1. Two types of solids
crystalline: highly ordered, regular arrangement (lattice/unit cell)
amorphous: disordered system

2. Fig

3. 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)

4. 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)

6. 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

8. Atomic Solids Atoms along lattice Metallic solid Network solid
Group 8A solid (noble gas)

9. 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

10. Fig

11. Alloys (Type of Metallic Solid)
Alloy: mixed metallic solid Substitutional (a) and Interstitial (b)

12. 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.

13. 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.

14. 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

15. Diamond Hardest naturally occurring substance
Each C is surrounded by Td of other C’s
Stabilized by covalent bonds (overlap of sp3)

16. 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)

17. 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

18. 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

20. 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

21. 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