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 Substitutional (a) and Interstitial (b)

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