A phase diagram is a diagram that shows the state of a substance at different temperatures and pressures. The sketch below shows a phase diagram for a hypothetical substance.
 

The line from the red A to the red B is the line along the liquid-vapor equilibrium line. At pressures and temperatures along this line, the substance is in equilibrium between the liquid and vapor phases. The line ends at the critical temperature and pressure of the substance, beyond which the there is no difference between the liquid and gaseous phases. Substances past this point are known as supercritical fluids.

In the same vein, the A-C line is the line where solid and liquid co-exist, and A-D is the area where solid and vapor coexist. The three lines meet at the triple point of the material.

The slope of the A-C line depends on the relative densities of the solid and the liquid. A positive slope means that the solid is more dense than the liquid, a negative slope means the opposite. In the diagram above, as the pressure increases the solid phase becomes stable at higher and higher temperatures: this is because it is the denser phase and under higher pressures it requires more energy to overcome the exterior pressure and change to the less dense liquid phase. Only a few materials have a liquid phase that is denser than the solid: water is the most important example.

Example: If we begin at the point marked A on the phase diagram below and increase the temperature, what will happen to the substance?

Solution: The point marked A on the phase diagram is in the solid region, so the substance begins as a solid. As we increase the temperature at constant pressure, we move right along the red line in the diagram below. The line passes first through the liquid area, then the vapor, so the substance moves from solid->liquid->vapor.

Information provided by: http://learn.chem.vt.edu