1. Dipole-dipole:
a. normal attraction between polar-polar
molecules; an electrostatic force; opposite sides of the molecules attract each
other (opposite charges attract)
b. Hydrogen bonding
A.
relationship of structure and properties? structure affects intramolecular and intermolecular forces,
knowledge of these predicts the properties (like melting and boiling points)
B. what do molecules have in common that do not behave as expected?
1. contain hydrogen
2. hydrogen is bonded to a highly electronegative atom (N, O, or F)
a. high polarity
b. the bare proton of hydrogen gives a high
∂+
C.
why is hydrogen the only element that shows this? hydrogen has no internal electrons to cover
its proton, once it has lost its one electron it is only a bare proton
D.
define hydrogen bonding:
attraction of the ∂+ of the hydrogen for the negative
portion of another molecule
E.
what is the strength of a hydrogen bond? weaker than chemical bonds but stronger than van der Waals, holds
molecules together hydrogen bonding:
seen only when hydrogen is bonded to nitrogen, oxygen or fluorine; have
very high dipole-dipole force, caused by the small size of the hydrogen
allowing atoms to get very close and the high partial positive charge of the
exposed bare proton of hydrogen, and the great polarity of the bond accounts
for the special properties of water:
1. water can surround
and attract both positive and negative ions which allows it to dissolve most
ionic compounds and many polar covalent compounds
2. water molecules
have a very strong attraction for each other and it takes a lot of energy to
break this force, therefore the majority of water on Earth is in the form of a
liquid
2. London dispersion forces - force increases with an
increase in the size of the molecules - the more electrons there are the more
distorted the electron cloud can become causing the force to be greater
a. Induced dipole: polar-nonpolar
b. Temporary or momentary dipole: nonpolar-nonpolar
How
particles are attracted (what force
holds them together) defines what type
of solid is formed.
I. Types of Solids
A.
Covalent Molecular Crystal:
discrete, covalent compounds (molecules) held together by relatively
weak intermolecular forces, collectively called van der Waals forces
3. properties:
a. relativity low melting points
b. may or may not be soluble in water and these
solutions may or may not conduct electricity
c. not conductors of electricity in any phase
B.
Ionic Crystal: ions are held
together by an electrostatic force;
between dissimilar atoms, usually a main group metal and a nonmetal or a
metal and a polyatomic ion; make the
compounds commonly called salts properties:
1. relatively high melting points (NaCl melting point is 800 ºC)
2. usually soluble in water and these water
solutions conduct electricity
3. molten and gas phases conduct electricity
4. form crystalline solids from their solutions
C.
Covalent Network crystal (macromolecules): strong, directional covalent bonds; the covalent bonds holding these together are ten times stronger
then van der Waals; ie. diamond,
silicon dioxide, silicon carbide
D.
Metallatic Crystal: strong and
nondirectional bonding; formed by the
sharing of delocalized electrons forming a sea of electrons that the metal
cation is attracted to
1. lustrous, shiny
2. Generally solids at room temperature
3. Malleable and ductile because positive ions can move within metal
structure (it is difficult to separate
the atoms in a metal but the positive ions are easily moved around as long as
they are in contact with one another)
4. High melting points
5. Electrical conductors in all phases, because of electrons moving
in the metal
a. Electrons must be in the conduction band to
move
b. Moving electrons are called free or
delocalized electrons
c. Degree of properties determined by the number
of delocalized electrons - more delocalized electrons = metal is harder, more
dense and has a higher melting point
6. alloys - mixtures of metals such as brass, bronze, sterling
silver and 14 karat gold
II. Structure of Solids
A.
The study of solids is the study of crystals, all "true"
solids are crystals.
B.
All intensive properties of crystals remain the same: definite melting point, density, angles
faces, etc.
C.
A crystal is a rigid structure whose particles are arranged in a
regular, repeating pattern.
1. the shape and properties of crystals are determined by the
attractive forces between the particles
2. There are seven crystal systems, based on their angles and the
length of their axis.
a. cubic - unit cells axes are all equal
lengths, all the unit cell axes meet in 90º angles
b. tetragonal - unit cells axes have two of
equal lengths and one not equal to the other two, all the unit cell axes meet
in 90º angles
c. orthorhombic - unit cells axes have none of
equal lengths, all the unit cell axes meet in 90º angles
d. rhombohedral (trigonal)- unit cells axes are all equal lengths, none of the
unit cell axes meet in 90º angles
e. monoclinic - unit cells axes have none of equal
lengths, two of the unit cell axes meet in 90º angles and the other doesn't
f. triclinic - unit cells axes have none of
equal lengths, none of the unit cell axes meet in 90º angles
g. hexagonal - three of the unit cell axes are
of equal length and the fourth is not equal to these three, three unit cell
axes meet at 60º angles forming 120º angles and the other meets these at a 90º
angle
2. A unit cell is the smallest repeating segment in a crystal, it is
a mental model
a. unit cells are arranged in a space lattice -
a three dimensional representation of repeating unit cells, a mental model
b. there are 14 unit cells
c. three basic units cells in the cubic system
1. simple cubic (SC)
2. body-centered
cubic (BCC)
3. face-centered
cubic (FCC)
III. Solid descriptors
A.
allotropic - a substance that occurs in more than one form
B.
isomorphic - different substances having the same crystal structure
C.
polymorphic - a substance that has different crystal structures
D.
mesomorphic - a middle form, between a solid and liquid
E.
amorphic - no regular repeating pattern and no definite melting point
F.
hydrate - a crystal that has water molecules bonded to it, anhydrous
means without water
G.
hygroscopic - a substance that absorbs water from the atmosphere or
surroundings
H.
deliquescent - a substance that absorbs enough water from the atmosphere
or surroundings to dissolve in that water
I.
desiccant - a substance that absorbs water from another and is used as a
dehydrating agent
J.
efflorescent - a substance that releases water into the atmosphere or
surroundings
IV. Liquid crystals
A.
mesomorphic substance - does not have order in all three dimensions
1. nematic - 1 dimension of order, 2 dimensions disordered
2. smectic - 2 dimension of order, 1 dimensions disordered
B.
exhibit anisotropy - has a property in one direction but not another
C.
used in calculators and watches for display
1. high frequency current makes transparent
2. low frequency current makes opaque
V. Amorphous solids - not true solids but
appear to be solids
A.
sometimes called super-cooled liquids
B.
no definite melting point but a wide band of temperatures where it gets
softer and softer (or harder and harder)
C.
viscosity - the resistance of a liquid to flow, as amorphous solids cool
they become more and more viscous, like glass
D.
metastable - a long lasting stable form of an amorphous solid
VI. Semi-conductors (see pgs 768-769) -
metalloids that under the right circumstances can conduct an electrical current
A.
crystal doping - deliberately adding an impurity to a crystal
B.
transistor - a doped crystal, conducts electricity better, ability to
control where electrons move, dope the
semi-conductor with atoms that have more electrons to create an excess of
electrons (an n-type semiconductor) or with atoms that have less electrons to
create electron holes (a p-type semiconductor) - allows electrons to more
C.
chip - an integrated circuit module
Steps to producing a chip: purify crystal - zone purification; cut wafers ; oxidize surface; apply photoresist; apply
stencil pattern; expose to light; etch pattern; expose to doping material;
clean surface
I. Kinetic Theory
A.
if you increase the temperature, what happens to the speed of particles?
increasing the temperature causes an increase in the speed of particles
B.
if particles move with more speed, what happens to the force of their
collisions? what does this cause? more speed causes particles to collide
with more force which forces them to move apart
C.
therefore: most solids and
liquids expand when heated
II. van der Waals
A.
what attraction force holds particles together? van der Waals forces holds particles together
B.
what happens if the temperature increases?
1. an increase in temperature causes an increase in the energy of
the particles
2. the increased energy is enough to overcome van der Waals
forces
3. particles can slide past each other
4. melting occurs
C.
what happens if the temperature decreases?
1. a decrease in temperature causes an decrease in the energy of the
particles
2. the decreased energy causes the particles to slow down enough for
van der Waals forces to take affect
3. particles can no longer slide past each other
4. freezing occurs
D.
where does this happen? there is
a definite point where melting and freezing occur, the melting point of a substance equals the freezing point of the
substance
III. Temperature and kinetic energy
A.
what is temperature? temperature
is a measure of the average kinetic energy
B.
relationship? at a given
temperature the particles have the same kinetic energy
C.
Boltzmann's law: at a given
temperature, some particles have more kinetic energy than the average and some
have less kinetic energy than the average à average distribution
curve of energy – see figure 12-13 on page 379
IV. Vapor
A.
what happens if particles have above the average kinetic energy? if particles have above the average kinetic
energy they may have enough energy to overcome the attraction forces of its
neighbors and escape from the surface forming a gas or vapor
1. gas: at room temperature,
the substance is normally a gas
2. vapor: at room temperature,
the substance is normally a liquid or solid
B.
reverse: the reverse can also
happen, particles can lose kinetic energy and return to the liquid or solid
state - doesn't happen often in an open system
V. Equilibrium
A.
what happens in a closed container?
in a closed system, some particles leave from the surface of the liquid
and at the same time some particles return to the surface of the liquid
B.
what is equilibrium? equilibrium
is when the number of particles leaving the liquid equals the number of
particles returning to the liquid state, the net change is zero
C.
what is dynamic equilibrium?
implies that the process of equilibrium is continuing to happen, it is
not static
D.
what does saturated mean?
saturated means there is an equilibrium between the liquid and vapor
phase, the vapor is saturated with the vapor of the liquid, rate leaving = rate
returning
VI. Le Chateliers principle
A.
what is vapor pressure? vapor
pressure is the pressure exerted by the vapor phase
B.
what is it dependent upon? vapor
pressure temperature dependent
1. high temperature = high vapor pressure
2. low temperature = low vapor pressure
C.
relationship? a closed container
reaches equilibrium at a specific pressure for a given temperature
D.
what is Le Chateliers principle?
when an external stress is applied to a system at equilibrium, the
system adjusts so the stress is reduced and an equilibrium is re-established
E.
examples: ice skater on
"ice" and wire through a block of ice
VII. Vapor pressure and intermolecular forces
A.
vapor pressure is caused by?
vapor pressure is caused by particles leaving the surface of a liquid or
solid
B.
how does this happen? particles
have overcome van der Waals forces
C.
relationship
1. high vapor pressure: easy
to leave the surface, must have low intermolecular forces
2. low vapor pressure: hard
to leave the surface, must have high intermolecular forces
VIII. Melting point
A.
what happens in a closed solid - liquid system? there is a dynamic equilibrium between the
solid and the liquid
B.
what is the relationship between vapor and solid and liquid? the solid is in equilibrium with its vapor
and the liquid is in equilibrium with its vapor, but it is the same vapor so
they have the same vapor pressure
C.
define melting point: melting
point is the temperature where the vapor pressure of the solid equals the vapor
pressure of the liquid
IX. Sublimation
A.
definition: sublimation is when
at room temperature a solid goes directly to a gas without ever becoming a
liquid
B.
relationship to vapor pressure?
substances that sublime have very high vapor pressures
C.
examples: CO2 (dry
ice) and iodine crystals
X. Boiling point
A.
define evaporation: in an open
system, particles leave the surface of the liquid which leads to the
disappearance of the liquid (happens only at the surface
B. if the temperature is increased, what happens to the kinetic
energy and resultant vapor pressure? as
the temperature is increased, particles gain more kinetic energy and can leave
the surface which increases the vapor pressure
C. how does this affect the internal pressure? increasing kinetic energy helps particles
the internal pressure of the substance
(internal pressure is the caused by the air pressing down on the liquids
surface)
D.
what does this cause? this
causes particles to collide and push apart which causes bubbles of gas/vapor to
form
E.
what happens? the density of the
gas/vapor is less than the liquid so the bubbles tend to rise, boiling
F.
how is boiling different from evaporation? boiling happens throughout the liquid, evaporation happens just
at the surface of the liquid
G.
define boiling point: boiling
point is the temperature where the vapor pressure of the liquid is equal to the
atmospheric pressure
H.
relationship?
1. high atmospheric pressure = high boiling point
2. low atmospheric pressure = low boiling point
I.
define normal boiling point:
normal boiling point is when the atmospheric pressure is at the standard, 1.00 atm
XI. Condensation
A.
what is condensation?
condensation is the reverse of boiling
B.
relationship to energy?
particles lose kinetic energy and return to the liquid state
C. relationship to boiling point?
the boiling point of a liquid is equal to the condensation point of the
vapor of the liquid
XII. Volatile versus nonvolatile
A.
what is volatile?
1. a substance boils at a low temperature
2. evaporates rapidly at room temperature
3. have a high vapor pressure
4. low intermolecular forces
B.
what is nonvolatile?
1. a substance boils at a high temperature
2. evaporates slowly at room temperature
3. have a low vapor pressure
4. high intermolecular forces
XIII. Liquefaction
A.
definition: condensation of a gas (substance that is normally a gas at room
temperature)
B.
when does this happen? when the
gases get close enough and slow enough to let van der Waals forces take
effect, must cool below a certain point
and sometimes compress the gas
C.
what is Tc? critical
temperature
1. definition: the
temperature above which no amount of pressure will cause a gas to liquefy
2. relationship to intermolecular forces?
a. high:
high Tc = high intermolecular forces
b. low:
low Tc = low
intermolecular forces
XIV. Phase diagram
A.
what is a phase diagram? a phase
diagram is a diagram that shows the relationship between temperature, pressure
and states of matter
B.
what is the triple point?
triple point is the temperature and pressure where all three state of
matter (solid, liquid and vapor) are in equilibrium
C.
solid-liquid slope line
1. what does it mean if it is negative?
a. the solid state is less dense than the
liquid state
b. if you increase the pressure, the
melting/freezing point is lowered
c. the substance expands when it freezes
2. what does it mean if it is positive?
a. the solid state is more dense than the
liquid state
b. if you increase the pressure, the
melting/freezing point is raised
c. the substance contracts when it freezes
XV. Energy and change of state
A.
energy diagram
B.
what do these mean? as the
temperature increases, the kinetic energy increase up to the melting
point; then energy is stored (potential
energy) to make the substance more disorganized; the temperature remains constant until all of the substance has
melted: the same is true at boiling: the reverse happens at condensation and
freezing
C.
what is enthalpy of fusion?
also called heat of fusion, the
energy needed to melt one gram of a substance at its melting point (= heat of crystallization)
D.
what is enthalpy of vaporization?
also called heat of vaporization,
the energy needed to vaporize one gram of a substance at its boiling point
(= heat of condensation)
XVII. Water and hydrogen bonding
A.
attraction of molecules?
1. water molecule is attracted to four other water molecules
2. the H is attracted to the O of another water while the H of
others is attracted to its O
B.
what happens at freezing? at
freezing this attraction causes an open crystalline structure to form, it
causes the molecules to spread apart
C.
what happens at melting? the
lattice structure starts to collapse and the molecules move closer together, it
becomes more dense
D.
how does the density of water change?
ice is less dense than water, as
ice melts the density increases until it reaches a maximum at 3.98 ºC, at this
point the water molecules start to move apart again and the density starts to decrease
XVIII. Surface tension
A.
definition: the apparent
elasticity of the surface of a liquid
B.
what is the relationship between surface tension and forces?
1. forces on the surface are unbalanced
2. the net force acts perpendicular into the liquid pulling the
surface inward
C.
what does this cause?
1. liquid forms spheres when dropped, sphere has least surface area
2. capillary rise - attraction of a liquid for solid sides of a
container, breaks surface tension