
Bonds can be classified as being either polar or
non-polar.

Polarity: tendency of a molecule, or compound, to be attracted
or repelled by electrical charges because of an asymmetrical
arrangement of atoms around the nucleus.

Think of it like a game of tug of war, if one end of the
compound is pulling on the electrons more than the other,
there is an unequal pull, and therefore, the substance is polar.
If there is an equal pull, then the substance is non-polar.

This concept of polarity is determined by electronegativity.
Ionic Bond
Attraction between oppositely charged ions
 Occurs when electrons are transferred from one ion
(charged particle) to another
 Electronegativity difference 1.7+
 Metals react with Nonmetals to form ionic
compounds


Always
Polar !!!
Lewis Dot Structure of Ionic
Compounds
 KCl
 CaBr2
 KNO3
 (NH4)3PO4
Properties of Ionic Compounds





Hard
Good conductors of electricity in liquid or
aqueous form only, because ions can move
in solution and in liquid form, but not in
solid form.
High melting and boiling points
Solid at room temperature
Dissolve in polar substances: like water.
(Polar – opposite charges).
Covalent Bonds
Formed when 2 atoms (both nonmetals) share electrons.
[Example Cl2 or H2O]
 Neither atom pulls strongly enough to remove an electron
from the other
 The EN difference is < 1.7
 Unpaired electrons pair up in such a way that the atoms
complete their outer shells
 Covalent compounds
also referred to as molecular
compounds

Properties of Covalent Bonds
Gases, liquids or solids
 Soft
 Nonmetals
 Poor conductors of heat and electricity because
they are not charged particles. (No ions or mobile
electrons)
 Low melting and boiling points because of weak
attraction between molecules.

Polar vs. Non-Polar Covalent Bonds



Unlike an ionic compound, a covalent compound can be
classified as either a polar covalent bond, or a non-polar
covalent bond.
If the EN of the atoms are different then it is a polar
covalent bond.
If the EN of the atoms are the same or very similar then it is
a non-polar covalent bond.
0.0 - 0.4 = non-polar covalent
0.5 - 1.6 = polar covalent
Number of Covalent Bonds



Single covalent bond: one pair of shared electrons; 2
electrons total
-Double covalent bond: two pairs of shared electrons; 4
electrons total
-Triple covalent bonds: three pairs of shared electrons; 6
electrons total
Must know how to draw these!!
HF
HCl
HBr
HI
H2O
H2S
CO2
CS2
SiO2
SiS2
CH4
CF4
CCl4
CBr4
CI4
NH3
PH3
Partially Positive & Negative

In a polar covalent bond, both of the elements are
non-metals, and therefore there is no “true” + or –
charges; instead there are partially (+) and partially
(-) charges.
The element with the
higher EN is partially (-)
and the one with the
lower EN is
partially (+)

This is a SNAP!
•Symmetric are
•Nonpolar
•Asymmetric are
•Polar
Other Types of Covalent Bonds
Network Solids:



Solids that have covalent
bonds between atoms linked
in one big network or one big
macromolecule with no
discrete particles. This gives
them some different
properties from most covalent
compounds.
They are hard, poor
conductors of heat and
electricity, and have high
melting points
Examples include: Diamond
(C), silicon carbide (SiC), and
silicon dioxide (SiO2)
Metallic Bond

Occurs only in metals (Example Copper)

Metals have low ionization energy meaning they hold onto
their valence electrons very loosely

As a result the electrons in metallic substances move about
very easily and are not associated with any particular atom

Therefore, the particles of a metal are usually positive ions

The attraction between the positive cations and the moving
electrons is what holds the metal together

Properties of Metallic Bonds are that of metals: hard, good
conductors of heat & electricity, malleable, ductile, etc . . .
surrounded by a mobile sea of electrons
Dipole-Dipole Attractions
• Positive end of a polar molecule is
attracted to the negative end of an
adjacent polar molecule.
Hydrogen Bonding
• An intermolecular attraction between a hydrogen
atom in one molecule to a nitrogen, oxygen, or
fluorine atom in another molecule
• The strongest intermolecular force
• Substances with hydrogen bonds tend to have
much higher melting and boiling points than
those without hydrogen bonds
• Example: The boiling point of H2O is much
higher than H2S
van der Waals Forces
Weak intermolecular forces between
non-polar molecules (like diatomic molecules)
 Dispersion forces make it possible for small, nonpolar molecules to exist in both liquid or solid
phases under conditions of high or low
temperatures.
 Increases with molecular size, Ex. As you go down
group 17, dispersion forces increase and boiling
point increases.

Molecule-Ion Attraction


Attraction between the ions of
an ionic compound such as
NaCl, and a molecule such as
water (or any other polar
covalent compound).
When you put NaCl into
water, the Na+ from the salt is
attracted to the O from the
water which is partially (-),
and the Cl- from the salt is
attracted to the H+ of the
water.