The spontaneity of a reaction is determined by the Gibbs free energy change (ΔG), given by the equation: ΔG = ΔH − TΔS.

For a reaction to be spontaneous, ΔG must be negative (ΔG < 0).

When a small positive number is subtracted from a large negative number (or a large positive TΔS term) is subtracted from a small ΔH, the resulting ΔG. Therefore, a reaction with a positive (endothermic) and small ΔH, and a positive (increase in disorder) and large ΔS, will be spontaneous, especially at high temperatures where the TΔS term dominates.

A. Non-spontaneous: A non-spontaneous reaction would have a positive ΔG. Under the given conditions, ΔG is negative, so the reaction is not non-spontaneous.

B. favour forward reaction only: A spontaneous reaction inherently favors the forward reaction, leading to the formation of products. Thus, "spontaneous" is a more complete and accurate thermodynamic description than just "favor forward reaction only," which could imply specific equilibrium conditions not fully defined by spontaneity alone (though they are related).

C. favour backward reaction only: The backward reaction is spontaneous if the forward reaction is non-spontaneous (ΔG>0). Since the forward reaction is spontaneous under these conditions, the backward reaction is non-spontaneous.