What is the difference between "spontaneous" and "energetically favourable"? I read somewhere that the difference between energetically favorable reactions and spontaneous reactions is that energetically favorable reactions are ones where energy is released, i.e., $\Delta H&lt;0$, whereas spontaneous reactions are ones where the change in Gibbs free energy is negative, i.e., $\Delta G&lt;0$. Thus it is possible for a reaction to be energetically favorable but not spontaneous if $\Delta H$ is negative, and yet $\Delta G$ is positive because $\Delta S$ is negative and the product of it and T is large enough to compensate for $\Delta H$ as per this equation: $\Delta G = \Delta H - T(\Delta S)$. 

My questions are:

1) Is this true?

2) If so, what would be an example of a reaction that is energetically favorable but not spontaneous?
Here is a very simple example of an exothermic, nonspontaneous process at room temperature:

$$\ce{H2O_{(l)} -&gt; H2O_{(s)}}$$

You need to extract heat from water to freeze it, but it is obvious that at $25\ ^{\circ}\mathrm{C}$, this does not happen spontaneously. Only if you decrease the temperature to decrease the contribution of the negative $\Delta S$ term will this reaction occur spontaneously.