A student studying the relationship between electrical work, standard free energy, and chemical equilibrium compiles the data in the table below for three redox reactions carried out at \(298\text{ K}\).

Reaction	Balanced equation	\(n\)	\(E^\circ\text{ (V)}\)	\(\Delta G^\circ\text{ (kJ/mol}_{\text{rxn}}\text{)}\)	\(K\text{ at } 298\text{ K}\)
1	\(\text{X}^{2+}\text{(aq)} + \text{Y(s)} \rightarrow \text{X(s)} + \text{Y}^{2+}\text{(aq)}\)	\(2\)	\(+0.40\)	\(-77\)	\(3.3 \times 10^{13}\)
2	\(2\,\text{W}^{3+}\text{(aq)} + 3\,\text{Z(s)} \rightarrow 2\,\text{W(s)} + 3\,\text{Z}^{2+}\text{(aq)}\)	\(6\)	\(+0.20\)	\(-116\)	\(1.9 \times 10^{20}\)
3	\(\text{Q}^{2+}\text{(aq)} + \text{R(s)} \rightarrow \text{Q(s)} + \text{R}^{2+}\text{(aq)}\)	\(2\)	\(-0.15\)	\(+29\)	\(8.5 \times 10^{-6}\)

Which of the following best explains why Reaction 2 is more thermodynamically favorable under standard conditions and has a larger equilibrium constant than Reaction 1, even though \(E^\circ\) for Reaction 2 is smaller?