Lewis Dot Structures

Using Lewis dot structures and the octet rule, we can predict the electronic structure of covalently bonded molecules.

Diatomic Molecules

These molecules are called homonuclear diatomic molecules, since all atoms are the same. They are also non-polar molecules.

 Below is an example of the polar HCl molecule.

Chlorine is more electronegative than Hydrogen so HCl has a polar covalent bond.

 For diatomic oxygen, the Lewis dot structure predicts a double bond.

Actually, this is not quite right. While the Lewis diagram correctly predict that there is a double bond between O atoms, it incorrectly predicts that all the valence electrons are paired (i.e. it predicts that each valence electron is in an orbital with another electron of opposite spin). We know this is not right because substances that contain unpaired electrons exhibit a behavior called paramagnetism. Paramagnetic substances are attracted to a magnet. For example, Iron metal is paramagnetic, and it can be picked up with a magnet. On the other hand, blackboard chalk is not paramagnetic, so it cannot be picked it up with a magnet. When we pour liquid nitrogen through the poles of a magnet, we see that it just passes through. Nitrogen is not paramagnetic, thus it doesn't have any unpaired electrons.

 On the other hand, if we pour liquid O₂ through the poles of the magnet, a solid O₂ bridge forms because liquid O₂ is paramagnetic. In the next chapter, we will learn a more advanced theoretical approach to chemical bonding called molecular orbital theory that can predict double bonds and unpaired electrons for O₂. In general, the Lewis-dot structures have the advantage that they are simple to work with, and very often present a good picture of the electronic structure.

 For a diatomic Nitrogen, the Lewis-dot structure is:

This triple bond between nitrogens is very strong. The strength of the this triple bond makes the N₂ molecule very stable against chemical change. It is often called an inert gas.

There is a relationship between the number of shared elelctron pairs and the bond length.

The distance between bonded atoms decrease as the number of shared e^- pairs increase.
 

Rules for Lewis Dot Structures

1. Count the number of valence e^- each atom brings into the molecule.

 

 
 
 

For ions, the charge must be taken into account.

 
 
 
2. Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule).

 
 

One exception is H (surrounded by only 2e^-s).  
 

Sometimes it's necessary to form double and triple bonds. Only C, N, O, and S (rarely Cl) will form multiple bonds.

 
 
 

Exceptions to the Octet Rule

3. If there is not enough electrons to follow the octet rule, then the least electronegative atom is left short of electrons.

 e.g. BeF₂ number of valence e^- = 2+ 2(7) = 16 e^- or 8 pairs.

Neither Be or F form multiple bonds readily and Be is least electronegative so

 
 
 
4. If there are too many electrons to follow the octet rule, then the extra electrons are placed on the central atom.

 

 
 
   
 
 

How can this happen?

The octet rule arises because the s and p orbitals can take on up to 8 electrons. However, once we reach the third row of elements in the periodic table we also have d-orbitals, and these orbitals help take the extra electrons.

NOTE: You still need to know how the atoms are connected in a polyatomic molecule before using the Lewis-Dot structure rules.