Chapter 10 - Solids and Liquids

I. Properties of Solids and Liquids

• They have a high density
• They are Essentially Incompressible
• They undergo little thermal expansion
• They have fixed volume.
• No KMT to explain
  • Each particle has kinetic energy
  • The basic particles have significant attractions for each other (held close together).
  • Particles occupy a significant portion of the volume
  • the basic particles are not in random motion (motion restricted by interactions with other particles)

II. Intermolecular forces

• force of attraction that holds molecules together (forces within a molecule)

• London Dispersion Forces - (nonpolar molecules) induced dipole will cause attraction of the partial neg. of one molecule to the partial positive of another.
  
• Dipole-Dipole attractions - (polar molecules) molecules with a permanent dipole can align themselves so that the negative end of one molecule is attracted to the positive end of another (dipole attractions).
  
• Hydrogen Bonding - strongest dipole dipole attraction. Found with molecules that have highly electronegative atoms bonded to hydrogen. Electrostatic attraction between a hydrogen bonded to F, O or N.
  

III. Solid State

• Amorphous solid - solids with no defined shape
• Crystalline solids - the molecules or ions are arranged in a regular, symmetrical structure called a crystal lattice.
  
  • IONIC Solids - crystalline solids in which ions are the basic particles making up the crystal lattice. ION-ION forces very strong - NaCl
  • MOLECULAR Solids - crystalline solid in which molecules are the basic particles and are held together by London dispersion forces.LONDON DISPERSION forces - Sugar
  • NETWORK Solids - atoms are covalently bonded to on another throughout the entire crystal.
  • METALLIC Solids - positive metal ions occupy regular positions in the crystal lattice with the valence electrons moving freely among these positive ions.

IV. Liquid State

• Surface tension - the forces that causes the surface of a liquid to contract.
• Viscosity - a measure of a liquids resistance to flow.
• Vapor Pressure
  • Vaporization - liquid state changes to the gaseous state.
  • Evaporation - vaporization below the boiling point.

Requirements to vaporize

1. Must be at the surface of the water
2. Must have enough kinetic energy to overcome the intermolecular forces.
3. Increasing the temperature decreases the energy barrier

Condensation - change in state from gaseous to liquid state.

Equilibrium vapor pressure - The pressure exerted by the vapor above a liquid at a given temperature

Sublimation - The vaporization of a solid.

Boiling point - the temperature where the vapor pressure = atmospheric pressure. High altitude cooking requires different cooking times.

Normal boiling point - the temperature where the vapor pressure = 760 torr.

V. Energy Changes Vs. Changes in State

• Energy is required to break intermolecular forces to boil or melt
• Melting and Freezing
• HEAT of FUSION - the amount of heat (in cal or J) required to melt one gram of a substance (DH_fus).

• Boiling and Condensation
• Heat of Vaporization - the amount of heat required to vaporize one gram of a substance (DH_vap)

• Combine with specific heat from chapter 2 - examples

VI. Heating Curve of water

• Graphical representation of the temperature of a solid is heated through the two phase changes plotted as a function of time heating.

1. Heating Ice - (-10ºC to 0ºC) increase in KE and increase in temp.
2. Melting Ice - (OºC) water and ice mixed together
3. Heating water - levels off then begins to increase temp. again (waits for all ice to turn into water).
4. Boiling water (100ºC) - vaporizing water begins
5. Heating vapor - levels off then increases again (waits for all water to convert to vapor).