MICSP - Resonance and Formal Charge (Lesson)
Resonance
Sometimes having just one Lewis structure for a compound is insufficient. In these cases, a resonance structure is created. Resonance is the concept that no one Lewis structure actually exists, but is instead an "average" of more than one Lewis structure. This average is known as the resonance hybrid. Molecules that exist in resonance show an increased stability. This is known as the resonance energy.
Consider the Lewis Structure for the carbonate ion:
This shows the double bond on the right. Just as easily it could show the double on the left. We could also show single bonds on the right and left with the double bond below or above. All would correctly represent the Lewis structure of the carbonate ion, but not the reality of the bonds.
Experimentally, it has been determined that ALL three C-O bonds in carbonate are identical in energy and length. There are NOT two alike and one different, but all three are identical. Additionally, none match the expected bond length for a single bond nor the shorter expected length for a double bond. All three bonds are shorter than the expectation for single bonds and longer than a double bond! What the model predicts is not the same as the experimental model.
Because all three of the resonance structures are equally appropriate, we would show resonance using a double headed arrow between these three structures to show the contributing factor of each. This indicates the real model is an average of all three.
The bond order for the carbonate ion would be 1 & 1/3 bond - three bonds each a bit shorter than a single bond but not a full short double bond.
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Formal Charge
Sometimes there are multiple ways to draw the Lewis structure for a molecule. In situations like this, chemists need a way to compare these structures. Formal charge is a bookkeeping method of showing electron distribution in a Lewis structure. Formal charge also shows exactly where the charge is located. For example, nitrate has a net charge of -1. This is the overall charge on the ion, but it does not tell us how the electrons are distributed in order to achieve this charge. And calculating formal charge will give us an idea of how reasonable a particular structure is.
Formal charge is the difference between the original number of valence electrons and the actual number of electrons used in filling the octet. The actual number of electrons is the addition of the lone pair electrons and one electron from each bond. Only one electron is counted in a bond because there are two electrons being shared in the bond but only one is being possessed by each atom. The formula is as follows:
If two or more Lewis dot structures can be drawn which satisfy the octet rule, the most stable one will be the structure where:
- The number of nonzero formal charges is minimized (ion totals will equal ion charge).
- Any negative charges are located on the more electronegative atoms.
- The sum of all the formal charges is equal to the charge of the chemical (neutral for molecules and the charge expected on polyatomic ions).
Consider these two possible Lewis structure for HClO4. Assigning formal charges to each atom in the formula gives a basis to determine the preferred Lewis structure for this compound.
Calculating Formal Charge: Perchloric Acid A
Calculating Formal Charge: Perchloric Acid B
The second structure has a negative charge on the chlorine as expected as the more electronegative element of Cl, O and H. As well the second structure assigns zero to both situations for the oxygen-carbon bonds and would the more preferred structure.
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