Five hypothetical redox reactions are listed in the table below

CHEMISTRY · Multi-part · AP · Chemistry · Thermodynamics and Electrochemistry · Thermodynamics and Electrochemistry

Problem

Five hypothetical redox reactions are listed in the table below, each carried out under standard conditions (298 K, 1 bar, 1 M for all aqueous species). Values of the standard cell potential Ecell∘E^{\circ}_{\text{cell}} and the standard enthalpy change ΔH∘\Delta H^{\circ} are provided for each reaction.
ReactionEcell∘E^{\circ}_{\text{cell}} (V)ΔH∘\Delta H^{\circ} (kJ\,mol−1^{-1})
I. A3+^{3+} + 3 B   ⟶  \;\longrightarrow\; A + 3 B+^{+}+0.56–72
II. 2 C2+^{2+} + D2−^{2-}   ⟶  \;\longrightarrow\; 2 C+^{+} + D–0.12–95
III. 3 E+^{+} + F3−^{3-}   ⟶  \;\longrightarrow\; 3 E + F+0.03+44
IV. 2 G3+^{3+} + 3 H2−^{2-}   ⟶  \;\longrightarrow\; 2 G2+^{2+} + 3 H−^{-}–0.41–18
V. J+^{+} + K   ⟶  \;\longrightarrow\; J + K+^{+}+1.02+115
Use the information in the table to analyze the thermodynamic favorability of each reaction. (a) Identify every reaction that is thermodynamically favored at standard conditions. Briefly state the criterion you use in making this determination. (b) Rank the thermodynamically favored reactions from the strongest to the weakest driving force. Support your ranking with quantitative evidence drawn from the table. (c) Select one reaction from the table whose ΔH∘\Delta H^{\circ} sign gives a different impression of favorability than the reaction’s actual spontaneity. Explain, in chemical terms, why enthalpy alone is insufficient to predict spontaneity in redox processes.

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