Table 2.1 shows electrode potentials for some electrode reactions involving manganese compounds. (i) Aqueous manganate(VI) ions, $$$\mathrm{MnO}_{4}{ }^{2-}$$$, are unstable in acidic conditions and undergo a disproportionation reaction. The $$$E_{\text {cell }}^{\ominus}$$$ for this reaction is +1.14 V . Construct an overall ionic equation for this disproportionation reaction. ................................................................................................................................. (ii) Suggest and explain how the $$$E_{\text {cell }}$$$ value of the disproportionation reaction changes with an increase in pH . ....................................................................................................................................... . ....................................................................................................................................... . .................................................................................................................................

Chemistry

IGCSE&ALevel

CAIE

Exam No：9701_s24_qp_42 Year：2024 Question No：2(c)

Answer:

Knowledge points:

13.2.1.1 homologous series

13.2.1.2 saturated and unsaturated

13.2.1.3 homolytic and heterolytic fission

13.2.1.4 free radical, initiation, propagation, termination (the use of arrows to show movement of single electrons is not required)

13.2.1.5 nucleophile, electrophile, nucleophilic, electrophilic

13.2.1.6 addition, substitution, elimination, hydrolysis, condensation

13.2.1.7 oxidation and reduction (in equations for organic redox reactions, the symbol [O] can be used to represent one atom of oxygen from an oxidising agent and the symbol [H] one atom of hydrogen from a reducing agent)

13.2.2.1 free-radical substitution

13.2.2.2 electrophilic addition

13.2.2.3 nucleophilic substitution

13.2.2.4 nucleophilic addition (in organic reaction mechanisms, the use of curly arrows to represent movement of electron pairs is expected; the arrow should begin at a bond or a lone pair of electrons)

24.2.1.1 standard electrode (reduction) potential

24.2.1.2 standard cell potential

24.2.10 understand and use the equation

24.2.2 describe the standard hydrogen electrode

24.2.3.1 metals or non-metals in contact with their ions in aqueous solution

24.2.3.2 ions of the same element in different oxidation states

24.2.4 calculate a standard cell potential by combining two standard electrode potentials

24.2.5.1 deduce the polarity of each electrode and hence explain/deduce the direction of electron flow in the external circuit of a simple cell

24.2.5.2 predict the feasibility of a reaction

24.2.6 deduce from values the relative reactivity of elements, compounds and ions as oxidising agents or as reducing agents

24.2.7 construct redox equations using the relevant half-equations

24.2.8 predict qualitatively how the value of an electrode potential, E, varies with the concentrations of the aqueous ions

Answer:

Knowledge points:

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