Energy is priceless

The Gibbs free energy $\Delta G$ of a reaction determines whether a reaction will occur at a given temperature. It can be calculated with the formula,

$\Delta G = \Delta H - T \Delta S,$

where $T$ is the temperature in kelvin. When the Gibbs free energy of the reaction is negative, the reaction is said to be exergonic , which basically means that the reaction occurs in the direction it is written in.

Since the given reaction occurs at room temperature, we know that $\Delta G < 0$ and $T = 293 \text{ K}.$ Thus,

$\begin{aligned} \Delta H - (293 \text{ K}) \Delta S &< 0 \\ \Delta H &< (293 \text{ K}) \Delta S \\ \dfrac{\Delta H}{\Delta S} &< 293 \text{ K}. \end{aligned}$

The direction of the less-than sign does not change because we are given that $\Delta S > 0.$ We conclude that $\boxed{\dfrac{\Delta H}{\Delta S} < 293 \text{ K}}$ is the only conclusion we can derive from the given information.

(EDIT: I noticed that the answer choices comparing $\Delta H$ and $\Delta S$ directly are not appropriate because they have different units and, thus, cannot be compared directly. Can this be fixed somehow?)