CHAPTER 12: LIQUIDS AND SOLIDS
I. Properties of Liquids
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particles are close together and lower KE than gases
intermolecular attractive forces are significant (definite volume)
can flow (shape of container)
high density compared to gases
relatively incompressible
will diffuse into another liquid if it can dissolve in it
Surface tension
 force that tends to pull parts of a liquid’s surface together (attraction toward the center of the
liquid)
 cohesion
Adhesion
 liquids can be attracted to the sides of their containers
 can move up narrow tubes (until gravity stops it)
 capillary action
Phase changes:
 vaporization: liquid  gas (heat must be ABSORBED)
o evaporation: occurs at surface of liquid (not boiling)
o boiling: occurs throughout the liquid
o COOLING PROCESS
 freezing: liquid  solid (heat must be REMOVED)
II. Properties of Solids
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particles very closely packed
intermolecular attractive forces are VERY significant (definite volume)
particles in fixed positions (vibrational motion possible)
high density
incompressible
very low rate of diffusion
Two Types:
 crystalline solids (crystals)
o particles arranged in orderly, repeating, geometric pattern (page 369 Fig. 12-8)
o definite, rigid shape
o definite melting point (depends on type of atoms involved—see page 370 Table 12-1)
 amorphous solids (supercooled liquids)
o no distinct geometric shape
o can change shape (particles arranged randomly)
III. Phase Diagrams
For a pure substance IN A CLOSED CONTAINER, a phase diagram will show the conditions under which
the phases of a substance exist (are stable).
 graph of pressure (y-axis) vs. temperature (x-axis)
 triple point: all 3 phases exist and are stable
 critical temperature (tc): highest temperature that the liquid phase can exist at
 critical pressure (Pc): LOWEST pressure that the liquid can have when at the critical temperature
 critical point: point where the critical temperature and critical pressure intersect
 NOTE: Denser phase (solid or liquid) will have greater area
IV. Phase Changes
When a substance changes phases, several things MUST happen:
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the temperature of the substance stays CONSTANT
there must be a change in ENERGY
complementary phase changes occur at the same T
Heating Process (energy released to ENVIRONMENT)
 freezing (liquid  solid + heat)
 condensation (gas  liquid + heat)
 deposition (gas  solid + heat)
Cooling Process (energy absorbed from ENVIRONMENT)
 melting (solid + heat  liquid)
 vaporization (liquid + heat  gas)
 sublimation (solid + heat  gas)
V. Equilibrium Systems
In a closed system (matter cannot enter or leave), a dynamic condition can arise:
forward reaction  reverse reaction
Example:
 water in a closed container will evaporate, but the water level will not change very much because
condensation also occurs
The water level is eventually becomes constant, even though BOTH processes are taking place because:
forward rate = reverse rate
When this condition arises, a state of EQUILIBRIUM is reached.
Le Chatelier’s Principle
If a system at equilibrium is subjected to a stress, the system will SHIFT to try and relieve that stress (get
rid of it).
Example:
liquid + heat  vapor
Stresses:
Increase T of system:
Decrease T of system:
Add vapor:
Remove vapor:
Decrease V (increase P)
Increase V (decrease P)
System Shifts:
VI. Vapor Pressure
As a liquid evaporates, molecules at the surface of the liquid escape into the gas phase with an upward
force.
This force causes an upward pressure.
*** Vapor Pressure: pressure caused by the evaporation of a liquid.
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as T increases, the average KE of the molecules in the liquid increases
more molecules can escape into the vapor phase
 vapor pressure increases
Vapor pressure also depends on the strength of intermolecular forces in the liquid.
 weaker the forces, HIGHER the v.p.
Volatile liquids—high vapor pressure (evaporate quickly)
VII. Boiling
A liquid will boil when
V.P of the LIQUID = ATMOSPHERIC PRESSURE
NOTE: the highest temperature a liquid can get is its boiling point.
Boiling point: TEMPERATURE when v.p. = a.p.
Normal b.p.: temperature when v.p. = 1 atm
VIII. Heats of Fusion and Vaporization
Molar heat of FUSION: the amount of heat energy needed to MELT one MOLE of a substance at its
melting point.
Molar heat of VAPORIZATION: the amount of heat energy needed to VAPORIZE one MOLE of a
substance at its boiling point.
NOTE: all energy added goes toward breaking intermolecular attractions (T does not change during a phase
change)
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Units: kJ/mol
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Values will depend on the strength of the intermolecular forces (stronger the force, larger the
value)
Molar heat of Solidification: the amount of heat energy RELEASED when one MOLE of a substance
FREEZES.
Molar heat of Condensation: the amount of heat energy RELEASED when one MOLE of a substance
CONDENSES.
NOTE:
heat of solidification = -heat of fusion
heat of vaporization = -heat of condensation
For water (at 1atm):
 molar heat of fusion = 6.009 kJ/mol
 molar heat of vaporization = 40.79 kJ/mol
1 mole H2O = 6.009 kJ
1 mole H2O = 40.79 kJ
Problem:
How much heat will be needed to melt 3.5 moles of water at its melting point?
How much heat will be needed to boil 3.5 moles of water at its boiling point?
IX. Water
Structure:
Properties:
 polar molecule
 strong hydrogen bonding between molecules
 MM = 18.01 g/mol
 liquid at room temperature
 relatively low vapor pressure
 relatively high b.p.
 high heat of fusion
 high heat of vaporization
 high surface tension
 solid less dense than liquid phase
Density of water:
Temperature
0oC (solid)
0 oC (liquid)
3.98 oC
NOTE: Density of water will DECREASE above 3.98 oC.
Structure of ICE:
Density (g/cm3)
0.917
0.99984
1.000