Types of Solids* Intermolecular Force(s) Between Particles
1. Metallic Crystals (Metals)
Examples: Na, Cu, Fe, Mn
Metallic bonding: Valence electrons form mobile
sea of electrons which comprise the metallic
2. Ionic Crystals (Ionic Solids)
Examples: NaCl, MgCl2, MgO
Ionic Bonding: Attraction of charged ions for one
another. Lattice energy is a measure of ionic bond
3. Covalent Crystals (Network
Examples (small class!):
C(diamond), SiC(s), SiO2 (quartz)
Network covalent bonding. Network solids are
extremely hard compounds with very high melting
and boiling points due to their endless 3-
dimensional network of covalent bonds.
4. Molecular Crystals
Examples: One or more of the following:
(a) Need H bonded to O, N or F:
H2O, HF, NH3.
(a) Hydrogen bonding: Hydrogen bonds are
weaker than covalent bonds, but stronger than (b)
or (c) below.
(b) C6H6 (benzene), polyethylene,
I2, F2, and all the compounds from
(b) Dispersion forces (induced dipole – induced
dipole or London dispersion forces): universal
force of attraction between instantaneous dipoles.
These forces are weak for small, low-molecular
weight molecules, but large for heavy, long, and/
or highly polarizable molecules. They usually
dominate over (c) below.
(c) CHF3, CH3COCH3 (acetone)
and H2O, HF, NH3.
(c) Dipole-dipole forces: these forces act
between polar molecules. They are much weaker
than hydrogen bonding.
Note: Van der Waals Forces is a category which includes both categories (b) and (c)
5. Atomic Crystals
Examples: He, Ne, Ar, Kr, Xe
Dispersion forces: See Section 4(b) above.
*Note: Many of the compounds given as examples are not solids at room temperature.
But if you cool them down to a low enough temperature, eventually they will become