Covalent Bonding, The Octet Rule and Multiple Bonds
Ionic substances exist as a metal cation and a non-metal anion that are attracted to one
another. Covalent molecules occur when non-metals combine with non-metals and electrons
are shared in order to lower energy. The process of forming a covalent bond can be
summarized as follows:
1. Two atoms approach one another as each has a high electron affinity
2. The electron density in each atom begins to shift towards the nuclei of the other atom
3. There is now a much greater possibility of finding an electron in the region between the 2
atoms
4. Each of the atoms appears to have a share of the two electrons between them
5. A balance between the two electron repulsions and the two nuclear repulsions occurs. This
creates a lower energy situation in a delicate balance between repulsion and attraction
Every covalent bond has 2 important characteristics:
1. The average distance between the nuclei (bond length)
2. The energy required to separate the 2 nuclei and create 2 neutral atoms (bond energy)
NOTE: when a bond is formed, energy is released. When a bond is broken, energy is
required. This will be discussed in the thermochemistry section.
Covalent bonds can be shown as 2 dots between atoms or as a dash
Example:
H· + H· → H : H or H – H
The Octet Rule
We know from electron configurations that noble gases are very stable and low energy. When
ions form, electrons tend to be gained or lost until a noble gas configuration is reached. The
same is true for the number of electrons shared in covalent molecules. The octet rule states
that when atoms react, they tend to achieve a valence shell having eight electrons.
Drawing Lewis Diagrams of Covalent Compounds
Remember: A Lewis Diagram consists of the element’s symbol with dots surrounding it,
symbolizing the valence electrons.
1. The central atom is usually the least electronegative.
2. Place other atoms equally spaced around the central atom (H is often attached to O, not
the central atom).
3. Count the number of valence electrons. This is the number of electrons that must be in
your final diagram. Add or remove electrons for charge.
4. Place a pair of electrons between each pair of bonding atoms (i.e., one pair of electrons
per bond).
5. Finish the octets of all atoms except the central atom.
6. Count the number of electrons in the diagram and compare to the total from Step 3.
Add any extra electrons as pairs on the central atom.
7. Check the central atom for an octet. If it is short, make multiple bonds with surrounding
atoms to complete its octet. Recall:
- B is happy with 6 electrons
- Be is happy with 4 electrons
- O will make up to three bonds if they are all single
- N won’t make more than 3 bonds
8. The central atom can have more than 8 electrons if it has available empty d orbitals (i.e.,
Si or larger atom).
Draw Lewis Dot Diagrams for the following:
1. Methane (CH4)
2. Ammonia (NH3)
3. Water (H2O)
Multiple Bonds
Bonding that we have encountered so far only have single bonds. There are however many
covalent molecules in which more than one pair of electrons are shared between 2 atoms.
Such an example would be nitrogen gas, N2:
CO2 works in the same fashion (remember: we drew it the first week!):
Practice - Draw Lewis dot diagrams for the following:
a) LiBr
b) K2S
c) PH3
d) F2
e) Na3N
f) Al2S3
g) CaS
h) BaO
i) O2
j) SiO2
k) CS2
l) BF3