1. Quick Facts ~Chapter 11 It takes energy to cause a substance to go from the liquid to the gaseous state. The boiling point and melting point of a substance is a good measure of the strength of the forces that hold the particles together. Most covalent compounds melt at lower temperatures than ionic compounds do. Ionic substances generally have much higher forces of attraction than covalent substances Ionic compounds with small ions have high melting points.

2. Covalent Compounds Intermolecular forces are the forces of attraction between molecules of covalent compounds. Intermolecular forces include: dipole-dipole forces & London dispersion forces.

3. Dipole-Dipole Dipole-dipole forces are interactions between polar molecules. When molecules are very polar, the dipole-dipole forces are very significant. The more polar the molecules are, the higher the boiling point of the substance.

4. Hydrogen Bonds A hydrogen bond is a dipole-dipole force occurring when a hydrogen atom that is bonded to a highly electronegative atom of one molecule is attracted to two unshared electrons of another molecule. In general, compounds with hydrogen bonding have higher boiling points than comparable compounds. Hydrogen bonds are strong dipole-dipole forces.

5. London Dispersion A substance with weak attractive forces will be a gas because there is not enough attractive force to hold molecules together as a liquid or a solid. However, many nonpolar substances are liquids. London dispersion forces are the dipole-dipole force resulting from the uneven distribution of electrons and the creation of temporary dipoles. The strength of London dispersion forces between nonpolar particles increases as the molar mass of the particles increases. Nonpolar molecules can experience only London dispersion forces. Polar molecules experience both dipole-dipole forces and London dispersion forces.

6. LDF vs Dipole-Dipole Dipole-dipole forces are generally stronger than London dispersion forces. However, both of these forces between molecules are usually much weaker than ionic forces in crystals. When the forces between ions are spread out over large distances, as with large ions or oddly shaped ions that do not fit close together, they do not have as great of an effect.

7. What is a Phase? A phase is a region that has the same composition and properties throughout. Example: ice water is a system that has a solid phase and a liquid phase.

8. Equilibrium A dynamic equilibrium exists when particles are constantly moving between two or more phases yet no net change in the amount of substance in either phase takes place. For an enclosed gas and liquid in equilibrium, the gas particles above the liquid exert pressure when they strike the walls of the container. The pressure exerted by the molecules of a gas, or vapor, phase in equilibrium with a liquid is called the vapor pressure. The boiling point is the temperature at which the vapor pressure equals the external pressure.

9. Vapor Pressure As you increase the temperature of a closed system, more liquid particles escape into the gas phase. Thus, as you increase temperature, the vapor pressure of the substance also increases.

10. Phase Diagram A substance’s state depends on temperature and that pressure affects state changes. A phase diagram is a graph that shows the relationship between the physical state of a substance and the temperature and pressure of the substance.

11. Phase Diagrams A phase diagram has three lines: • One line shows the liquid-gas equilibrium. • Another line shows the liquid-solid equilibrium • A third line shows the solid-gas equilibrium. The three lines meet at the triple point, the temperature and pressure at which the three states of a substance coexist at equilibrium.

12. Phase Diagram for Water At critical point the liquid and vapor phases are identical. The three lines meet at the triple point , the temperature and pressure at which the three states of a substance coexist at equilibrium.

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