Selection of Lewis Structures
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Lewis structures are only theoretical explanations and predictions based of observations.
The distance between nuclei of a bond is called a bond length, and the energy need to split the bond is called bond energy.
Bond lengths and energies are different for single, double, and triple bonds.
This is necessary for the obervation of how atoms are bonded.
For example, the using the rules in the previous section, this drawing of sulfuric acid would be:
But based on observation, the bonds between the S and the O are actually shorter than regular single bonds.
Sulfur-oxygen double bonds are shorter than single bonds.
Therefore the real structure is like this:
It doesn't seem to follow Lewis rules, but sulfur has a valence shell that has 3s, 3p, and 3d subshells, which can accomodate more than 8 electrons.
Sometimes, while drawing Lewis structures, you don't have observational data to use.
Formal charges can be used to find out what to choose.
Formal charge = (number of electrons in valence shell of isolated atom) - (number of bonds to the atom) - (number of unshared electrons)
For example, lets take the first incorrect drawing of sulfuric acid.
Formal charge on S = 6 - 4 - 0 = 2
Formal charge on H = 1 - 1 - 0 = 0
Formal charge on O = 6 - 2 - 4 = 0 (on the ones bonded to H)
Formal charge on O = 6 - 1 - 6 = -1 (on the isolated ones)
We can disregard the ones with 0 formal charge.
The ones that do have formal charge are the sulfur and the isolated oxygens.
Since the 2 oxygens are -1 and sulfur is 2, another bond goes from each oxygen to the sulfur to cancel the formal charge.
(2 + 2(-1) = 0)
When several Lewis structures are possible, those with the smallest formal charges are the most stable and are preferred.