1. Lewis Structures
Two dimensional pictures of covalent species that show how the atoms are joined together with covalent bonds

2. Remember covalent bonds?
Covalent bonds occur when two or more nonmetals bond
Consists of a shared pair of electrons
Very strong bonds – large amount of energy is necessary to break them
Note the difference between covalent and ionic! Covalent is shared electrons, while ionic is taken electrons.

3. Step 1: Draw the skeleton structure
Join the bonded atoms with a single dash representing a shared pair of bonding electrons H
C O
Shows only how the atoms are linked together but you have to figure out the central atom

4. Step 1.5: Identify the Central Atom
Neither hydrogen nor fluorine is ever a central atom.
The central atom usually appears only once in the formula. For example: PF3, SO42-, CCl4
If there are two atoms that could be the central atom, most often the atom of lower electronegativity will be the central atom

5. Step 2: Total the overall valence e-
Example: H2CO C = 4 valence electrons O = 6 valence electrons H = 1 valence electron Total = 12 valence electrons

6. Step 3: Subtract two electrons for each single bond drawn in the skeleton structure
Example: H2CO 12 total valence e- - 6 e- = 6 e- H C O

7. Step 4: Distribute the remaining electrons as nonbonding electron pairs until each atom bonded to central has eight electrons (except H)
H C O
At this point, if the central atom does not have an octet, convert one or more nonbonding pairs to bonding pairs.

8. Step 5: Finalize your Lewis Structure

9. Sometimes the central atom breaks the “octet rule”!
SF6 has 48 total electrons!
The central atom must be in the third energy level (at least) in order for this to happen. The d orbitals can hold more electrons, and do. D orbitals are not present until the 3rd energy level.

10. Resonance Structures
Adds stability to a species because of multiple possibilities!

11. Which resonance structure is correct?
All three of them! Simultaneously. The true structure of this nitrate ion is actually an average of the two structures:

12. Where will you observe resonance structures?
Polyatomic ions: not always, but it is NOT uncommon for PI to have resonance structures
Some organic compounds: not always, but very possible
**Remember, only draw resonance structures when there’s more than one possible place the multiple bonds can be shown!

13. Valence Shell Electron Pair Repulsion
VSEPR
Valence Shell Electron Pair Repulsion
Basically means that the outer electrons on an atom repel each other.
All e- have negative charge, this repulsion makes sense.

14. VSEPR tells us how atoms are arranged in a molecule.
The electrons in these atoms will arrange in whatever way separates them as much as possible.

15. s, sp, sp2, sp3, sp3d, sp3d2 Hybridization
The electron pairs on each atom all exist in something called hybrid orbitals.
Consist of some combination of the s-,p-,etc orbitals that allow the molecule to achieve angles keeping the atoms far apart from each other.
s, sp, sp2, sp3, sp3d, sp3d2

16. “things” attached to central atom (bonds or lone pairs)
Determining VSEPR
How many “things” are attached to the central atom?
“things” attached to central atom (bonds or lone pairs)
hybridization 1 s 2 sp 3
sp2 4
sp3 5
sp3d 6
sp3d2

17. Polarity
Property of molecules that can be explained by their molecular shape
the atoms that make up a molecule have different electronegativities
The result?
Polar bonds – the electrons are shared unequally
When there is this “unbalance” the molecule is having a dipole moment.

19. Molecules that contain polar bonds may not be polar overall if the polar bonds cancel each other out. Their charge separations are precisely in opposite directions, therefore, a zero dipole moment!

20. The bent H2O molecule, on the other hand, IS a polar molecule
The bent H2O molecule, on the other hand, IS a polar molecule. The two polar bonds are also equal in magnitude, but do not cancel each other.

22. Sigma & Pi bonds
Sigma bonds (σ) are “end-to-end” bonding Pi bonds (π) are “side-to-side” bonding

23.
Misconception: many students may have this wrong notion that a sigma bond is the result of the overlapping of s orbitals and a pi bond is the result of the overlapping of p orbitals because they may relate the 's' to 'sigma' and the 'p' to 'pi'. However, it is seen that sigma bonds can be formed by the overlapping of both the s and p orbitals and not just s orbital.