College Prep' Chemistry Ch' 12 p' 1

1
College Prep. Chemistry Ch. 12 p. 1

• Liquid particles are closer together than in a
  gas, and have lower kinetic energy. Fluids are
  liquids and gases that flow and take the shape of
  their container.
• Liquids have a higher density than gases and are
  relatively less compressible.
• Liquids diffuse through other liquids, but the
  process is much slower than for gases, since the
  particles are closer together.
• Surface Tension- force that pulls parts of a
  liquids surface inward, which decreases the
  surface area.
• Capillary Action- attraction of the sides of the
  solid to

2
College Prep. Chemistry Ch. 12 p. 2

• The liquid. Ex. Meniscus on a graduated
  cylinder.
• Vaporization- the process by which a liquid or
  solid changes to a gas. Evaporation the
  process by which particles escape from the
  surface of a non boiling liquid and enter the gas
  state. Ex. Perspiration.
• Boiling- change of liquid to a gas, bubbles form
  throughout the liquid.
• Freezing- the physical change of a liquid to a
  solid by removal of heat. Water turns to ice at
  0 0C
• Section Review p. 366

3
College Prep. Chemistry Ch. 12 p. 3

• The particles of a solid are more closely packed
  than a liquid or a gas. More ordered
  arrangement, vibrations. Definite shape and
  volume
• Crystalline Solid -particles are arranged in an
  orderly, 3-D repeating pattern. NaCl, sugar
• Amorphous Solid particles are arranged
  randomly. Ex. Glass, clay, They do not have a
  regular shape.
• Melting- is the physical change of a solid to a
  liquid by the addition of heat. Melting Point-
  the temperature at which a solid becomes a
  liquid. Ex. 0 0 C for ice to water.
  Supercooled liquids retain properties of a
  liquid at temperatures where they should be a
  solid. Ex. Glass

4
College Prep. Chemistry Ch. 12 p. 4

• Solids are slightly more dense than liquids. The
  particles are more closely packed together.
  Crystal Structure- 3-D arrangement of particles,
  7 crystal structures, cubic, tetragonal,
  hexagonal, trigonal, orthorhombic, monoclinic,
  triclinic
• Unit Cell smallest portion of a crystal lattice
  that shows the three dimensional pattern of the
  entire lattice.
• Ionic Crystals positive and negative ions are
  arranged in a regular pattern. Ions can be
  monatomic or polyatomic. Ex. NaCl or NaNO3
• Covalent Network solids a large array of
  neighboring atoms. Ex. Diamond, Silicon
  Carbide,

5
College Prep. Chemistry Ch. 12 p. 5

• SiC. They have high melting points, and are
  hard, and brittle. They are nonconductors or
  semiconductors.
• Metallic Crystals metal atoms are surrounded by
  a sea of valence electrons. High electrical
  conductivity.
• Covalent Molecular Crystal Structures- covalent
  bonds, held together by intermolecular forces,
  smaller melting points, Sucrose, C12H22O11
• Section Review p. 371

6
College Prep. Chemistry Ch. 12 p. 6

• Equilibrium a dynamic process in which two
  opposing changes occur at equal rates in a closed
  system. In a closed system, matter cannot enter
  or leave, but energy can. Matter and energy can
  escape from an open system.
• Phase any part of the system that has uniform
  composition and properties.
• Condensation- process by which a gas changes to a
  liquid. The rate of condensation depends on the
  numbers of particles that pass from the gas to
  the liquid phase.
• Liquid to Vapor it absorbs heat from its
  surroundings. Double Arrow means it is a
  reversible

7
College Prep. Chemistry Ch. 12 p. 7

• Reversible reaction
• Le Chateliers Principle When a system at
  equilibrium is disturbed by a stress, it will
  shift to attain a new equilibrium that minimizes
  the stress. A stress is a change in
  concentration, pressure, or temperature.
• Endothermic Systems absorb heat energy, add heat,
  favors the products.
• Exothermic Systems release heat energy, cool
  them, to favor the products.

8
College Prep. Chemistry Ch. 12 p. 8

• Larger volume will favor a shift in terms of
  whatever side of the equation has more moles of
  gas.
• Equilibrium Vapor Pressure the pressure exerted
  by a vapor in equilibrium with its corresponding
  liquid at a given temperature is called the
  equilibrium vapor pressure.
• Increasing the temperature causes more liquid
  particles to escape into the vapor phase.

9
College Prep. Chemistry Ch. 12 p. 9

• Volatile Liquids evaporate readily have weak
  forces of attraction between the particles. Ex.
  Ether , Evaporates Rapidly
• Nonvolatile liquids have strong attractive forces
  between the particles. H2O, Evaporate slowly
• Boiling the conversion of a liquid to a vapor
  within the entire liquid, including its surface.
• The equilibrium vapor pressure of the liquid
  equals the atmospheric pressure. The lower the
  atmospheric pressure, the lower the boiling
  point.
• Foods take longer to cook at higher elevations.
• Pressure cookers increases the boiling point of
  the

10
College Prep. Chemistry Ch. 12 p. 10

• Water, resulting in shorter cooking times.
• When water boils, the temperature remains
  constant. Only the amount of heat changes that
  was added.
• Molar heat of vaporization The amount of heat
  energy needed to vaporize one mole of liquid at
  its boiling point. The stronger the attractive
  forces are between the liquid, the more heat
  energy is needed to overcome it.
• Ex. Water, hydrogen bonding is extensive.

11
College Prep. Chemistry Ch. 12 p. 11

• Freezing loss of heat energy occurs. The
  normal freezing point is the temperature at which
  the solid and liquid are in equilibrium at 1 atm.
  (760 torr 101.3 kpa) pressure.
• Melting and freezing occur at the same
  temperature, Equilibrium
• The amount of heat energy required to melt one
  mole of solid at its melting point is its molar
  heat of fusion. This depends on the attraction
  of the solid particles.

12
College Prep. Chemistry Ch. 12 p. 12

• Sublimation change of state from a solid
  directly to a gas. (dry ice) Deposition is the
  reverse process, change of a gas to a solid
  directly.
• Phase Diagram graph of pressure versus
  temperature that shows the conditions under which
  all 3 phases of a substance exist. P. 381 Phase
  Diagram for Water
• AB Curve Ice and Water Vapor coexist,
• AD Curve Ice and liquid water coexist
• AC Curve Liquid water and water vapor coexist.

13
College Prep. Chemistry Ch. 12 p. 13

• Triple Point the temperature and pressure of a
  substance at which all three phases can coexist
  at equilibrium. The critical point indicates the
  critical temperature and critical pressure. The
  critical temperature, tc is the temperature above
  which the substance cannot exist in the liquid
  state.
• The critical pressure Pc is the lowest pressure
  at which the substance can exist as a liquid at
  the critical temperature. For water 217.75 atm
• Section Review p. 382 (1-8)

14
College Prep. Chemistry Ch. 12 p. 14

• Water bent, polar, undergoes hydrogen bonding.
  Ice has a hexagonal arrangement to the molecule,
  it forms an open framework. Ice has a very low
  density, .917 g/mL. Ice floats in liquid water,
  density of water 1 g/mL Important to aquatic
  life in arctic regions
• Odorless, tasteless, colorless liquid
• Heat of fusion of ice 6.009 kJ/mol
• Waters Heat of vaporization 40.79 kJ/mol,
  which is quite high

15
College Prep. Chemistry Ch. 12 p. 15

• Sample Problem 12-1 p. 386
• ? kJ 47 g H2O x 1mole H2O/18 g H2O x 6.009
  kJ/mol 15.7 kJ
• ? kJ 47 g H2O x 1mole H2O/18 g H2O x 40.79
  kJ/mol 106 kJ
• Practice p. 386 (1,2)
• Section Review p. 386 (1-4)
• P. 388 1, 2, 4 ,5, 7-9, 11-13, 16, 18, 19