(taken in part from Lowry and Richardson)

Lewis Model of Bonding-most elementary bonding model in common use; based on idea that ionic bonding forces arise from electrostatic interaction between ions of opposite charge; covalent bonding forces arise from sharing of electron pairs

Valence shell occupancy (electrons) -total # of noncore electrons in the immediate neighborhood of atom; should not exceed 2 for hydrogen; 8 for second row elements (Li -F) = octet rule; for third row and later, may exceed 8 (e.g. P and S)

Formal Charge is given by: (valence electrons for an isolated atom)-(electron ownership)

*valence electrons for an isolated atom: 1 for H, 4 for C, 5 for N, 6 for O, 7 for F, etc. (this is the number of electrons in the highest energy shell for the isolated atoms)

*electron ownership = # of unshared electrons (non-bonding electrons = lone pair electrons) + 1/2 of all electrons in bonds leading to the atom being considered

Total charge of molecule = sum of all formal charges

Rules for writing Lewis Dot Structures (a notation used to show electron distribution about atoms-example given for diazomethane: H2CN2)

1. Count the total number of valence electrons contributed by the electrically neutral atoms. If the species being considered is an ion, add one electron to the total for each negative charge; subtract one electron for each positive charge.

(overall charge is zero so no electrons added or subtracted)

2. Write the elemental symbols for the atoms and fill in the number of electrons determined in step 1 as dots around the atom. The electrons should be added so as to make the valence-shell occupancy of hydrogen 2 and the valence shell occupancy of other atoms 8 wherever possible. (note that in some cases, an atom will necessarily have to have less than 8)

3. Valence shell occupancy must not exceed 2 for hydrogen and 8 for a second-row atom; for a third-row atom it may be 10 or 12.

4 . Maximize the number of bonds, and minimize the number of unpaired electrons always taking care not to violate Rule 3.

5. Find the formal charge on each atom using the equation given above.

H=>valence electrons =1 and each H has 1 [[sigma]] bond therefore, 1-1=0; no formal charge on H

C => valence electrons = 4 and it has 3 bonds ([[sigma]]) to it (1/2 x 6 e's = 3) and one lone pair (non-bonding pair= 2e's) therefore, 4-5= -1; C has a -1 charge in this resonance structure

N(on left) => valence electrons= 5 and it has 2 [[sigma]] and 1 [[pi]] bond to it (1/2 x 6 e's = 3) & 1 lone pair (non-bonding pair of e's = 2 e's) therefore, 5-5 =0; no formal charge on N (on left)

N(on right) => valence electrons = 5 and it has 1 [[sigma]] bond and 1 [[pi]] bond (1/2 x 4 e's = 2 e's) & 1 lone pair (2e's) therefore , 5-4 = +1; N(on right) has a +1 charge in this resonance structure

Resonance Structure- the superposition of two or more Lewis dot structures with different electron distributions but identical nuclear positions into a composite or hybrid structure

*resonance structures must adhere to the octet rule (atoms in the second row cannot exceed 8 e's in their valence shell but in some cases may have less than 8 e's)

*[[sigma]] bonds are not altered between resonance structures only [[pi]] e's and non-bonding e's (i.e. valence e's) are "moveable" in resonance structures

*hybridization of atoms does not change between resonance structures

Resonance structures of diazomethane that satisfy the above rules:

Arrows in Organic Chemistry