Intermolecular Forces
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
bond.
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
strength.
3.
Covalent Crystals (Network
Solids)
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
(a) above.
(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)
above.
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
solids.
