.

[1]

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Question 33

HL • Paper 1B

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HL • Paper 1B

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A student considered constructing a standard cell from nickel and silver half-cells. The relevant standard reduction potentials are shown.

Half-cell	Reduction half-equation	E $^\circ$ / V
Silver	$\text{Ag}^+(aq) + e^- \to \text{Ag}(s)$	+0.80
Nickel	$\text{Ni}^{2+}(aq) + 2e^- \to \text{Ni}(s)$	−0.25

Identify the cathode in the spontaneous cell.

[1]

Calculate $E^\circ_{\text{cell}}$ using $E^\circ(Ag^+/Ag)=+0.80\ \text{V}$ and $E^\circ(Ni^{2+}/Ni)=-0.25\ \text{V}$ .

[2]

Write the overall equation for the spontaneous reaction.

[1]

State whether the forward reaction in part (c) is spontaneous under standard conditions.

[1]

Question 34

HL • Paper 1B

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HL • Paper 1B

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Iron(II) ions can reduce dichromate(VI) ions in acidic solution. The relevant standard reduction potentials are shown.

Half-equation (reduction)	E $^\circ$ / V
$\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^+ + 6e^- \to 2\text{Cr}^{3+} + 7\text{H}_2\text{O}$	+1.33
$\text{Fe}^{3+} + e^- \to \text{Fe}^{2+}$	+0.77

State the number of moles of electrons transferred per mole of balanced reaction.

[1]

Calculate $\Delta G^\circ$ for the reaction using $E^\circ_{\text{cell}}=+0.56\ \text{V}$ and $F=9.65\times10^4\ \text{C mol}^{-1}$ .

[2]

Interpret the sign of $\Delta G^\circ$ for the reaction under standard conditions.

[1]

Question 35

HL • Paper 1B

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HL • Paper 1B

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Aqueous copper(II) sulfate was electrolysed using two different electrode arrangements. The same current was passed for the same time in each experiment.

Electrodes	Cathode observation	Anode observation	Blue intensity / a.u. (before $\to$ after)
Inert electrodes	Reddish-brown solid forms; cathode gains 0.32 g	Colourless gas bubbles; relights a glowing splint	8.0 $\to$ 6.4
Copper electrodes	Reddish-brown solid forms; cathode gains 0.32 g	Anode gets smaller; loses 0.32 g	8.0 $\to$ 7.9

Write the cathode half-equation common to both experiments.

[1]

Identify the gas formed at the inert anode.

[1]

Write the anode half-equation when copper electrodes are used.

[1]

Explain why the blue colour remains approximately constant when copper electrodes are used.

[1]

Question 36

SL • Paper 2

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SL • Paper 2

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Vanadium forms several oxo-ions in acidic aqueous solution. Two reactions involving vanadium species are shown.

Reaction 1: $VO_2^+(aq) + 2H^+(aq) + e^- \to VO^{2+}(aq) + H_2O(l)$

Reaction 2: $Zn(s) + 2VO_2^+(aq) + 4H^+(aq) \to Zn^{2+}(aq) + 2VO^{2+}(aq) + 2H_2O(l)$

For reaction 1:

I.

Deduce the oxidation state of vanadium in $VO_2^+$ and in $VO^{2+}$ .

[2]

II.

State whether vanadium is oxidized or reduced.

[1]

For reaction 2:

I.

Identify the reducing agent and the oxidizing agent.

[2]

II.

Explain, using electron transfer, why the two agents identified in (b)(i) have these roles.

[2]

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Question 37

SL • Paper 2

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SL • Paper 2

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A student tests three metals, $A$ , $B$ and $C$ , with aqueous solutions containing their ions. A tick indicates that a visible displacement reaction occurs.

solid metal	$A^{2+}$ (aq)	$B^{2+}$ (aq)	$C^{2+}$ (aq)
$A$ (s)	–	✓	✓
$B$ (s)	–	–	✓
$C$ (s)	–	–	–

Use the data to deduce the order of ease of oxidation of the metals, from greatest to least.

[2]

For the reaction between solid $A$ and $B^{2+}(aq)$ :

I.

Write the oxidation half-equation, assuming $A$ forms $A^{2+}$ .

[1]

II.

Write the reduction half-equation for $B^{2+}$ .

[1]

III.

Deduce the overall ionic equation.

[1]

Explain the trend in ease of reduction of halogens down Group 17, in terms of atomic structure.

[2]

Question 38

SL • Paper 2

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SL • Paper 2

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A student adds excess dilute hydrochloric acid separately to equal masses of magnesium, zinc and copper. Only two of the metals produce a gas.

For the reaction of zinc with dilute hydrochloric acid:

I.

Write a balanced equation, including state symbols.

[2]

II.

State a test for the gas produced and the positive result.

[1]

Explain why copper does not react with dilute hydrochloric acid to produce hydrogen.

[2]

Identify the reducing agent in the reaction of magnesium with dilute acid and justify your answer.

[1]

Question 39

HL • Paper 1B

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HL • Paper 1B

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Aqueous sodium chloride was electrolysed using inert graphite electrodes. The products depended on the concentration of sodium chloride. Selected standard electrode potential data and observations are shown.

Half-equation	E $^\circ$ / V	Condition	Observation
$\text{Na}^+(aq) + e^- \to \text{Na}(s)$	-2.71	cathode	—
$2H_2O(l) + 2e^- \to H_2(g) + 2OH^-(aq)$	-0.83	cathode	colourless gas; squeaky pop
$Cl_2(g) + 2e^- \to 2Cl^-(aq)$	+1.36	anode, concentrated $\text{NaCl}(aq)$	pungent gas; bleaches damp litmus
$O_2(g) + 2H_2O(l) + 4e^- \to 4OH^-(aq)$	+0.40	anode, dilute $\text{NaCl}(aq)$	colourless gas; relights glowing splint

Deduce the cathode product in both concentrated and dilute aqueous sodium chloride.

[1]

Write the cathode half-equation.

[1]

Deduce the main anode product for concentrated aqueous sodium chloride.

[1]

Explain why sodium metal is not produced at the cathode.

[2]

Question 40

SL • Paper 2

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SL • Paper 2

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Acidified manganate(VII) ions, $MnO_4^-$ , oxidize ethanedioate ions, $C_2O_4^{2-}$ , to carbon dioxide. Manganese is reduced to $Mn^{2+}$ .

Deduce balanced half-equations for the reaction in acidic solution.

I.

Deduce the reduction half-equation for $MnO_4^-$ to $Mn^{2+}$ .

[3]

II.

Deduce the oxidation half-equation for $C_2O_4^{2-}$ to $CO_2$ .

[2]

Deduce the overall ionic equation for the reaction.

[2]

Explain why this titration is self-indicating.

[1]

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Question 41

SL • Paper 2

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SL • Paper 2

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A primary cell is constructed using a zinc half-cell and a silver half-cell. The cell diagram is:

$Zn(s)\mid Zn^{2+}(aq)\mid\mid Ag^+(aq)\mid Ag(s)$

For this primary cell:

I.

Identify the anode and cathode.

[2]

II.

Write the half-equation occurring at each electrode.

[2]

Explain the direction of electron flow in the external circuit.

[2]

Explain the movement of ions in the salt bridge during operation of the cell.

[2]

Question 42

SL • Paper 2

Hard

Calculator Permitted

SL • Paper 2

Hard

Calculator Permitted

Butan-1-ol, butan-2-ol and 2-methylpropan-2-ol are separately heated with an oxidizing agent. Organic equations may use $[O]$ .

For butan-1-ol:

I.

Write an equation for its oxidation to an aldehyde.

[1]

II.

Write an equation for further oxidation of the aldehyde.

[1]

III.

State the experimental technique used to obtain the aldehyde as the main product.

[1]

Explain why reflux is used when converting butan-1-ol to butanoic acid.

[2]

Compare the oxidation of butan-2-ol and 2-methylpropan-2-ol under similar conditions.

[2]

Question 43

HL • Paper 2

Hard

Calculator Permitted

HL • Paper 2

Hard

Calculator Permitted

The standard reduction potentials for three half-cells are shown.

$Fe^{3+}(aq) + e^- \rightleftharpoons Fe^{2+}(aq) \quad E^\circ = +0.77\ \text{V}$

$I_2(aq) + 2e^- \rightleftharpoons 2I^-(aq) \quad E^\circ = +0.54\ \text{V}$

$Sn^{2+}(aq) + 2e^- \rightleftharpoons Sn(s) \quad E^\circ = -0.14\ \text{V}$

Using the standard electrode potential data:

I.

Identify the strongest oxidizing agent from the species shown on the left of the half-equations.

[1]

II.

Identify the strongest reducing agent from the species shown on the right of the half-equations.

[1]

III.

Explain your answers to (a)(i) and (a)(ii).

[1]

Predict whether $I_2(aq)$ will oxidize $Sn(s)$ under standard conditions. Justify your answer quantitatively.

[3]

Discuss why a platinum electrode is used in the standard hydrogen electrode.

[2]

Question 44

HL • Paper 2

Hard

Calculator Permitted

HL • Paper 2

Hard

Calculator Permitted

A student proposes the following cell under standard conditions.

$Mg(s)\mid Mg^{2+}(aq)\mid\mid Cu^{2+}(aq)\mid Cu(s)$

$Mg^{2+}(aq) + 2e^- \rightleftharpoons Mg(s) \quad E^\circ = -2.37\ \text{V}$

$Cu^{2+}(aq) + 2e^- \rightleftharpoons Cu(s) \quad E^\circ = +0.34\ \text{V}$

For the proposed cell:

I.

Identify the cathode and write the cathode half-equation.

[2]

II.

Calculate $E^\circ_{\text{cell}}$ .

[2]

Explain why the reaction is spontaneous in the direction shown by the cell diagram.

[2]

The student multiplies the copper half-equation by 2 and states that $E^\circ$ for the copper half-cell becomes $+0.68\ \text{V}$ . Evaluate this statement.

[2]

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Question 45

HL • Paper 2

Hard

Calculator Permitted

HL • Paper 2

Hard

Calculator Permitted

The standard cell potential for the reaction below is $+1.10\ \text{V}$ .

$Zn(s) + Cu^{2+}(aq) \to Zn^{2+}(aq) + Cu(s)$

Use $F = 9.65 \times 10^4\ \text{C mol}^{-1}$ .

For this reaction:

I.

Deduce the value of $n$ in $\Delta G^\circ = -nFE^\circ_{\text{cell}}$ .

[1]

II.

Calculate $\Delta G^\circ$ in $\text{kJ mol}^{-1}$ .

[3]

Explain the sign of $\Delta G^\circ$ in relation to the cell potential and spontaneity.

[2]

State why the unit calculation is consistent with energy per mole.

[1]

Question 46

HL • Paper 2

Hard

Calculator Permitted

HL • Paper 2

Hard

Calculator Permitted

A steel key is to be electroplated with copper using aqueous copper(II) sulfate. A copper strip is used as one electrode.

For the electroplating cell:

I.

Identify which electrode is the cathode and which is the anode.

[2]

II.

Write the half-equation for the reaction at the key.

[1]

III.

Write the half-equation for the reaction at the copper strip.

[1]

Explain why the concentration of $Cu^{2+}(aq)$ remains approximately constant during electroplating.

[2]

Suggest one visible observation at each electrode during the process.

[1]

Question 47

HL • Paper 2

Hard

Calculator Permitted

HL • Paper 2

Hard

Calculator Permitted

Concentrated aqueous sodium chloride and aqueous copper(II) sulfate are electrolysed using inert graphite electrodes. Standard electrode potentials for relevant competing reductions are available in the data booklet.

For concentrated aqueous sodium chloride:

I.

Deduce the main product at the cathode and write the cathode half-equation.

[2]

II.

Deduce the main product at the anode and write the anode half-equation.

[2]

Explain why electrolysis of molten sodium chloride gives a different cathode product from aqueous sodium chloride.

[2]

For aqueous copper(II) sulfate with inert electrodes:

I.

Deduce the product at the cathode.

[1]

II.

Deduce the product at the anode.

[1]

Question 48

HL • Paper 2

Hard

Calculator Permitted

HL • Paper 2

Hard

Calculator Permitted

A simplified secondary lithium cell has the following discharge half-equations.

Anode during discharge: $Li(s) \to Li^+(aq) + e^-$

Cathode during discharge: $Li^+(aq) + CoO_2(s) + e^- \to LiCoO_2(s)$

For discharge of the cell:

I.

Deduce the overall equation.

[1]

II.

State the direction of electron flow in the external circuit during discharge.

[1]

III.

Identify the energy conversion during discharge.

[1]

For charging of the cell:

I.

Deduce the half-equation at the electrode where lithium metal is formed.

[1]

II.

Deduce the half-equation at the other electrode during charging.

[1]

III.

Explain why charging is described as an electrolytic process.

[1]

Evaluate one advantage and one disadvantage of using a secondary cell rather than a primary cell for powering a portable computer.

[2]

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R3.1 Proton transfer reactions

R3.3 Hydrogen sharing reactions

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