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States of Matter
Kinetic Molecular Theory
Particles are always in motion.
Temperature is a measure of
average kinetic energy of particles.
Intermolecular forces hold particles
together. Stronger forces require
more energy (higher temp.) to
overcome.
Intermolecular Forces
Always weaker than
chemical bond
Affect structure and
state of matter
Dipole-Dipole Forces
Positive and negative ends of
polar molecules attract each
other.
About 1% as strong as covalent
or ionic bonds
Weaken as distance between
molecules increases
Hydrogen Bonding
Especially strong dipole-dipole
force
Occurs when H bonds to a
strongly electronegative atom—O,
N, or F
Very strong because 1) molecule
is very polar & 2) small size of H
H Bonding
Example—water
More pronounced in molecules
formed from small atoms
(dipoles can come closer)
High boiling point
London Dispersion Forces
Forces that exist in all atoms and
molecules but that are significant
only among Noble gases and
nonpolar molecules
Result from temporary dipoles
formed when electrons distribute
themselves unevenly—can induce a
dipole in a neighboring atom
VERY WEAK
London Dispersion Forces
Stronger in larger atoms or
molecules due to the greater
chance of the formation of
instantaneous dipoles.
Physical Properties
Melting and boiling points are
higher when IM attractions are
stronger
More energy required to
separate molecules
Which would have the
higher boiling point & Why?
Cl2 or F2?
H2O or H2S?
SiBr2 or SBr2?
CH4 or C10H22?
O2 or NO?
States of Matter
Gases—weak IM forces (like
London Dispersion Forces)
Liquids—intermediate IM
forces
Solids—strong IM forces
Liquids & IM Forces
Surface tension—result of IM
forces that resist an increase in
surface area
Capillary action—result of
cohesive forces within liquid
and adhesive forces between
liquid and tube
Liquids (cont’d)
Viscosity—the ability of a liquid
to resist flow (resist change in
shape)
All effects are higher with more
polar molecules.
Solids
Amorphous—without
definite structure
Crystalline—definite
structures
Solids (Crystals)
Ionic solids—made of charged
particles; ions at lattice points
Molecular solids—made of neutral
particles; molecules at lattice points
Atomic solids—made of neutral
particles; atoms at lattice points; 3
types
An atomic solid, an ionic
solid & a molecular solid
Ionic Solids
Ions at lattice points
Closest packed spheres
Arranged to minimize
repulsions and maximize
attractions
Conducts only when melted
Molecular Solids
Lattice positions occupied by
molecules
Internal covalent bonds are strong,
but intermolecular forces are weak
IM force: dipole/dipole if polar
covalent bond; London dispersion
forces (larger in larger molecules)
Atomic Solids
1. Network—directional covalent
bonds; forms giant molecules
(diamond, graphite,and silicon);
highest melting points
2. Metallic—delocalized covalent
bonds; atoms have closest
packing structure; high melting
points
Atomic Solids
3. Group 8A—Noble gases—
London dispersion forces
only; low melting points.
Network Atomic Solids
Strong, directional bonds
Form giant “molecules”
Typically brittle & poor
conductors
Examples—carbon and silicon
Carbon Network
Follows a molecular orbital
(not atomic orbital) model
Diamond
Tetrahedral--sp3 hybridized
bonds stabilize structure
Large gaps exist between
filled and unfilled molecular
orbitals—hard for electrons
to move—no conductivity
Graphite
Fused carbon rings form sheets
Trigonal planar—sp2 hybridized (1
p orbital remains unhybridized)
Delocalized electrons in orbital
causes graphite to be conductive
Figure 10.22: The
structures of diamond
and graphite. In each
case only a small part
of the entire structure
is shown.
Closest Packed Solids
aba pattern—alternating layers—
atoms in 3rd layer lie directly above
atoms in 1st layer—hexagonal unit
cell—body centered
abca pattern—atoms in 1st and 4th
layers are in line; 2nd & 5th layer; 3rd
& 6th layer—face-centered cubic cell
aba Packing
abca Packing
Density of Closest
Packed Solids
To calculate density, you need to
know:
MASS
VOLUME
Mass
Figure out how many atoms
in one unit cell
Multiply by molar mass
Divide by Avogadro’s number
You now know a mass in
grams
Face-Centered Cubic Unit Cell
If these are atoms of calcium, what is
the mass of the cell?
Volume
Determine the length of one
side of the cube by using the
atomic radius (varies depending
on type of unit cell)
Cube the side length.
Simple Cubic--aaa
If the atomic radius of
this atom is 122 pm,
what is the volume?
Body-Centered Cubic--aba
If the atomic radius of the
atom is 246 pm, what is
the volume of the cell?
Face-Centered Cubic--abca
If the atomic radius is 291
pm, what is the volume of
the cell?
Sample Problem
Silver crystallizes in a face-centered cubic
closest packed structure. The radius of the
silver atom is 144 pm. Calculate the density
of silver.
Bonding in Metals
Strong, non-directional bonds
Atoms are hard to separate
but easy to move.
“Electron sea” model
Mobile electrons carry heat or
electricity easily
Band Model or
Molecular Orbital Model
Electrons travel around metal
crystal in a molecular (instead
of atomic) orbitals
Result is a continuum of levels
that eventually merge to form
a band.
Figure 10.19:
The molecular
orbital energy
levels produced
when various
numbers of
atomic orbitals
interact.
Band or MO Model
Empty orbitals close in energy
exist.
Electrons are very mobile into
and out of these similar-energy
orbitals--CONDUCTIVITY.
Semiconductors
Some electrons can cross the
“energy gap” between molecular
orbitals—somewhat conductive
Higher temperatures result in
more electrons’ being able to
reach conductive bands
Doping
Adding other elements with one
more or one less electron than a
semiconductor can increase
conductivity
n-type semiconductor
An element with one more valence
electron is added
More valence electrons are
available to move into conduction
bands
What could be used to dope Si?
p-type semiconductor
An element with one less
valence electron is added
The absence of a valence
electron creates a hole through
which electrons can travel
Alloys
Introducing other elements into
metallic structure is easy.
Substitutional—some “host” atoms
are replaced by atoms of similar
size (Brass = copper + zinc)
Interstitial—small atoms occupy
spaces between metal atoms (such
as carbon used to harden steel)
Phase Changes
Phase Changes
Condensation/Evaporation
Boiling point
Freezing/Melting
Melting point
Deposition/Sublimation
Energy of Phase Changes
Heating/Cooling Lab
Energy of Phase Changes
Heating Curve
Temperature changes when only
one phase is present
No temperature change during
phase change
For Water
Heat of fusion—334 J/g
Heat of vaporization—2260 J/g
Specific heat
1.84 J/goC for gas
4.18 J/goC for liquid
2.1 J/goC for solid
What amount of energy is needed to
completely melt 45 g of ice at its
melting point?
What amount of energy is needed to
completely vaporize 45 g of water at
its boiling point?
Calculate the amount of energy
required to raise the temperature of
125 g of ice from -12oC to 75oC.
What is the final temperature/state
of 95 g of ice at -15oC if 65 kJ of
energy are added?
The heat of vaporization for carbon
dioxide is 571 J/g. How much energy
would be required to change 6.92 g
of dry ice to carbon dioxide gas?
Vapor Pressure
In a closed system, #
of liquid molecules
decreases as they
enter the gas
phase. Eventually,
an equilibrium is
reached—constant
number of
molecules in both
phases.
Equilibrium
Vapor Pressure
At equilibrium evaporation rate
is exactly the same as
condensation rate.
Vapor pressure = the pressure
of the vapor present at
equilibrium
Measuring Vapor Pressure
Difference between
atmospheric pressure with a
vacuum above
and with vapor
above.
Effect of IM Forces
Low IM forces lead to high vapor
pressure; less attraction allows
molecules to vaporize more easily
High IM forces lead to low vapor
pressure; molecules are tightly held
in liquid
Enthalpy (DH) of Vaporization
Change in energy
when a liquid
vaporizes
Vapor pressure
varies with
temperature.
(Why??)
Note: Function is
not linear
Linear Function
Graphing ln of
vapor pressure
and 1/T gives
a linear graph.
Can use slopeintercept form
of equation to
solve for
points.
Line Equation
DHvap (1)
ln(Pvap) =+C
R
(T)
Format is y = mx + b
R is the universal gas constant.
C is a constant (unique for each
substance)
Clausius-Clapeyron Equation
Since C is temperature
independent, by measuring vapor
pressure at several different
temperatures, heat of vaporization
can be calculated.
Problem 2
Ethanol’s vapor pressure at 30.o C
is 100. torr. At 62o C the vapor
pressure has increased to 400.
torr. What is DH for ethanol?
In Your Scientist’s Notebook
What is ethanol’s vapor
o
pressure at 15 C?
Heating Curve
Changes of
state are
evident on a
heating
curve:
Heating Curve
As heat is added
temperature typically
increases.
At the plateaus, all
energy is going into
the phase change, so
temperature is
constant.
Enthalpy (DH) of Fusion
Enthalpy change that occurs at
the melting point when a solid
melts or freezes
Problem
The enthalpy of fusion for water is
6.02 kJ per mole. How much
energy would be required to
change 43.2 g of ice to liquid
water?
What is the enthalpy change when
52.8 g of CCl4 freezes? Enthalpy of
fusion is 2.51 kJ/mol.
Boiling Point
The temperature at which the
vapor pressure of a liquid is
equal to the atmospheric
pressure
Varies with pressure (Think
about a pressure cooker.)
Phase Change Enthalpy
The amount of heat gained or lost—
DH
Endothermic: boiling (evaporating),
melting, sublimating
Exothermic: condensing, freezing,
depositing
Phase Change Entropy
The amount of disorder in a
system
Increasing entropy: boiling,
melting, sublimating
Decreasing entropy: condensing,
freezing, depositing
Phase Diagram
Lines show equilibrium boundaries
for phase changes
Equilibrium—the system shows no
net change
Dynamic equilibrium—two
processes of change occur at
exactly the same rate
Typical Phase Diagram
Points of Interest
Critical temperature—the temp.
above which a gas cannot be
liquefied at any pressure
Critical pressure—the pressure
required to liquefy the gas at
critical temperature
Points of Interest
Triple point—the point at which
solid, liquid & gas have equal
vapor pressure and exist in
equilibrium
“Normal” boiling or melting
point—read at 1 atm pressure
What is the enthalpy /
entropy change when…
Evaporation occurs?
Sublimation occurs?
Freezing occurs?
Melting occurs?
Deposition occurs?
Calculations with Enthalpy
To calculate total energy change,
consider each stage in the process.
Ice at -10o C to steam at 110o C:
Heating ice from -10o to 0o C
Melting
Heating water from 0o to 100o C
Boiling
Heating steam from 100o to 110o C
Finding the Numbers
Melting or Freezing: heat of fusion
(DHfus)
Vaporizing or Condensing: heat of
vaporization (DHvap)
Temperature changes: specific
heats (q = mcDT)
Sign depends on whether heat is
gained or lost in the process.
Example
What is the total energy required to
change 45.0 g of ice at -5o C to water at
50o C?
Specific heat of ice: 2.1 J/g oC
Specific heat of water: 4.21 J/g oC
DHfus: 6.02 kJ/mol
DHvap: 40.7 kJ/mol