IB Chemistry SL Electron Transfer Reactions

Review IB Chemistry electron transfer reactions through oxidation states, half-equations, electrochemical cells, electrolysis and organic redox.

Exam points

deduce oxidation states and identify oxidized, reduced, oxidizing and reducing species
write half-equations for electrodes in voltaic cells, molten electrolysis or organic redox

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: 6]

(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}

[ 2 ]

(i)

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

[ 2 ]

(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 $e^{-}$ .

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Covered syllabus points

3.2 Electron transfer reactions question 1

3.2 Electron transfer reactions question 2

Question (a)

Question (b)

Question (c)

3.2 Electron transfer reactions question 3

Question (a)

Question (i)

Question (b)