Two types of solids

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

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Fig. 12.26
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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)

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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)

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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

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Atomic Solids

• Atoms along lattice
• Metallic solid
• Network solid
• Group 8A solid (noble gas)

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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

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Fig. 12.34
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Alloys (Type of Metallic Solid)

• Alloy mixed metallic solid
• Interstitial and Substitutional

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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.

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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.

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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

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Diamond

• Hardest naturally occurring substance
• Each C is surrounded by Td of other Cs
• Stabilized by covalent bonds (overlap of sp3)

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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)

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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 lt1 hardness Black High
conductivity ?Hfo 0 kJ/mol
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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

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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

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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