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AP Chem ch. 13 Solids and Liquids

General Chemistry edition 7, Thomson
Condensed phases
solids and liquids
1. Have definite shape
2. Are nearly incompressible
3. Usually have higher density than liquids
4. Are not fluid
5. Diffuse only very slowly through solids
6. Have an ordered arrangement of particles that are very close together; particles usually have only vibrational motion
1. Have no definite shape
2. Have definite volume
3. Have high density
4. Are fluid
5. Diffuse through other liquids
6. Consist of disordered clusters of particles that area quite close together; particles have random motion in three dimensions
1. Have no definite shape
2. Are compressible
3. Have low density
4. Are fluid
5. Diffuse rapidly
6. Consist of extremely disordered particles with much empty space between them; particles have rapid, random motion in three dimensions
when IMFs overcome KE of gas particles, condensation occurs
Why is it hard to compress a liquid?
particles are so close together that there is no empty space between them and thus it is hard to
liquids diffuse into other liquids with which they are miscible
if temperature is lowered sufficiently, at ordinary pressures, stronger but shorter-range attractive interactions overcome the reduced kinetic energies of the molecules to cause solidification
at a given pressure depends on the nature of short-range interactions among particles; characteristic of each substance ~ very specific solidification, very ordered
Intermolecular Forces
forces between molecules (weaker)
Intramolecular Forces
forces within compounds (stronger)
ion-ion interactions
according to Coulomb's law, force of attraction between two oppositely charged ions is directly proportional to the charges on the ions, q+ and q-, and inversely proportional to the square of the distance between them, d.
F = q+q-d^(-2)
Energy = Fd so energy of attraction between two oppositely charged ions is directly proportional to the charges on teh ions and inversely proportional to the distance of separation
E = q+q-d^(-1)
ionic bonding
strong, causes high melting points; forms rigid lattice structures which do not conduct electricity
molten ionic compounds
ions are free to move, thus they are great conductors of electricity
As charges on ions increase..
product of q+q- increases; ionic substances containing multiply charged ions, such as Al3+, Mg2+, O2-, S2- usually have higher melting and boiling points
MgO, CaO, BaO
highest boiling point to lowest boiling point (larger the size, lower boiling point?)
Dipole-Dipole interactions
Permanent dipole-dipole interactions occur between polar covalent molecules because of the attraction of partially positive atoms of one molecule to the partially negative atoms of another molecule.
Rate by which dipole dipole forces vary as d increases
(electrostatic forces decrease by 1/d62 as d increases)
Average energy of a dipole-dipole bond
4 kJ per mole
Average energy of a ionic and covalent bond
400 kJ per mole
Hydrogen bonding
case of dipole-dipole bonding; strong hydrogen bonding occurs among polar covalent molecules containing H and one of the three small, highly electronegative elements - F, O, or N
Hydrogen bond donor
molecule with the partially positive H
Hydrogen bond acceptor
molecule receiving the partially positive H
Average hydrogen bond energy
15 to 20 kJ per mole
Properties caused by hydrogen bonding
high melting and boiling points for: water, methanol, ammonia; keeps double helix structure of DNA intact (AT = 2 bonds, CG = 3 bonds); causes high surface tension of water; causes solid water (ice) to be less dense than liquid water
Dispersion forces
weak attractive forces, only important over extremely short distances
They result from attraction of positively charged nucleus of one atom for the electron cloud of an atom in nearby molecules. This induces temporary dipoles in neighboring atoms or molecules. As electron clouds become larger and more diffuse, they are attracted less strongly by their own (positively charged) nuclei. Thus, they are more easily distorted or polarized by adjacent nuclei.
Dispersion forces vary by
1/d^7 as distance increases
IMF present in symmetrical nonpolar substances
dispersion forces
Stronger dispersion forces
larger molecules or have more electrons; polarizability increases with increasing numbers of electrons and therefore with increasing sizes of molecules. Therefore, dispersion forces are genearlly stronger for molecules that are larger or have more electrons. For large polarizible molecules, total effect of dispersion can be greater than dipole-dipole interactions or hydrogen bonding.
What boils at a higher temperature?
polar covalent compounds boil at higher temperatures than nonpolar compounds
among nonpolar compounds, heavier compounds boil at higher temperatures than lighter compounds
resistance to flow of a liquid
what allows a liquid to flow
molecules must be able to slide past one another
stronger the IMFs, viscosity
longer molecules are, viscosity
the unit used to express viscosity
temperature increases, viscosity
surface tension
measure of the inward forces that must be overcome to expand the surface area of a liquid
why liquid droplets are spherical
molecules inside are affected by IMF attractions from all over, but those on the surface are only attracted to the interior which pull the surface layer into the center. The most stable situation is one in which the surface area is minimal. The sphere has the least possible surface area, so drops of liquid assume spherical shapes.
cohesive forces
hold a liquid together
adhesive forces
attraction between a liquid and another surface
surface of water; concave shape in glass
mercury's meniscus
convex shape because its cohesive forces are stronger than its attraction to glass
capillary action
one end of a capillary tube is immersed in a liquid; if adhesive forces exceed cohesive forces, the liquid creeps up the sides of the tube until a balance is reached between adhesive forces and the weight of the liwuid (how plants suck up nutrients and water with their roots)
evaporation or vaporization
process by which molecules on the surface of a liquid break away to go into the gas phase while at a temperature lower than boiling point
as temperature increases, rate of vaporization
evaporative cooling
only energized particles leave the surface of a liquid, remaining particles have lower average KE and thus are cooler
gas to a liquid
dynamic equilibrium
rate of evaporation equals rate of condensation in a closed container; continuously occuring, do not equal 0, but no net change occurs
LeChatelier's Principle
a system at equilibrium, or changing toward equilibrium, responds in the way that tends to relive or "undo" any stress placed on it
vapor pressure
partial pressure of vapor molecules above the surface of a liquid at equilibrium at a given temperature
as temperature increases, vapor pressure
ALWAYS increases; rate of evaporation increases with increasing temperature
volatile liquids
easily vaporized, high vapor pressures
stronger cohesive forces, vapor pressure
lowers; water's hydrogen bonds cause low vapor pressure
tempature is raised enough for vapor pressure to be high enough so that bubbles can persist, rise to the surface, and burst, releasing vapor into the air
boiling point
temperature at which its vapor pressure equals the external pressure
normal boiling point
temperature at which the vapor pressure of a liquid is equal to 1 atm (760 torr)
as heat energy is added to a pure liquid at boiling point, temperature
remains constant; excess heat gets converted into potential energy used to overcome the cohesive forces in the liquid to form vapor
separating a mixture of liquids with different boiling points by heating the mixture slowly until the temperature reaches a point where the most volatile liquid boils off
specific heat or molar heat capacity
amount of heat that must raise one gram of a substance by one degree Celsius
molar heat or enthalpy of fusion
delta H fus (solids to liquids, liquids to solids)
molar heat or enthalpy of vaporization
delta H vap (liquids to gases, gases to liquids)
Heats of vaporization
increase as boiling points and IMFs increase, and vapor pressures decrease
heat of condensation
reverse of heat of vaporization
distillation = economical?
no! high heat of vaporization of water makes it too expensive to vaporize large volumes of water to purify it
(would take a lot of heat for water to become a gas because it has such a high specific heat and can absorb a lot of heat without changing its temperature)
Clausius Clapeyron equation
1. to predict the vapor pressure of a liquid at a specified temperature
2. to determine the temperature at which a liquid has a specified vapor pressure
3. to calculate delta H vap from measurement of vapor pressures at different temperatures
volatile liquids' properties
low- cohesive forces, viscosity, surface tension, specific heat, boiling point, heat of vaporization
high - vapor pressure, rate of evaporation

nonvolatile = opposite
melting point
(freezing point) temperature at which its solid and liquid phases coexist at equilibrium
normal melting point
melting point at 1 atm; changes in pressure have small effects on melting point, have large effects on boiling points
solid straight to a gas
gas straight to a solid
condensation, freezing, deposition
evaporation, melting, sublimation
how sublimation works
solids exhibit vapor pressure as liquids do, but have lower vapor pressures
solids with high vapor pressures
sublime easily
phase diagrams
show equilibrium pressure-temperature relationship among the different phases of a given pure substance in a closed system
important components of a phase diagram
A = triple point, temperature and pressure at which solid, liquid, and gas phases occur at equilibrium with each other
D = pressure below which substance cannot exist as a liquid
AB line = solid-liquid equilibrium line; negative slope is solid is less dense than liquid, positive slope is liquid is less dense than solid; melting curve
C = critical point, temperature after which gas cannot be recompressed into a liquid
AC line = liquid-gas equilibrium line
AD line = solid-gas equilibrium line; sublimation curve
supercritical fluid
a substance above its critical temperature
critical pressure
the pressure required to liquefy a gas at its critical temperature; upper limit for distinct phase changes
amorphous solids
noncrystalline solids
examples of amorphous solids
glasses (flow slowly), some kinds of plastics, rubber, and amorphous sulfur
crystalline solids
well deined shape, sharp melting temperatures
amorphous solids' IMFs
vary in strength within a sample
amorphous solids' melting points
soften over a temperature range
shattering a crystalline solid
cleaving occurs preferentially along crystal lattice planes
shattering an amorphous solid
shatters irregularly to yield pieces with curved edges and irregular angles
way to test purity of crystalline solid
sharpness of melting point, impurities disrupt the IMFs and cause melting to occur over a considerable temperature range
14 unit cells
cubic - simple, body centered, face centered
tetragonal - simple, body centered
otrhorhombic - simple, body centered, face centered, end centered
monoclinic - simple, end centered
triclinic - simple
hexagonal - simple
rhomboderal - simple
x-ray diffraction by crystals
arrangement of particles in crystalline solids are determined indirectly by x-ray diffraction (scattering)
monochromatic x-ray beam is defined by a system of slits and directed onto a crystal. the crystal is rotated to vary the angle of incidence theta. at various angles, strong beams of deflected x-rays hit a photographic plate. upon development, the plate shows a set of symmetrially arranged spots due to deflected x-rays. different crystals produce different arrangements of spots.

consider reflection grating instead of diffraction grating
more electrons, more it scatters
different substances that crystallize in the same type of lattice with the same atomic arrangement
single substance that can crystallize in more than one arrangement
simple cubic
1 atom per unit cell
8 * 1/8 = 1
body centered cubic
2 atoms per unit cell
8 * 1/8 = 1
+ 1 at center of cell
face centered
4 atoms per unit cell
8 * 1/8 = 1
+ (6 * 1/2)

1 + 3 = 4
4 types of solid bonding
covalent solids, molecular solids, ionic solids, metallic solids
metallic solids
1. body centered cubic, face centered cubic, and hexagonal close packed
2. particles of unit cell = metal ions in "electron cloud"
3. strongest interparticle forces = metallic bonds attraction between cations and electrons
4. properties = soft to very hard; good thermal and electrical conductors; wide range of melting points
ionic solids
1. face-centered cubic, body-centered cubic, face0centered cubic lattice
2. particles of unit cell = anions, cations
3. strongest interparticle forces = electrostatic
4. properties = hard; brittle; poor conductors of heat and electricity (except when molten or put into solution); high melting points
molecular solids
1. particles of unit cell = molecules
2. strongest interparticle forces = dispersion, dipole-dipole, and/or hydrogen bonds
3. properties = soft; poor thermal and electrical conductors; low melting points
covalent solids
1. particles of unit cell = atoms
2. strongest interparticle forces = covalent bonds
3. properties = very hard; poor thermal and electrical conductors; high melting points
band theory of metals
overlap interaction of two atomic orbitals = one bonding, one antibonding
creates band of orbitals that belong to crystal as a whole
ability of metal to conduct electricity
based on ability of electrons to jump to higher energy vacant orbitals in the same band when an electric field is applied; resulting net flow of electrons through the crystal is in the direction of the applied field

electrons = delocalized
element or compound with filled bands that are only slightly below, but do not overlap with empty bands; only difference between insulator and semiconductor = size of energy gap
conduction band
band within which electrons move to allow electrical conduction
as temperature decreases, electrical conductivity
do not conduct electricity
band gap
energy difference between bands
filled bands slightly below, but do not overlap with empty bands; do not conduct electricity at low temperatures, but a small increase in temperature is enoguh to excite some high energy electrons into the empty band
lustrous appearance
mobile electrons absorb wide range of wavelengths of radiant energy as they jumpt ot higher energy levels; emit photons of visible light and fall back to lower levels within the conduction band
malleable or ductile
crystal is easily deformed; metal ions are identical, and imbedded in a sea of electrons, and as bonds are broken, new ones are readily formed with adacent metal ions
different forms of the same element in the same physical state
coordination number
number of nearest neighbors of an atom or ion (in a crystal)