LIQUIDS AND SOLIDS
LIQUIDS: Why are they the least
common state of matter?
1. Liquids and K.M.T.
 Are particles in constant motion? Spacing?
Kinetic Energy? Attractive forces?
 Fluid: a substance that flows and hence
takes the shape of its container.
Properties of Liquids
1. High Density: 1000x greater than gases,
10% less dense than solids.
2. Relatively Incompressible: Water’s
volume only decreases 4% under
1000atm of pressure!
3. Can diffuse:
Slower in liquids than gases due to:
slower motion and attractive forces.
4. Surface Tension: a force that pulls
adjacent parts of a liquid’s surface,
thereby decreasing surface area.
Ex: bug “walking” on water.
Capillary Action: related to surface
tension; attraction of the surface of
a liquid to the surface of a solid.
Ex: water transport from roots to leaves
5. Vaporization: Process by which liquid
gas.
Evaporation: Process by which particles
escape from the surface of a nonboiling
liquid.
Boiling: change of a liquid to vapor
bubbles appearing throughout the liquid.
SOLIDS
1. Solids and K.M.T.




More closely packed than liquids or gases.
Intermolecular forces are VERY effective.
Only vibrational movement.
Crystalline vs. Amorphous (glass) solids.
Properties of Solids
1. Definite shape and volume
2. Melting point:
 Crystalline Solids: Definite melting point, KE of
particles overcome attractive forces of solid.
 Amorphous Solids: No definite melting point,
Supercooled liquids.
3. High Density and Incompressibility
4. Low diffusion rate: very slow
Crystalline Solids
1. Crystal structure = 3D arrangement of
particles of crystals.
a) 7 types of crystals- pg. 369
2. Unit Cell = smallest portion of a crystal
that shows the 3D structure.
Binding Forces in Crystals
1. Ionic Crystals:
 NaCl
 Strong electrostatic forces holds it together.
 Hard, brittle, high melting pts.
2. Covalent Molecular Crystals:
 Nonpolar: H2, CH4 vs. Polar: H20, NH3
 Covalently bonded molecules held together
by intermolecular forces.
 Low melting points, soft, easily vaporized.
3. Covalent Network Crystals:
 Diamond (C)X , Silicon Carbide (SiC)X
 Giant molecules that extend indefinitely- each
atom is covalently bonded to neighboring atom.
 Hard, Brittle, High Melting Points
4. Metallic Crystals:
 Metal atoms surrounded by sea of valence
electrons.
 High electrical conductivity, Melting Points vary.
Amorphous Solids
1. No regular pattern of atoms.
2. Large range of melting points.
3. Examples: glass, plastics
Equilibrium
Equilibrium = a dynamic condition in which 2
opposing changes occur at equal rates in a
closed system.
Equilibrium and State Changes:
ex: Evaporation of water in a closed container
(assuming constant temp.)
Equilibrium Equations:
liquid + heat energy
vapor
Le Chatelier’s Principle
When a stress is applied to a system at
equilibrium, the system will respond in a way
to minimize that stress. (Stress= change in
temp, pressure, concentration)
ex: liquid + heat
vapor
Equilibrium Vapor Pressure of a
Liquid
 The pressure exerted by a vapor in
equilibrium with its liquid at a given temp.
Increases as temp. increases (How can we
explain this using KMT?)
Volatile vs. Nonvolatile Liquids:
Volatile liquids have WEAK forces of attraction,
therefore they evaporate readily. Ex: ethanol
Vapor Pressures of varying
substances at different temps.
Boiling
The conversion of a liquid vapor within
the ENTIRE liquid. Occurs when the
vapor pressure in the bubble =
atmospheric pressure.
Phase Diagram for Water
Shows the
conditions
under
which the
phases of a
substance
can exist.
Triple Point:
Indicates the
temp and
pressure at
which the solid,
liquid, gas
coexist.
Critical Point:
indicates the
critical temp.
and pressure
of a substance
Critical Temperature = Substance can’t exist as a
liquid above this temperature (only as a gas).
Critical Pressure = Lowest pressure at which the
substance can exist as a liquid at the critical
temperature. (any lower P, it’s a gas)