Resonance & Formal Charge

Resonance & Formal Charge

Resonance & Formal Charge Chapter 8 part 5 Resonance What if more than

one valid Lewis dot structure is possible? Consider Nitrate ion. Nitrogen bound to 3 oxygen atoms, one with a double bond. But is that the only possible structure that obeys the octet rule? Possible structures for NO3 -

Another illustration The readers of the AP do not recognize this method of writing the resonance structure, but I think it is important to be familiar with it.

More Examples Why the Resonance model ? A word about bond

A word about bond length: Experimentally it has been shown: Single bonds are longer than double bonds Double bonds are longer than triple bonds. But

The Resonance model explains this Since there is more than one correct position for the double bond, the model for resonance allows that double bond to shift locations. The

resonance bond is shorter than a single bond and larger than a double bond. Question Which of the two resonance structures has the shorter bond length and Why?

Odd Electron Molecules Relatively few molecules formed with nonmetals contain odd

numbers of electrons. These examples usually involve N. These are not very stable. Formal Charge

The difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule. Need to determine the valence electron of the free atom. Need to determine the number of valence electrons assigned to the atom in the molecule. Formal charge

Assigning electrons: Lone pairs and single electrons on the atom each count for that atom. Bonded electrons are shared and therefore half go to each atom bound.

Subtract this number from the valence number. The result and its sign is the atoms formal charge. Example In The Sulfate ion,

what are the possible Lewis dot structures? Which one is the most likely? Use formal charge. Final example: Give the possible Lewis dot structures for XeO3. Which one is the most likely given the formal

charge? Summary 1. 2. To calculate the formal charge of an atom: Take the sum of the lone pair electrons and one half the shared

electrons of the atom in the molecule. Subtract the number of valence electrons on the free neutral atom. Summary continued The sum of the formal charge of all atoms in a given molecule or ion must equal the overall charge of that species.

If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule.

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