Lewis structures are another way to represent molecules. Lewis Structures were introduced by Gilbert N. Lewis in 1916. Lewis suggested the use of lines between atoms to indicate bonds, and pairs of dots around atoms to indicate lone or non-bonding pairs of electrons.
In the example above, 3 hydrogen atoms with one valence electron each form three bonds with one nitrogen atom with 5 valence electrons. By forming three bonds, nitrogen gains 3 electrons to make a total of 8 surrounding it. This satisfies the octet rule allowing nitrogen's valence shell of electrons to look just like the noble gas neon's. The hydrogens on the other hand gain one electron each in the formation of the bonds and thus their valence shell now appears like heliums. The unused pair of electrons are assigned as a lone pair to the nitrogen forming a stable molecule of ammonia (NH3)
As demonstrated in the example above, the guiding principle behind the formation of Lewis structures is the fulfillment of the octet rule: all atoms would like to be surrounded with an octet of electrons. Of course, there are, some exceptions: very small atoms (H, Be and B) have less than an octet, and some main group atoms in the third period and below (P, S, Cl, Br, and I) may have more than an octet but most elements still strive for the completion of their outer (valence) shell with 8 electrons. Drawing correct Lewis structures takes practice but the process can be simplified by following a series of steps:
How to Draw Lewis Structures in 5 Easy Steps
Gregory Hodgkins (YouTube)
Step 1. Count all the valence electrons for each atom. Add or subtract electrons if the structure is an anion or cation, respectively.
Example
SO42-
Sulfate is a polyatomic ion with 1 sulfur (6 valence electrons), 4 oxygens (4 x 6 valence electrons = 24 e-) and a charge of -2 (2 valence electrons). If we add all the electrons together we get 32 valence electrons with which to make bonds and lone pairs around the atoms in the ion.
Step 2. Determine which atoms are bonded to one another. Draw a skeletal structure.
Example
SO42-
A rule of thumb is that the central atom (atom which all the other atoms will be bound) is the one furthest to the left or bottom in the periodic table. Sulfur and oxygen are in the same group but sulfur is below oxygen so it will most likely be the central atom in the structure:
How do we know that all of the oxygen atoms are bound to the sulfur and none to each other? Well, oxygen only needs to share two electrons to complete its octet so it really does not want or need to make more than one bond. Also, the structure shown is the most symmetrical way to build the molecule and with a few rare exceptions, nature tends to make molecules in a symmetrical manner.
Step 3. Connect the atoms with a pair of electrons in each bond. Subtract the number of bonding electrons from the total number of valence electrons.
Each line in the drawing above represents the use of 2 electrons. Knowing we started with 32 and have now used 8 leaves us with 24 electrons to distribute.
Step 4. Add electron pairs to complete octets for all peripheral atoms attached to the central atom. Beware of hydrogen – hydrogen never has more than one bond or one pair of electrons.
Using the remaining 24 electrons we need to add them to the oxygen atoms until each has an octet.
Step 5. Place remaining electrons on the central atom, usually in pairs. The octet rule may be exceeded for P, S, Cl, Br, or I.
In this case there are no remaining electrons as we have used them all and each atom has an octet so this is a good Lewis structure as shown.
If the ion had been SO32- (sulfite) instead of sulfate, then the structure would have had a lone pair of electrons on the central atom:
The structure of PCl5 is a good example of a molecule that exceeds the octet:
Phosphorus is located in Period 3 (3rd row) of the periodic table and thus is capable of exceeding an octet.
Step 6. If the central atom does not have an octet, form double or triple bonds by moving lone pairs from one or more peripheral atoms to form bonds with the central atom and achieve an octet.
In the case of sulfate or sulfite we do not need to make double bonds, but another example in which a double bond would be necessary is C2H4. In this example, several rules come in to play: the most symmetrical structure is best and hydrogen cannot be bound to more than one atom (form 1 bond). Knowing these rules makes it easier to determine that the two carbons must be bound to each other and the hydrogens distributed symmetrically around them. But let’s first run through the proper steps to make this Lewis structure:
Number of valence electrons: 2 carbons (2 x 4 valence electrons= 8e-) and 4 hydrogens (4 x 1 valence electron = 4 e-) added up gives us 12 valence electrons to work with:
Counting the bonds we have made shows that we have used up 10 of the 12 valence electrons leaving us with 2 to complete the octets. Hydrogen has made its one bond and is sharing two electrons (like Helium) so it does not need any more electrons, but both carbons currently only have 6 which means they both need another 2 electrons to reach their octet. So we need 4 electrons and only have 2 to work with. How do we solve this problem? SHARE MORE. If we add another bond between the two carbons (a double bond) they will share 4 between them and have another 4 bonded to the hydrogen atoms and therefore will have 8; the octet that they need.
Double and even triple bonds are formed when there is a lack of electrons needed to make full octets for all of those atoms that need them. We don’t add double or triple bonds when we have more than enough electrons, only when we run short.
Step 7. Add brackets to ions to indicate their charge
Going back to our sulfate ion example, notice that the Lewis structure we drew does not in any way indicate to the viewer that it is a charged molecule:
This is a problem because as a knowledgeable chemist you would automatically realize that there are 2 more electrons forming this molecule than can be accounted for by the atoms represented. In other words if we just count the valence electrons for the sulfur and oxygen shown we get a total of 30 valence electrons but we are showing 32 in the Lewis structure. In order to avoid confusion, we have to indicate in some manner where we get the other two electrons. We do this by adding brackets around the ion and showing the charge:
Let's Practice: