Chapter 8 sections 1-7
Chapter 9 sections 1-6
Ionic – Formed by the electrical attraction of a
positive cation to a negative anion.
Don’t worry about “Energetics of Ionic Bond Formation” for now!
Nonpolar Covalent – A bond between atoms
with equal (or nearly equal)
electronegativities (the electrons are shared
equally – neither atom has a slight charge)

Polar Covalent – A bond between atoms with
different electronegativities (the one with the
higher electronegativity pulls the electrons more
and thus that atom is slightly negative and the
other atom slightly positive)



The ability of an atom in a molecule to attract
electrons to itself.
Generally increases as we move across a series
and generally decreases down a group.
(Fluorine has the highest!)
The larger the difference in electronegativity
values between two atoms, the more polar the
bond between them will be.

Determination of Bond Type
Step 1: Check for a metal - if there is a metal the bond is
usually IONIC
Step 2: If there is no metal subtract the electronegativity
values (bigger – smaller)
If the difference is:
0.4-ish or less then the bond is NONPOLAR
0.5-ish or greater then the bond is POLAR
Electronegativity
Difference
The bond is:
Example
(difference)
0-0.4
Nonpolar
Cl-Cl
(0.0)
0.4-1.0
Polar
H-Cl
(0.9)
1.0-2.0
Very polar
H-F
(1.9)
>2.0
Ionic
NaCl
(2.1)
Ex: Si & N
3.0 – 1.8 = 1.2
very polar


The atom with the higher electronegativity is
partially negative
The atom with the lower electronegativity is
partially positive

Polar Molecules are also called “dipoles”
 Si-N 



Dipole Moment is a quantitative measure
of the magnitude of a dipole…the higher
the dipole moment the more polar the
molecule.


A way to determine which Lewis Structure
is the best representation of a molecule
The best one has atoms with formal charges
closest to zero.

Valence Bond
Theory – a
covalent bond
is an overlap
of atomic
orbitals that
allows two
electrons to
share a space
in both atoms.

Sigma (σ) bond – the end-to-end overlap of
orbitals.

All single bonds are sigma bonds

Pi (π) bond – A side-to-side overlap of
orbitals


Double bond – 1 sigma & 1 pi
Triple bond – 1 sigma & 2 pi


Consider methane
(CH4)
The carbon makes
4 bonds, but there
are only 2 available
orbitals!
↑↓
1s
↑↓ ↑ ↑ __
2s
2p


1 of the electrons
from the 2s
sublevel jumps to
the 2p sublevel.
Now there are 4
orbitals available
for bonding!
↑↓
1s
↑_
2s
↑ ↑↑_
2p

One of these orbitals is an “s” orbital and
the other 3 are “p” orbitals….when
scientists look at methane, all the bonds
look the same. They become sp3 hybrid
orbitals.
Bonds + Lone Pairs
around the central
atom
Common
hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
*Multiple bonds count as one!!


Electrons are spread out over a number of
atoms in a molecule rather than localized
between a pair of atoms.
Often happens in molecules with
resonance.