Transcript File

Two types of solids
– crystalline: highly ordered, regular arrangement
(lattice/unit cell)
– amorphous: disordered system
Fig. 12.26
Components of unit cell/lattice?
• At lattice points can have
– ions = ionic solid
– covalent molecules = molecular solid
– atoms = atomic solid
• 3 types depending on bonding
Metallic solid
Network solid
Group 8A solid (noble gas)
Molecular Solids
• Molecules along lattice
• Mostly covalent (either polar or np)
• Intramolecular bonding is stronger than
intermolecular
– If np, London forces only (S8, P4, CO2 (s))
– If polar, London and dipole-dipole (H2O)
• Range of IM forces, gives wide range of physical
properties (as discussed with liquids)
Ionic Solids
• Ions along lattice
• Stable, high mp
• Packing is done in a way to minimize
repulsion and maximize attractive forces
– Fixed ion position
– Very strong interionic forces
– High lattice energies, mp
– Low electrical conductivity
Atomic Solids
• Atoms along lattice
– Metallic solid
– Network solid
– Group 8A solid (noble gas)
Metallic Solids
• Metal along lattice
• Pack to minimize empty space
– packing efficiency determines things like mp
and hardness
• Delocalized electrons
– high thermal and electrical conductivity
– luster
– malleability
– ductile
Fig. 12.34
Alloys (Type of Metallic Solid)
Alloy: mixed metallic solid
Interstitial and Substitutional
Interstitial Alloys
Form between atoms of different radii, where the
smaller atoms fill the interstitial spaces between the
larger atoms.
Steel is an example in
which C occupies
the interstices in Fe
The interstitial atoms make the lattice more rigid,
decreasing malleability and ductility.
Substitutional Alloys
Form between atoms of comparable radii, where one
atom substitutes for the other in the lattice.
Brass is an example
in which some Cu
atoms are substituted
with a different element,
usually Zn
The density typically lies between those of the
component metals, and the alloy remains malleable
and ductile.
Network Solids
• In contrast to metallic solids, they
– Are brittle
– Do not conduct electricity or heat
• Classic Examples are solids of
– C: diamond and graphite
– Si: silica, SiO2 (quartz, sand) and silicates,
SiO4
Diamond
• Hardest naturally occurring substance
• Each C is surrounded by Td of other C’s
• Stabilized by covalent bonds (overlap of sp3)
Graphite
• Layers of sp2 C (fused 6 member rings)
• Conductor: unhybridized p can  bond
(resonate e- density/charge)
• Slippery: strong bonding within layers, but
weak between layers (slide)
Diamond
d = 3.51 g/mL
10 hardness
(hardest natural substance)
Colorless
No conductivity
Hfo = 1.90 kJ/mol
Graphite
d = 2.27 g/mL
<1 hardness
Black
High conductivity
Hfo = 0 kJ/mol
Conductors and Semiconductors
• Solids that have delocalized electrons
– Electrons that are free to move between HOMO
(valence band) and LUMO (conductance band)
• Examples include:
– Metallic solids and alloys
– Some network solids like graphite
– Note: ionic solids are not conductive because the
ions are in fixed positions. Only conduct when
melted or dissolved in water (ions can move
freely)
– Molecular solids tend to be nonconductive
Doping
• Can dope a semiconductor to make it
more conductive
• n-type semiconductor: increase #val e– Ex. dope Si with P: extra e- from P enters
conductance band and lowers E gap
• p-type semiconductor: decrease #val e– Ex. dope Si with Ga: Ga has 1 less e-, some
of the orbitals in valence band are empty,
creates a positive site. Si e- can migrate to
these sites
p-n junction (transistors/solar cells)
When the two come in contact, get p-n
junction. Hook n-type to (-) end of battery (or
light source) and p-type to (+) terminal, get
e- flow