In a reduction reaction characterized by an equilibrium constant (K) of \(4.8 \times 10^2\), it is essential to understand the implications of this value. An equilibrium constant greater than 1 indicates that the reaction favors the formation of products, suggesting that the reaction is spontaneous. This spontaneity is linked to the Gibbs free energy change (\( \Delta G \)), which must be negative for spontaneous reactions. Therefore, in this case, the reaction indeed has a negative change in standard Gibbs free energy.

Furthermore, spontaneous reactions in electrochemistry are typically associated with galvanic cells, which harness redox reactions to generate electricity. The relationship between Gibbs free energy and cell potential (\(E\)) is given by the equation:

\[ \Delta G = -nFE \]

Here, \(n\) represents the number of moles of electrons transferred, and \(F\) is Faraday's constant. Since the reaction is spontaneous, the standard cell potential (\(E\)) must be positive, contradicting any assertion that it could be negative.

In summary, the correct interpretation of the reaction is that it is spontaneous, has a negative standard Gibbs free energy, and produces electricity, confirming that it operates as a galvanic cell. Thus, the only accurate statement regarding this reaction is that it is spontaneous and has a negative change in Gibbs free energy.