IB Chemistry SL 3.2 Electron Transfer Reactions Questions

Use IB Chemistry SL 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.

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

[Maximum number: 1]

Chlorine undergoes many reactions.

2.67 g of manganese(IV) oxide was added to $200.0 \mathrm{~cm}^{3}$ of $2.00 \mathrm{moldm} \mathrm{m}^{-3} \mathrm{HCl}$ .

\mathrm{MnO}_{2}(\mathrm{~s})+4 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{Cl}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\mathrm{MnCl}_{2}(\mathrm{aq})

Deduce, referring to oxidation states, whether $\mathrm{MnO}_{2}$ is an oxidizing or reducing agent.

oxidizing agent AND oxidation state of Mn changes from +4 to +2 / decreases $\boldsymbol{}$

[Maximum number: 4]

(a)

Ethanol is a primary alcohol that can be oxidized by acidified potassium dichromate(VI). Distinguish between the reaction conditions needed to produce ethanal and ethanoic acid.

Ethanal:
Ethanoic acid:

[ 2 ]

(b)

Deduce the half-equation for the oxidation of ethanol to ethanal.

[ 1 ]

$\mathrm{CH_3CH_2OH \rightarrow CH_3CHO + 2H^+ + 2e^-}$

Half-equation required. Do not accept:
$\mathrm{C_2H_5OH+2[O]\rightarrow CH_3CHO+H_2O}$

Accept e for $e^-$ .

(c)

Deduce the overall redox equation for the reaction of ethanol to ethanal with acidified potassium dichromate(VI) by combining your answer to part (c) (iii) with the following half-equation:

\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq})+14 \mathrm{H}^{+}(\mathrm{aq})+6 \mathrm{e}^{-} \rightarrow 2 \mathrm{Cr}^{3+}(\mathrm{aq})+7 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})

[ 1 ]

$\mathrm{3CH_3CH_2OH(aq)+Cr_2O_7^{2-}(aq)+8H^+(aq) \rightarrow 2Cr^{3+}(aq)+3CH_3CHO(l)+7H_2O(l)}$

Award [1] for correct reactants and products.
Award [1] for correct balancing;
M2 can only be scored if M1 is correct.
Ignore state symbols.

[Maximum number: 7]

(a)

Consider the following spontaneous reactions.

\begin{aligned} & \mathrm{Fe}(\mathrm{~s})+\mathrm{Ni}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ni}(\mathrm{~s}) \\ & \mathrm{Zn}(\mathrm{~s})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Fe}(\mathrm{~s}) \\ & \mathrm{Ni}(\mathrm{~s})+\mathrm{Pb}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Ni}^{2+}(\mathrm{aq})+\mathrm{Pb}(\mathrm{~s}) \end{aligned}

[ 3 ]

(i)

Deduce the order of increasing reactivity of the metals based on the reactions above.

[ 2 ]

(ii)

Identify the strongest oxidizing agent in the reactions above.

[ 1 ]

$\quad \mathrm{Pb}^{2+} /$ lead(II) (ions);

Marking guidance:

Do not accept Pb/lead.

(b)

Draw a diagram of an electrolytic cell for the electrolysis of molten nickel(II) bromide, $\mathrm{NiBr}_{2}(\mathrm{l})$ . Include the direction of the electron flow, the polarity of electrodes and the half-equations at each electrode.

Negative electrode (cathode):

Positive electrode (anode):

[ 4 ]

Figure for Answer to Question (b) — IB Chemistry SL

power source and direction of $\mathrm{e}^{-}$ movement;
labelled +/positive electrode/anode and -/negative electrode/cathode and (molten) electrolyte/ $\mathrm{NiBr}_{2}(\mathrm{l})$ ;
Accept polarity of electrodes given at the power source.

Award [1 max] for M3 and M4 if equations are given at wrong electrodes. Ignore state symbols and reversible sign.
Allow e instead of $e^{-}$ .

[Maximum number: 2]

Consider the following spontaneous reactions.

\begin{aligned} & \mathrm{Fe}(\mathrm{~s})+\mathrm{Ni}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ni}(\mathrm{~s}) \\ & \mathrm{Zn}(\mathrm{~s})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Fe}(\mathrm{~s}) \\ & \mathrm{Ni}(\mathrm{~s})+\mathrm{Pb}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Ni}^{2+}(\mathrm{aq})+\mathrm{Pb}(\mathrm{~s}) \end{aligned}

Deduce the order of increasing reactivity of the metals based on the reactions above.

[Maximum number: 2]

Two hydrides of nitrogen are ammonia and hydrazine, $\mathrm{N}_{2} \mathrm{H}_{4}$ . One derivative of ammonia is methanamine whose molecular structure is shown below.

Outline, using an ionic equation, what is observed when magnesium powder is added to a solution of ammonium chloride.

bubbles

OR

gas

OR

magnesium disappears
$2 \mathrm{NH}_{4}^{+}(\mathrm{aq})+\mathrm{Mg}(\mathrm{~s}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+2 \mathrm{NH}_{3}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})$

Marking guidance:

Do not accept "hydrogen" without reference to observed changes. Accept "smell of ammonia".

Accept $2 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Mg}(\mathrm{s}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq}) +\mathrm{H}_{2}(\mathrm{~g})$

Equation must be ionic.

[Maximum number: 2]

Iron(II) sulfide can be produced by heating powdered iron and sulfur together.

A student hypothesised that iron(II) sulfide is ionic and therefore can be converted back to iron and sulfur by electrolysis of the molten salt, using inert electrodes.

Annotate the electrolytic cell with the terms anode and cathode, and show the direction of ion movement.

[Maximum number: 1]

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

A voltaic cell is made from a half-cell containing a magnesium electrode in a solution of magnesium nitrate and a half-cell containing a silver electrode in a solution of silver(I) nitrate.

Negative electrode (anode):

Positive electrode (cathode):
(ii) Outline one function of the salt bridge.

[Maximum number: 1]

Which statement about electrochemical cells is correct?

Primary (voltaic) cells use electrical energy to produce a redox reaction.

Electrolytic cells are rechargeable.

Secondary cells can use electrical energy to produce a redox reaction.

Primary (voltaic) cells are rechargeable.

[Maximum number: 4]

Figure for Answer to Question 9 — IB Chemistry SL

[Maximum number: 2]

[Maximum number: 1]

Which product may be obtained by the reduction of $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH}$ ?

$\mathrm{CH}_{3} \mathrm{CH}(\mathrm{OH}) \mathrm{CH}_{3}$

$\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH}$

$\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OCH}_{3}$

$\mathrm{CH}_{3} \mathrm{COOCH}_{3}$

[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 SL

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

State the conditions required for this reaction to occur.

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Covered syllabus points

Question 1

Question 2

Question (a)

Question (b)

Question (c)

Question 3

Question (a)

Question (i)

Question (ii)

Question (b)

Question 4

Question 5

Question 6

Question 7

Question 8

Question 9

Question 10

Question 11

Question 12