Unit 3: Bonding and Nomenclature
The entirety of bonding is to fulfill the octet rule as well as possible. That is, all elements want
to have a full valence shell of electrons (0 or 8, and 2 if its the first row).
At this point you should know that there are a certain number of valence electrons and
oxidation numbers (charge) associated with each group. They follow the following pattern.
1
2
3
4
5
6
7
+1
+2
+3
±4
−3
−2
−1 0
H
8
He
Li
Be
B
C
N
O
F
Ne
Chemical Bonding

Ionic Bonds: atoms give up or gain e– and are attracted to each other by coulombic
attraction
Ionic compounds = salts

Metal + Nonmetal

Electronegativity difference > 1.7
Properties of Salts
1. very hard –
2. high melting points –
3. brittle –
For Lewis Dot Diagrams, simply move an electron from the cation to the anion.
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
…atoms share e– to get a full valence shell
Covalent Bonds
Lewis structure:
1. Two shared e– make a single covalent bond, four make a double bond, etc.
2. unshared pairs: pairs of unbonded valence e–
3. Each atom needs a full outer shell, i.e., 8 e–.
Exception: H and He need 2 e–
covalent compounds = molecular compounds
-- have lower melting points than do ionic compounds
 good insulators
 can be found as a gas/liquid/solid at room temperature
For Lewis Dot Diagrams you need to connect the lone pairs of electrons.

In metals, valence shells of atoms overlap, so v.e–
Metallic Bonds
are free to travel between atoms through material.
Properties of Metals
Lots of free electrons
Good conductors of heat and electricity
Malleable – can be turned into a sheet
Ductile – can be turned into a wire

Other Types of Bonds

dipole-dipole forces

hydrogen bonds ← The only important one in the entirety of this course as it holds our
DNA together.

London dispersion forces

ion-dipole forces
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To draw Lewis Dot Structures/Diagrams
For starters drawing LDD's is based upon the number of Valence electrons. You can find
them based on the column that they are in. The only real concern is that you follow the rule of
ONE electron on each side, before placing the SECOND. [Similar to electron box and Hund's
rule]
For example:
H has one v.e
.
H
..
. N.
.
N has 5 v.e.
There is only one exception to this rule and that is He which has 2 v.e → He:
When doing Ionic bonds draw the LDDs of the individual atoms
Salt (NaCl)
Na.
..
: Cl :
.
Note that Chlorine is close to fulfilling the octet rule if it had one more electron, so it will take
the electron from Na.
..
Na
: Cl :
..
Now we have Na+1 and Cl-1
When doing Covalent bonds draw the LDDs of the individual atoms and connect the LONE
electrons. For simple ones this works. In some cases, if there are not enough electrons you
can break apart a pair of e- to make new bonds.
Water (H2O)
SF6
H.
..
.O:
.
H.
F
F
F
S
F
F
F
O
H
H
Note how S has 6 v.e and 6 bonds
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To draw Lewis Dot Structures/Diagrams
1. Determine the type of compound.
a) Figure out if it is ionic or covalent.
2. Identify the arrangement of the atoms
a) Begin with the central atom.
b) Place the remaining terminal atoms around the central one.
c) Count the total number of valence electrons for all of the atoms.

Add or subtract the number of electrons based on the charge.
d) Draw a bond between the central atom and all exterior ones.
e) Place electron pairs around all terminal (exterior atoms) to complete octets

except H since it holds 2 v.e

If there are unused electron pairs you may place them on the central atom.
f) Count the total number of electrons and bonds (each is 2 e- ) to be sure you have
the right number.
g) Check the CENTRAL atom to see if there is the correct number of objects coming
out of it and rearrange accordingly.
Let's take SO42Before going on, these are between nonmetals and nonmetals → Covalent, so we can draw bonds.
Steps a/b
O
O
S
O
Step c
S=6
4O = 4x6 = 24
-2 = 2 more
Total = 32 v.e
O
Step d
Step e
O
|
O– S–O
|
O
..
:O:
..
| ..
:O – S – O :
..
| ..
:O:
Step f
There are 32
4 bonds = 8
24 electrons
4
..
Step g
S has 6 v.e so one needs to make sure that there are 6 bonds or electrons attached to S. But we are out of
electrons! So instead we take some of the pairs of electrons and turn them into a bond instead.
:O:
Step f
.. || ..
There are 32
:O – S = O :
4 bonds = 8
..
|
24 electrons
:O:
..
In this case I took a pair of electrons from the top and right most oxygens.
To check your answer make sure each outside atom has the right number of valence electrons (8 usually).
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