Chapter 9: Liquids and Solids
Properties of Solids
Ch 13. The Chemistry of Solids
Types of Solids
TYPE
Ionic
Metallic
Molecular
EXAMPLE
NaCl, CaF2, ZnS
Na, Fe
Ice, I2
Network
Diamond
Graphite
1. Molecules, atoms or
ions locked into a
CRYSTAL LATTICE
2. Particles are CLOSE
together
3. Strong IM forces
4. Highly ordered, rigid,
incompressible
FORCE
Ion-ion
Metallic
Dipole,
Dispersion
Extended
covalent
ZnS, zinc sulfide
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Crystal Lattices
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Cubic Unit Cells
(easy to visualize; common)
•  Regular 3-D arrangements of equivalent LATTICE
POINTS in space.
•  Lattice points define UNIT CELLS
–  smallest repeating internal unit that has the symmetry
characteristic of the solid.
Units Cells for Metals
(Solid adopts structure with the lowest energy)
3 different unit cells are possible for this 2-D solid
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Chapter 9: Liquids and Solids
Metallic Structure
Metallic Solids
• 
• 
• 
• 
• 
Lattice points occupied by metal atoms
Held together by metallic bonds
Range of hardness, range of melting point
• 
In general, the more valence e-, the stronger the bonding, higher mp (Na
mp=98°C; Cr mp=1890°C)
Good conductors of electricity (mobile electrons carry charge)
Good conductors of heat (closely packed lattice vibrations and mobile
electrons transmit vibration)
nucleus &
inner shell e-
Simple Ionic Compounds
Salts with formula MX can have primitive
cubic structure — but not salts with
formula MX2 or M2X
•  Because a Cs+ ion is much larger
than a Na+ ion, the lattices of
CsCl and NaCl are different.
mobile “sea”
of valence ePLAY MOVIE
Cross Section of a Metallic Crystal:
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Use X-ray crystallography to determine
interatomic distances in a crystal
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Electrostatic (Lattice) Energy
Typical Properties of Ionic Solids
Lattice energy holds ionic solids together; it is defined as the energy
required to completely separate one mole of a solid ionic compound into
gaseous ions:
•  Hard; brittle
•  High MP
•  Poor electric conductivity
as solid, good as liquid
•  Often water-soluble
NaCl(s) → Na+(g) + Cl-(g)
Lattice Energy ∝ Q+Qr
There is a rough correlation between lattice energies and MP. The larger
the lattice energy, the more stable the solid and the more tightly held the
ions. It takes more energy to melt such a solid, so the MP is higher.
MgCl2
MgO
LiF
LiCl
lattice energy MP
2527
714
3938
2800
845
1036
610
853
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Chapter 9: Liquids and Solids
Na(s) + ½Cl2(g) → NaCl(s)
Mg2+(g) + O2-(g)
ΔHfo NaCl(s)
737 EA
Mg2+(g) + O(g)
Mg2+(g) + ½ O2(g)
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Ionic Liquids
Why does Mg form
Mg2+?
•  Ionic compound which melts below 100oC (arbitrary
definition)
•  Typically contains large organic cation, small anion; e.g.
Electron affinity
Ionization energy
IE 2180
-3916
Lattice
energy
•  Diffuse asymmetric cation leads to low melting point;
many ILs are liquid at room temperature
•  No standout applications have been developed yet; low
volatility expected to be useful, but many are toxic
Na+(g) + F(g)
Na+(g) + ½ F2(g)
Na+(g) + F-(g)
EA -328
Lattice energy
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IE 495
Mg(g) + ½ O2(g)
Mg(s) + ½ O2(g)
150
-923
Lattice
energy
Na(g) + ½ F2(g)
Na(s) + ½ F2(g)
-602 Overall energy change
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-570 Overall energy change
NaF(s)
MgO(s)
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Network Solids
Molecular Crystals
• 
• 
• 
• 
Network Covalent
Solids
Lattice points occupied by molecules
Held together by intermolecular forces
Soft; low to moderate melting point
Poor conductor of heat and electricity
•  atoms attached to its nearest neighbors by covalent bonds
•  because of the strength of the covalent bonds, typically
have very high melting points
–  generally > 1000°C
•  dimensionality (2-D or 3-D) of the network affects other
physical properties
PLAY MOVIE
–  Range of hardness
–  Generally poor electric conductivity (exceptions)
A comparison of diamond (pure
carbon) with silicon.
Naphthylacetic acid: C12H10O2
Menthol: C10H20O
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Chapter 9: Liquids and Solids
Single and Multi-wall Carbon Nanotubes
Properties of Diamond
• 
• 
• 
• 
• 
• 
• 
Graphite: 2-Dimensional Network
each carbon atoms has 4 covalent bonds to
surrounding atoms
–  sp3, tetrahedral geometry
very high melting, ~3800°C
very rigid
–  directionality of the covalent bonds
very hard
–  strong covalent bonds holding atoms in position
–  used as abrasives
electrical insulator
thermal conductor
–  best known; vibrations transmitted through lattice
chemically very nonreactive
•  Carbon atoms in a sheet are
covalently bonded together
–  forming 6-member flat rings fused together
• 
• 
• 
• 
–  each sheet a giant molecule
•  the sheets are stacked and held together by
dispersion forces
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•  ~90% of earth’s crust
•  extended arrays of Si-O
•  hexagonal crystals
•  high melting, ~3800°C
–  need to overcome some covalent bonding
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An amorphous solid does not possess a well-defined
arrangement and long-range molecular order; e.g. glass
Silicates
Properties of Graphite
sp2
similar to benzene
each C has 3 sigma and 1 pi bond
trigonal-planar geometry
–  sometimes with Al substituted for Si to
make aluminosilicates
•  slippery feel
–  because there are only dispersion forces
holding the sheets together, they can slide
past each other
–  lubricants
•  Quartz
–  Very hard
–  Melts at ~1600°C
•  electrical conductor
–  Layer of π-orbitals parallel to sheets
•  thermal insulator
•  chemically very nonreactive
Crystalline
quartz (SiO2)
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Non-crystalline
quartz glass
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Chapter 9: Liquids and Solids
Phase Diagrams
Along the lines,
phases are in
equilibrium
Solid-liquid
equilibrium (mp) not
very sensitive to
pressure changes.
The P and T axes
not drawn to scale
to emphasize
significant features
of diagram
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Phase Diagram for I2
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Phase Changes at Constant Temp
How does L→ G at
constant T, P ?
Normal mp = 113.6°C,
Tc = 113.6°C
What point on the
phase diagram
corresponds to this
picture?
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Regelation: how skaters skate (maybe)
G
70oC, 1 atm air
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↑ V
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Chapter 9: Liquids and Solids
Solid-Liquid Equilibria
In any system, if you ↑ P the DENSITY will ↑.
Therefore — as P ↑, equilibrium favors
phase with larger density
Therefore — as P goes up, equilibrium favors
phase with the larger density.
usually less dense
Actually, friction is biggest factor in making ice melt
slope of phase boundary
In any system, if you ↑ P the DENSITY will ↑.
Liquid D
Regelation: how skaters skate (partly)
H2O Solid-Liquid Equilibria
slope of phase boundary
H2O(l) D H2O(s)
1 g/cm
more dense
favored at high P
Solid
usually more dense
0.917 g/cm3
less dense
favored at low P
For most substances, solid is more dense
than liquid, so ↑ P causes liquid → solid.
Also, ↑P causes ↑ MP
The pressure exerted by the skater on ice lowers its melting point, the ice
under the skate melts, and the film of water lubricates skate blade.
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Phase diagram for H2O
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Carbon Phase Diagram
Water forms at least ten
kinds of ice, depending on
how the H2O molecules fit
together; only one form is
stable at ordinary pressures.
(a) How many triple points does carbon have?
Kurt Vonnegut’s plot in
Cat’s Cradle centered on
fictional ice-nine, which melts
at 115oF. If ice-nine comes
into contacts with the ocean,
it will act as a seed crystal
and the ocean will freeze.
(d)  Synthetic diamond can be made from
graphite. How would you do this?
(b)  Which is more dense, graphite or diamond?
(c)  Graphite is the most stable solid phase
under normal conditions. Why do diamonds
exist under normal conditions?
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