Phase Diagrams

Chapter 11

Chemical Systems

• Undergo three main processes that change their energy– Chemical reactions– Heating/cooling– Phase transitions

Phase DiagramsPhase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

Phase Diagrams• The AB line is the liquid-vapor interface.• It starts at the triple point (A), the point at

which all three states are in equilibrium.

Phase DiagramsIt ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

Phase Diagrams• Each point along this line is the boiling point of the

substance at that pressure.– Pressure = Force / area– Increasing pressure causes the molecules to be condensed

making the molecular arrangement closer

Phase Diagrams• The AD line is the interface between liquid

and solid.• The melting point at each pressure can be

found along this line.

Phase Diagrams• Below A the substance cannot exist in the

liquid state.• Along the AC line the solid and gas phases

are in equilibrium; the sublimation point at each pressure is along this line.

Phase Diagrams• As temperature increases, the average

kinetic energy of the substance increases causing the spacing between the molecules to increase and eventually change phase

Phase Diagram of Water

• Note the high critical temperature and critical pressure:– These are due to the strong

van der Waals forces between water molecules.

• London dispersion force• Dipole-dipole forces• Hydrogen bonding

– Increase the melting and boiling point of water

Phase Diagram of Water

• The slope of the solid–liquid line is negative.– This means that as the

pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

Phase Diagram of Carbon Dioxide

The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (such as caffeine).

Intermolecular Forces and impact on ideal behavior

• Presence of intermolecular forces among gas particles (including noble gases) leads to deviations from ideal behavior– Therefore, at sufficiently low temperatures or

sufficiently high pressures, gases can condense

Heating/Cooling

• Between phase changes, all heat (energy) absorbed or released changes the average amount of KE (therefore, the T) of the substance– Areas on graph with a

slope

Heating/Cooling

• There is no slope in areas where a phase change is occuring– All energy is being

used to change the phase of matter (rather than change the temperature)

Heating/Cooling• Melting and Boiling

– Endothermic

• Freezing and Condensation– Exothermic

• The amount of energy needed to vaporize = amount of energy needed to condense– Reverse phase changes have equal but opposite

magnitudes of energy gained or released– Same applies for melting and freezing