Why is water so awesome?
• (Near) universal solvent
– The high polarity (and, therefore, hydrogen
bonding power) of water means it can dissolve so
many compounds – ionic compounds, polar, nonionic compounds and even non-polar gases.
Why is water so awesome?
• Thermal properties
– Water has a high heat capacity (higher than
almost any liquid). High intermolecular forces
mean water can absorb a lot of heat before it
boils. This allows the earth to remain at a steady
temperature.
– Water has a high heat of vaporization which gives
it enormous cooling power. Example:
Adding 4 kJ of heat to 1000g of water causes the
temperature to rise 1 °C, but only 2g of that water
has to evaporate to keep the remaining 998g at a
constant temperature.
Why is water so awesome?
• Surface properties
– High surface tension and high “capillarity” (both
results of hydrogen bonding) are critical to plant
life (land and aquatic).
Why is water so awesome?
• Density of solid and liquid water
– Because solid water (ice) is less dense than liquid
water (as a result of hydrogen bonding), ice floats
on water. This protects aquatic life, erodes rocks,
but sometimes freezes your pipes…
Chapter #16 – Liquids and Solids
16.1) Intermolecular Forces
16.2) The Liquid State
16.3) An Introduction to Structures and Types of Solids
16.4) Structure and Bonding of Metals
16.5) Carbon and Silicon: Network Atomic Solids
16.6) Molecular Solids
16.7) Ionic Solids
16.8) Structures of Actual Ionic Solids
16.9) Lattice Defects
16.10) Vapor Pressure and Changes of State
16.11) Phase Diagrams
What happens when substances
freeze into solids?
• Less thermal energy available
• Less motion of the molecules
• More ordered spatial properties
Crystals
(or Glasses)
Properties of Crystals
Unit Cell: The smallest repeating unit needed to describe
the complete extended structure of a crystal (through
repetition and translation).
Some Common Crystal Lattices
Figure 16.9
How do we know what crystals look like?
x-ray diffraction
Figure 16.10
Why x-rays?
To satisfy the diffraction conditions, the wavelength of
light needs to be comparable to the unit cell
dimensions
Ångstroms (10-10meters)
Diffraction Conditions
Figure 16.11
From trigonometry: nl = 2d sinq
Key Point:
Typical interatomic and intermolecular distances are d ≈ 1.0 to
20Å
Typical x-ray wavelengths are l ≈ 0.01 to 10Å
The X-ray Diffraction Experiment
Crystal
structure
Molecular
structure
3 Types of Crystalline Solids
Atomic
Solids
e.g. all metals, Si,
Carbon (diamond,
graphite)
Ionic Solids
e.g. salts like NaCl
Molecular
Solids
e.g. protein crystals, sucrose
Figure 16.12
Physical properties of crystals
Crystal structure
Physical Properties
Bonding forces
• melting point
• mechanical strength
• electrical properties
Example: Copper and Diamond are both atomic solids, but they
have very different physical properties:
Copper:
very soft, lower melting point (1083°C), excellent conductor
Diamond:
hardest known substance, higher m.p. (3500°C), insulator
The BIG Picture/A Summary
Rank by energy
Recalling Intra/Inter-molecular Forces
Structure and Bonding in Metals
Metals in solids can be treated as hard spheres that
(usually) pack in a way to minimize the empty space
between spheres. This is called closest packing.
Two distinct structures can be formed by closest
packing of atomsa cubic structure and a hexagonal structure
Cubic: x = y = z
All unit cell angles = 90°
Hexagonal: x = y ≠ z
some unit cell angles
≠ 90° (60° or 120°
instead)
Cubic Closest Packed Structure (ccp)
Figure 16.13
Figure
16.15
Face Centered Cubic (fcc)
Hexagonal Closest Packed Structure
Figure 16.13
Figure 16.14
Hexagonal Prism (hcp)
Both fcc and hcp have the same number
of nearest neighbor interactions = 12
Figure 16.16
How many atoms are in the fcc unit cell?
Figure 16.17
6(atoms on faces) + 8(atoms on corners)
= 6(1/2) + 8(1/8)
=3+1=4
Some Common Crystal Lattices
Figure 16.9
How many atoms are in:
Simple cubic unit cell?
8 atoms in corners
8(1/8) = 1
Body-centered cubic unit cell?
8 atoms in corners + 1 atom
8(1/8) +1 = 2
Packing Efficiency
The fraction of the volume (often expressed as %) of the
unit cell that is occupied by atoms, ions, or molecules
PE = fv =
volume of space occupied by particles
volume of unit cell
Example: the face centered cubic unit cell
PE = fv =
(#atoms in unit cell)(volume of atom)
(edge of cubic unit cell)3
(4)(4/3 pr3)
=
e3
e
Example (cont’d): the face centered cubic unit cell
4r
e
e
e
So,
e
e2 + e2 = (4r)2
e =r8
(4)(4/3 pr3)
PE = fv =
e3
= 0.74 or 74%
for the face centered cubic unit cell
Cubic Crystal Lattices
PE = 52%
6 nearest neighbors
PE = 68%
8 nearest neighbors
PE = 74%
12 nearest neighbors
Figure 16.9
Determining Atomic Radius
from a Crystal Structure
Problem: Barium it has a body-centered cubic unit cell and a density of
3.62g/cm3. What is the atomic radius of barium?
Plan: Since an atom is spherical, we can find its radius from its volume.
1. Volume of a sphere is V = 4/3 pr3
2. From the density (mass/volume) and the molar mass (mass/mole),
we find the molar volume of Ba metal.
3. Since it crystallizes in the body-centered cubic structure, 68% of this
volume is occupied by Ba atoms, and the rest is empty.
4. Dividing by Avogadro’s number gives the volume of one Ba atom,
from which we determine the atomic radius.
Solution:
Volume/mole of Ba metal =
1
1 cm3
137.3 g Ba
 MM =

density
3.62 g Ba
1 mol Ba
= 37.9 cm3/mol Ba
Volume/mole of Ba atoms = (volume/mol Ba) * (packing efficiency)
= 37.9 cm3 / mol Ba * 0.68 = 26 cm3/mol Ba atoms
26 cm3
1 mol Ba atoms
Volume/Ba atom =
*
1 mol Ba atoms
6.022 x 1023 Ba atoms
= 4.3 x 10 - 23 cm3/Ba atom
Finding the atomic radius of Ba from the volume of a sphere:
V of Ba atom = 4/3pr3 and r3 =
r=
3
3V
4p
=
3
3(4.3 x 10 - 23cm3)
4 x 3.14159
3V
4p
= 2.17 x 10 - 8 cm = 2.17Å
Properties of metals
• Conductivity, malleability and ductility are all a result
of the bonding forces between particles being strong
and non-directional
• It is difficult to separate metal atoms, but fairly easy
to move them as long as they stay in contact.
• MO theory helps explain this. Each atom brings its
own set of orbitals, the varying energy levels of each
set combine to make bands of filled and empty
orbitals
• As a result, there are large numbers of highly mobile
electrons  they move from filled orbitals to empty
ones.
Metallic Bond: Cations in a “Sea of Electrons”
Group 1A Metals (Li, Na, K…etc.)
Group 2A Metals (Be, Mg, Ca…etc.)
ENERGY BANDS IN SOLIDS
Band theory  MO theory applied to solid crystals (nearly
infinite groups of atoms).
• Instead of well-separated bonding, nonbonding, and antibonding MOs, a large group of close-packed atoms has very
closely spaced orbitals: a lower-energy valence band of filled
MOs and a higher-energy conduction band of empty MOs.
• Metals have no band gap (energy separation) between the
valence and conduction bands. Nonmetals and most
compounds have a large band gap. Metalloids
(semiconductors) have a small band gap. Metalloid elements
are B, Si, Ge, As, Sb, and Te (along periodic-table metalnonmetal line).
Orbital
energy
levels
Easily movable (“conduction band”)
Localized at atom
ENERGY BANDS IN SOLIDS
• Electrical conductivity (e movement across a crystal lattice)
requires excitation of a few e to mostly empty orbitals (the
conduction band). Because of the different-sized band gaps,
this occurs easily in metals, to some extent in metalloids, and
not at all in nonmetals and most compounds.
• Conductivity of metals decreases with increasing T because
atomic motion retards the cross-lattice electron movement.
Conductivity of metalloids increases with increasing T
because higher T provides more excitation energy.
• You can make an e- jump to the conduction band by giving it
energy. One way is for it to absorb light Ephoton= hn= hc/l
• In a laser: electrons jump down from the conductance band to
emit a photon of light with E=hn
CLOSE PACKING: METALS
• In alloys, some atoms of another element fit into closepacked lattice of a metal.
• Types:
o substitutional: where two metal atoms have similar
ratomic, one can simply replace the other (example: brass
= 1/3 Zn, 2/3 Cu). Strengthens metal by adding bond
polarity.
o interstitial: element with small ratomic can fit into metallattice holes (example: in steel, nC = 16% nFe).
Strengthens metal by adding new element-metal bonds.
Examples of
Metal Alloys
Crystals formed by
different atoms or ions;
Contains a mixture of
elements and has
metallic properties
Chapter #16 – Liquids and Solids
16.1) Intermolecular Forces
16.2) The Liquid State
16.3) An Introduction to Structures and Types of Solids
16.4) Structure and Bonding of Metals
16.5) Carbon and Silicon: Network Atomic Solids
16.6) Molecular Solids
16.7) Ionic Solids
16.8) Structures of Actual Ionic Solids
16.9) Lattice Defects
16.10) Vapor Pressure and Changes of State
16.11) Phase Diagrams
Network Atomic Solids
• Unlike metals, network atomic solids
contain strong covalent bonds.
• These solids tend to be brittle and
relatively non-conductive (heat and
electricity).
• Representative elements for these types
of solids are: carbon and silicon.
Network Solids: Carbon
• Network solid: not close-packed. Each atom’s
environment is determined instead by covalentbond geometry (think “VSEPR”).
• Carbon occurs in three different atomic forms
(allotropes):
− Diamond
− Graphite
− Fullerenes (molecular solid)
Diamond
• One giant molecule: web of
C–C single bonds, one
connecting each pair of C
atoms, tetrahedral, sp3 C
• Hardest natural substance;
must break bonds to deform
• mp = 4,440 C; secondhighest-melting natural
substance
Graphite
• Planar sheets of fused
hexagonal rings, sp2 C
• Sheets held together by
delocalized p bonds
(conducts electricity along
sheets). Soft
• mp = 4,492 C; highestmelting natural substance.
Atomic Networks
Fig. 16.26
(Typical metal)
Fig. 16.27
Graphite consists of layers of carbon atoms
p-Electron System in Graphite
Fig. 16.28
Network Solids: Silicon
• Silicon is to geology what carbon is to biology
• Silicon is right below carbon on the periodic table…
– So why is SiO2 so different from CO2?
SiO2 at room temperature
CO2 at room temperature
Si is too big to form strong p bonds to oxygen –
no double bonds as in CO2
SiO Bond Network in Quartz
Ring structures
Tetrahedral geometry
Examples of silicate anions, all of which are
based on SiO44- tetrahedra
Two-dimensional representations of
(a) a quartz crystal and (b) a quartz glass
When silica is heated above melting and then cooled
rapidly, the result is a glass – cooled too quickly for
regular crystalline patterns to form.
soda-lime
aluminosilicate
borosilicate
optical
Ceramic – another siliconbased substance
• Ceramics are make of clays
fired at high temperatures.
• They are brittle, non-metallic
materials that consist of
minute crystals of silicates
suspended in a glassy
cement.
• Unlike regular glass, ceramic
cannot be melted and remelted.
Silicon - continued
• Elemental silicon has the same structure as diamond,
but the gap between filled and empty MOs is smaller
in silicon:
Diamond
Silicon
The smaller band gap means some electrons can cross the
gap – silicon is a semiconductor.
Semiconductors
• Pure semiconductors (like silicon) allow only a few
electrons to cross the band gap, BUT they can be doped
with other elements to create greater or fewer valence
electrons available for movement
More electrons: n-type
Fewer electrons: p-type
n-type: conductivity is increased
by doping it with elements that
have more valence electrons
than the host crystal.
For example, silicon doped with
arsenic (1 more e-)
p-type: conductivity is increased by
doping it with elements that have
less valence electrons than the host
crystal, creating a hole.
For example, silicon doped with boron
(1 less e-)
Semiconductors
n-dopant
(electron rich, like arsenic)
p-dopant
(electron deficient, like boron)
Semiconductors
• Why are n-type and p-type semiconductors useful?
– When you put one of each together, you get a p-n
junction.
• When is a p-n junction useful?
– Only in those rare circumstances when you want
to plug something into the AC outlet in your wall!
– p-n junctions are used in rectifiers to convert AC
to DC
– They also form the building blocks of diodes,
transistors, solar cells, LEDs and integrated circuits.
p-n junction – good rectifier (converts AC to DC)
Charge buildup on p =
contact potential; prevents
further migration
Reverse bias = no
current flow through
system
Forward bias = current
flows easily
Chapter #16 – Liquids and Solids
16.1) Intermolecular Forces
16.2) The Liquid State
16.3) An Introduction to Structures and Types of Solids
16.4) Structure and Bonding of Metals
16.5) Carbon and Silicon: Network Atomic Solids
16.6) Molecular Solids
16.7) Ionic Solids
16.8) Structures of Actual Ionic Solids
16.9) Lattice Defects
16.10) Vapor Pressure and Changes of State
16.11) Phase Diagrams
Other types of solids
• While metals and networked solids can be thought of
often as one giant molecule, a few other types of solids
also exist:
– Molecular solids: covalently-bonded molecules
occupy the lattice positions and are held together in
the solid state by intermolecular forces. Examples:
ice, sulfur (S8) and white phosphorous (P4)
Molecular Solids
• The same intermolecular forces at work in liquids
exist in solids:
– London dispersion forces are fairly weak in nonpolar molecules (like CO2, I2, P4, S8), but increased
molecular weights causes many to be solids at r.t.
– Polar molecules have greater intermolecular forces
(especially when H-bonding is possible)
• These intermolecular forces are still not as strong as
the covalent bonds that hold each molecule together
as discrete units.
Other types of solids
– Ionic solids: stable, high-melting substances held
together by strong electrostatic forces between
oppositely charged ions. Examples: salt (NaCl), zinc
sulfide (ZnS), calcium fluoride (CaF2)
IONIC CRYSTALS
Ionic solid: array of B ions (not quite close-packed), with A ions in a fraction
of the holes between the B’s. A is usually (not always) the + ion and B the 
ion, since + ions are normally smaller than  ions.
Example: the two types of holes in a ccp/fcc array of B ions.
• Octahedral holes: total of 6 holes around each B ion, each shared with 5
other B ions  net of 1 hole per B (or 1 hole/cps)
• Tetrahedral holes: total of 8 holes around each B ion, each shared with 3
other B ions  net of 2 holes per B (or 2 holes /cps)
Think:
• Pack basketballs in a container, then fill holes with smaller balls
• One type of hole has 6 basketballs around it (oct), one has 4 (tet)
• Whether the holes are occupied depends on the relative sizes of the anions
and cations
Ionic Solids (NaCl)
The locations of the octahedral holes (gray x) in the
face-centered cubic unit cell
Tetrahedral Holes
(a) The location (x) of a tetrahedral hole in the face centered
cubic unit cell
(b) One of the tetrahedral holes
Tetrahedral Holes
(c) The unit cell for ZnS, S2- are closest packed, Zn2+ fill
alternate tetrahedral holes (half the tetrahedral holes are
filled)
(d) The unit cell for CaF2, Ca2+ are closest packed, F– fill
tetrahedral holes (all of the tetrahedral holes are filled)
SOLIDS AND LIQUIDS BY BOND TYPE
In solids and liquids, many atoms, molecules, or ion pairs join
together. Properties depend on bond type.
Ionic Compounds. Nearly always solids at room T, with:
• no discrete molecules, just alternating + and  ions;
• each ion surrounded by as many of opposite charge as fit.
Example: NaCl. Na+ and Cl alternate
This non-directional close packing occurs because electrical
force is equal in all directions outward from an ion’s center.
SOLIDS AND LIQUIDS BY BOND TYPE
Covalent Substances (some elements, some compounds). Can
be solids, liquids, or gases at room T.
3 types of solids/liquids:
• Molecular: Electrons held tightly within individual
molecules. Weak forces, large distances between usually
close-packed molecules. Examples: CO2, HCl, H2, Ne. For
the noble gases such as Ne, the “molecule” is a single atom.
• Metallic: “Sea” of valence e delocalized over (shared
equally among) close-packed nuclei. Examples: Mg, Cu.
• Network: All atoms interconnected by limitless web of
strong bonds, each localized between two atoms. Not
close-packed; specific e-group geometry around each
atom. Whole crystal is one giant molecule. Examples:
C(diamond), SiO2 (quartz, sand).
BOND TYPE AND PHYSICAL PROPERTIES
Structure
Room-T
Phase
Molecular g, l, or s
s
Network
Metallic
l or s
s
Ionic
MP/BP 1 Conduct
Electricity?
Mechanical
Properties
Low
Highest
Low 
High
High
Soft
Hard/Brittle
Soft  Hard/
Workable
Hard/Brittle
No
No 2
s, l 2
l,aq
1Molecular
substances melt at < 200 oC. Other substances melt
at 300 4000 oC, except certain metals: Column I, Ga, In, and
Hg [40 (Hg) to +180 oC (Li)].
2C
graphite
slightly.
and metalloids (B, Si, Ge, As, Sb, Te) conduct electricity
BOND TYPE AND PHYSICAL PROPERTIES
Element
S8
Si
Fe
Na
Bond Type
Purely Covalent (Molecular)
Purely Covalent (Network)
Metallic
Metallic
MP ( C)
113
1,414
1,538
98
Oxide
SO2
SiO2
Fe2O3
Na2O
Bond Type
Polar Covalent (Molecular)
Polar Covalent (Network)
Polar Covalent (Network)
Ionic
MP ( C)
 76
1,713
1,565
1,132