Lewis Structure

The Lewis structure is named after Gilbert N. Lewis, who developed the concept in 1916. The idea was later introduced and named by Irving Langmuir in 1920, who gave credit to Lewis. Lewis structures are also known as electron dot structures.

The Lewis structure only shows valence electrons. A more accurate representation of the truly shared bonding electrons can be drawn by using a ‘resonance’ structure that offers both kinds of electrons (Lewis and resonance).

Both structures represent the same molecule, but each is a little more complete than the other. The first one represents only the valence electrons, while the resonance hybrid also shows some bonding electrons. Their choice depends on which effect is more critical for the molecule under consideration.

In chemistry, it is conventional to use one structure for neutral molecules and another for ionised molecules or have an electric charge because of an imbalance of charge between atoms (or molecules) in them. You may have noticed that the Lewis dot structure is similar to the Lewis structure. The only difference is that in a Lewis structure, you show the valence electrons (the number of electrons in the outer shell of an atom). You leave out all nonbonding electrons in a Lewis dot structure, showing only the valence electrons.

How to draw Lewis Diagrams?

The nitrogen atom has five valence electrons, and we know from the chart that it needs three more electrons to achieve a full octet. The oxygen atom has six valence electrons, and we know that it requires two more electrons to complete a full octet. We can see from the chart that nitrogen will give up three valence electrons, and oxygen will gain two valence electrons by forming a triple bond between the two atoms.

Let us draw a Lewis structure for NO3

First, we need to calculate the number of valence electrons in a molecule. It can be calculated by adding the group number of all the atoms and their charge. After that, we would have obtained the number of valence electrons.

To draw the skeletal structure for the molecule, draw a rectangle and divide it into three equal parts. The central atom is bonded to two of the others by single bonds. The central atom is bonded to the other by a double bond. The molecule can be drawn with all single bonds or one double bond and one single bond, which would mean two lone pairs were not used in bonding. Let us now have a look at what the structure would be looking like:

After that, the next step is to add lone pairs of valence electrons to the structure. We have calculated that NO3 has 24 valence electrons. Of these, 24, 6 have been used to make the complete skeletal structure. You can add lone pairs of electrons to the terminal atoms. Keep adding until their octet is done, or the other case is when you run out of electrons. The first option, adding lone pairs to the atoms, is called resonance, and it is what we’ve been doing in these lessons. The second option, running out of electrons, is called a formal charge. Formal charge tells you how many extra electrons are needed for an atom to have its octet.

The octet rule requires that the central atom have eight electrons in its outer shell. If you run out of electrons before forming bonds, you have to create multiple bonds to get it done. The number of multiple bonds varies on what type of element you’re dealing with. Here, we can see that N is short with 2 electrons. Now, we can use one lone pair from the atom of O, located on the left. Hence, it will form a double bond and help complete the bond.

Now, we need to determine the formal charge of the molecule. The formal charge of each atom in a molecule is significant because it indicates the polarity of the molecule and the direction of bonding. All atoms with a positive formal charge are electrophiles and are usually the attacking species in electrophilic addition reactions. Atoms with a negative formal charge (indicated by a minus sign) are nucleophiles and are traditionally the attacking species in nucleophilic substitution reactions.

We cannot complete the Lewis structure without the formal charges. The things which have been needed in this are as follows:

• The fewest number of atomic charges as possible. In short, we need to have 0 formal charges for as many items as possible.
• It is necessary to have the formal charges the same as the atom’s electronegativity. It means more negative charges on the electronegative atoms and less positive charges on the electronegative atoms.

The formal charge is the charge it would have if all its bonds were utterly ionic. If you have a -1 and +1 charge close together, you can often rearrange things so that the charges cancel out. One way to do this is to use the lone pair of electrons on the -1 atom to form a double bond with the nearest +1 atom. This makes a double bond between them instead of a single bond, and there is no longer a net charge.

The Lewis structure for NO3:

You have determined that the “best” structure involves a double bond between the nitrogen and one of the oxygens (shown in red). That is fine, but there are several ways to display this information. 

Conclusion

In this material, we have discussed the concept of Lewis structure and how it is used. We also discussed the idea of skeletal structure, calculating formal charges, and much more.