A hypothetical reaction is observed to be non-spontaneous at low...

The spontaneity of a reaction is determined by the Gibbs free energy equation:

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

For a reaction to be spontaneous, $\Delta G$ must be negative ($\Delta G < 0$).

At low temperatures, the magnitude of the $T\Delta S$ term is small, so the sign of $\Delta G$ is largely determined by $\Delta H$. Since the reaction is non-spontaneous at low temperatures, $\Delta G > 0$, which implies $\Delta H > 0$.

At high temperatures, the magnitude of the $T\Delta S$ term becomes large and dominates the equation. For the reaction to become spontaneous ($\Delta G < 0$) at high temperatures, the term $-T\Delta S$ must be negative, which implies $\Delta S > 0$.

Therefore, the reaction has $\Delta H > 0$ and $\Delta S > 0$.

Answer: $\Delta H > 0$, $\Delta S > 0$