+ – 0 0
H Cl H H
The basic units: ionic vs. covalent
• Ionic compounds form repeating units.
• Covalent compounds form distinct molecules.
• Consider adding to NaCl(s) vs. H2O(s):
Na Cl Na Cl O
Cl Na Cl Na O
• NaCl: atoms of Cl and Na can add individually
forming a compound with million of atoms.
• H2O: O and H cannot add individually, instead
molecules of H2O form the basic unit.
I’m not stealing, I’m sharing unequally
• We described ionic bonds as stealing electrons
• In fact, all bonds share – equally or unequally.
• Note how bonding electrons spend their time:
H2 H H HCl H Cl LiCl [Li]+[ Cl ]–
0 0 + – + –
(non-polar) polar covalent ionic
• Point: the bonding electrons are shared in each
compound, but are not always shared equally.
• The greek symbol indicates “partial charge”.
• Recall that electronegativity is “a number that
describes the relative ability of an atom, when
bonded, to attract electrons”.
• The periodic table has electronegativity values.
• We can determine the nature of a bond based on
EN (electronegativity difference).
• EN = higher EN – lower EN
NBr3: EN = 3.0 – 2.8 = 0.2 (for all 3 bonds).
Basically: a EN below 0.5 = non-polar covalent,
0.5 - 1.7 = polar covalent and above 1.7 = ionic
Determine the EN and bond type for these: HCl, CrO,
Br2, H2O, CH4, KCl
HCl: 3.0 – 2.1 = 0.9 polar covalent
CrO: 3.5 – 1.6 = 1.9 ionic
Br2: 2.8 – 2.8 = 0 covalent
H2O: 3.5 – 2.1 = 1.4 polar covalent
CH4: 2.5 – 2.1 = 0.4 covalent
KCl: 3.0 – 0.8 = 2.2 ionic
Holding it together
Q: Consider a glass of water. Why do molecules
of water stay together?
A: there must be attractive forces.
forces are much
Intramolecular forces occur Intermolecular forces occur
between atoms between molecules
• We do not consider intermolecular forces in ionic bonding
because there are no molecules.
• The type of intramolecular bond determines the type of
Inter-molecular Forces (IMF’s)
• London Dispersion Forces
• Dipole – Dipole Forces
• Hydrogen Bonding
London Dispersion Forces
• London dispersion forces result from
instantaneous non-permanent dipoles created by
random electron motion. London dispersion forces
are present in all molecules and are directly
proportional to molecular size.
• LDF’s are the weakest of the IMF’s
Dipole – Dipole Forces
• Dipole-dipole interaction is the
attraction between a partially
negative portion of one molecule
and a partially positive portion of a
• Dipole-dipole interaction occurs in
any polar molecule as determined by
• Hydrogen bonding is the unusually
strong dipole-dipole interaction that
occurs when a highly
electronegative atom (N, O, or F) is
bonded to a hydrogen atom.
• Hydrogen bonding is the strongest
of the 3 IMF’s
Effects of IMF’s on Physical Properties
• The strength of intermolecular forces present in a substance
is related to the boiling point and melting point of the
substance. Stronger intermolecular forces cause higher
melting and boiling points.
• CH4 - Methane: has only very weak London dispersion
forces (lowest b.p. & m.p.)
• CHCl3 - Chloroform: has dipole-dipole interaction (moderate
b.p. & m.p.)
• NH3 - Ammonia: has hydrogen bonding and dipole-dipole
interaction (high b.p. & m.p.)
Name / Formulate the following Compounds
1. PbS 1. Lead (II) Sulfide
2. HClO4 2. Hydrogen Perchlorate
3. Zn(OH)2 3. Zinc Hydroxide
4. HBr(aq) 4. Hydro Bromic Acid
5. (NH4)2CO3 5. Ammonium Carbonate
6. sulfur hexafluoride 6. SF6
7. nitrous acid 7. HNO2(aq)
8. hydrogen chloride 8. HCl
9. lead(II) chloride 9. PbCl2
10. zinc sulfate 10. ZnSO4