A contradiction, probably from confusion
Processes in nature that lead to an increase in the entropy of the system are in general favourable, but entropy alone is not the deciding factor! 

If we were to follow a full proof we would have to consider Carnot engines and such but for now the result that we are interested in is called the Clausius statement. It is the Clausius statement that really defines the second law and the direction of spontaneity. The Clausius statement leads us to define two parameters depending on the experimental conditions: either the Helmholtz free energy $A$ (when working at constant volume) or the Gibbs free energy $G$ (under conditions of constant pressure). In practice it is much easier to work with $G$ since we can expose our reaction mixture to the atmosphere and call it constant pressure. 

It is these thermodynamic potentials $G$ and $A$ that determine the overall thermodynamic picture of spontaneity of a given process.  

If we work with the Gibbs potential, after some mathematics we can easily express by the following equation:

\begin{equation}
G=H-TS
\end{equation}
Where $H$ is the enthalpy of the system, $T$ is the temperature and $S$ is the entropy. For a given process it is the Gibbs energy that determines the spontaneity of the reaction. Crudely put, the starting material will have a certain $G$ value and the products will have a certain $G$ value, if the change in $G$ over the course of the reaction is negative then we say it is spontaneous. If it is positive it is not spontaneous. 

We can clearly see that the condition for spontaneity is not solely dependant on the entropy alone rather it is also a function of the enthalpy too. Further the entropic contribution to the $G$ is dependent on the temperature indicating that there will be a certain threshold temperature for an endergonic process.

An endergonic process is characterised by having a positive enthalpy change, requiring energy to occur. If the "amount" of energy the process requires is very large then the change in enthalpy will also be large and as we can see this makes the change in $G$ for the process become less favourable. 

Moving on, we can now see that a given process will have a certain energy barrier... an activation energy if you like. Whether a process is thermodynamically spontaneous or not one has to consider the Kinetics too! A simple yet ubiquitous model for Kinetics is the Arrhenius equation which in general says the rate of reaction decreases for processes with a higher energy of activation. 

Turning the later part of your question. The most stable phase (state of matter) is dependent on the conditions of the reaction. The most stable phase for the given combination of temperature and pressure will prevail in the reaction mixture. In practice this leads to the state with the lowest value of $G$. 

Solids actually liberate energy when they come together, especially ionic solids! This energy is called the Lattice energy and it is relative to all the reactants in their gaseous state (as ions). (look at Born-Lande equation for more info)! So although gaseous species will have a higher entropy due to all the rotational vibrational freedom they have ... the state of matter is dictated by $G$. 

Without the full mathematical argument behind what I have said I feel my answer will be incomplete, but I hope it helps you :)