1. Liquids & Solids
Chapter 14

2. Intermolecular Forces & Phase Changes
Section 14.1
Intermolecular Forces & Phase Changes

3. Properties…
Gases
Solids
Low Density
High Density
Highly Compressible
Slightly Compressible
Fill Container
Rigid (keeps its shape)
These properties indicate that the components of a solid are close together and exert large attractive forces on each other.
The properties of liquids lie somewhere between those of solid and of gases.
It takes about seven times more energy to change liquid water to steam (a gas) at 100⁰C than to melt ice to form liquid water.
Therefore, going from the liquid to the gaseous state involves a much greater change than going from the solid to the liquid.

4. Forces…
Intermolecular Forces – Attractive forces that occur between molecules.
Intramolecular Forces – attractive forces that occur between atoms in a molecule; chemical bonds.

5. Dipole Movement
When molecules with dipole moments are put together, they orient themselves to take advantage of their charge distributions. Molecules with dipole moments can attract each other by lining up so that the positive and negative ends are close to each other.
This is called a dipole-dipole attraction

6. Hydrogen Bonding
Special name for unusually strong dipole-dipole attractions that occur among molecules in which hydrogen is bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine).
The especially large electronegativity value of the oxygen atom compared with that of other group members causes the O-H much more polar than other similar bonds. This leads to very strong hydrogen bonding forces among the water molecules.
This is why an unusually large quantity of energy is required to overcome these interactions and separate the molecules to produce the gaseous state, which is why water has a very high boiling point.

7. London Dispersion Forces
Relatively weak intermolecular forces resulting from a temporarily uneven distribution of electrons that includes a dipole in a neighbor.
These forces exist among noble gas atoms and non-polar molecules.

8. Atoms can develop a temporary dipolar arrangement of charge as the electrons move around the nucleus. This instantaneous dipole can then induce a similar dipole in a neighboring atom.
The interatomic attraction thus formed is both weak and short-lived.

9. Normal Boiling Point – The boiling temperature of a liquid under one atmosphere of pressure. The temperature at which the vapor pressure of a liquid is exactly one atmosphere.
Normal Freezing Point – The freezing temperature of a liquid under one atmospheric of pressure.

10. Phase Change Diagram

11. Recall, that changes of state from solid to liquid and from liquid to gas are physical changes. No chemical bonds are broken in these processes.
Molar Heat of Fusion – The energy required to melt one mole of a solid.
Molar Heat of Vaporization – The energy required to vaporize one mole of a liquid.

12. Vapor Pressure & Boiling Point
Section 14.2
Vapor Pressure & Boiling Point

13. Vaporization aka Evaporation
The process in which a liquid is converted to a gas.
To escape into the vapor phase a given component must have sufficient speed to overcome the intermolecular forces of the liquid. Thus only the fastest moving components can escape the surface of the liquid.
Evaporation is a cooling process and therefore endothermic.

14. Condensation The process in which a vapor is converted to a liquid.
Eventually, the same number of molecules are leaving the liquid as are returning to it and the rate of condensation equals the rate of evaporation.
At this point no further changes occurs in the amounts of the liquid or vapor, because the two opposite processes exactly balance each other and when that occurs we say that the system is in equilibrium.

15. Vapor Pressure
The pressure exerted by a vapor in equilibrium with its liquid phase at a certain temperature.
Example: Water
The vapor pressure of the water must be equal to atmospheric pressure before boiling can occur.
Water will begin to form bubbles at 100⁰C, at this point the water molecules are energetic enough to sustain a pressure of 1 atm inside the bubbles.

16. Section 14.3
Properties of Solids

17. Types of Solids Crystalline Solid
A solid characterized by the regular arrangement of its components.
Ionic Solid
A solid that contains cations and anions
Molecular Solid
A solid composed of molecules
Atomic Solid
A solid that contains atoms at the lattice points

18. Ionic Solids
Are stable substances with high melting points that are held together by the strong forces that exist between oppositely charges ions.
When thinking about the structures of ionic solids, it is best to imagine the ions as spheres packed together as efficiently as possible.

19. Molecular Solids The fundamental particle is a molecule
They tend to melt at relatively low temperatures because the intermolecular forces that exist among the molecules are relatively weak.

20. Atomic Solids

21. Bonding in Metals Metals have familiar physical properties
They can be pulled into wires, hammered into sheets
They are efficient conductors of heat and electricity
They are durable and have high melting points
The bonding in most metals is strong but nondirectional.
The electron sea model, which pictures a regular array of metal atoms in a “sea” of valence electrons, that are shared among the atoms in a nondirectional way and that are quire mobile in the metal crystal.

22. Alloy
A substance that contains a mixture of elements and has metallic properties.
Substitutional
Interstitial