1. Covalent Lewis Structures
Add up the valence electrons.
Put the element with the fewest number of atoms in the middle and surround with the remaining elements.
Draw a line from the central atom to the adjacent atoms
Subtract two electrons for each line drawn.
Make octets as needed giving the most electronegative atoms electrons first. B only needs 6, H &He 2 e’s. P, S, Cl, Argon and everything below can violate the octet rule and have more than 8 e’s if you must.
If needed form double and triple bonds.
Assign formal charges to atoms.
Formal Charge = # of valence electrons – (# of dots+ # of lines)
Assign formal charges to molecules.

2. VSEPR Beryllium chloride Aluminum chloride Methane (CH4)
Phosphine (PH3)
Water
Niobium (V) Bromide
Sulfur hexafluoride

3. Notes
Water
CO2 CO

4. BeCl2
AlCl3
CH4
PH3
H2O
NbBr5
SF6

5. C2H3O2-1
HNO3
CO3-2
PO4-3
ClO2-1
Na2SO3
CN-1
(NH4)2O2

6. SeF6
NCl3
I2S
SiO2
BF3
PI5
CCl4

7. BeCl2 (Ionic)
AlCl3 (Ionic)
CH4 (Covalent)
PH3 (Covalent)
H2O (Covalent)
NbBr5 (Ionic)
SF6 (Covalent)

8. HCl
Rb2S
PI3
SiCl4 O3
C2H2 N2 O2

9. I2O
TeO4-2
BI3
SeS2
SbCl5
NO3-1
TeCl6
AsO4-3

10. H2S
SiI4
BBr3
SiO
PCl5
NCl3
TeF6
CaO

11. TlAt3
SrCl2
PbCl4
SbI5
Rb2O
NaF
RaBr2
Al2O3

13. Steps for Drawing Lewis Structures
Add up the total number of valence electrons available (EA). Be sure to take into account the total number of atoms present.
Exception: if the compound is negatively charged, add 1 electron for each negative value
Exception: if the compound is positively charged, remove 1 electron for each negative value
Example: CH4 C= 4 VE H= 1VE = 4 + 4(1) = 8 EA
Example: CO2
Example: SO42-
Draw out a skeletal arrangement of lewis structure. Place the least electronegative atom in the center, surrounded by the remaining atoms. Draw lines to represent the bonds between the central atom and each surrounding atom.
Exception: Hydrogen can never be a central atom because it is found in the first energy level and can only have 2 electrons at most.
Draw the skeletal arrangement in the following examples
CO2
H2SO4 (Just underline the central atom)
NCl3
PF5
Count and draw in the number of electrons necessary (EN) for each to have a full octet. These are the electrons needed. Be sure to include the bonded electrons.
Subtract the number of valence electrons available (from step one) from the number of electrons needed (EN –EA). This is the number of valence electrons that you still need. We make up for these by adding double or triple bonds. If you get a negative number, you have extra electrons. These may become part of your expanded octet.
To add double or triple bonds: divide the number you get from step 4 by 2 (why? Because we are adding more bonds and each bond represents 2 electrons). The number you get is the number of bonds you need to add in the form of either double or triple bonds.
Replace necessary electron pairs with double or triple bonds, in accordance with your answer from step 5. You may need to redraw your skeleton with any necessary double bonds and repeat steps 3-5.
Example: Add in bonds for CO2
But where do they go? Double? Triple? Considering the following:
How many bonds does this atom want to make?
Octet rule
Exceptions: 3 reasons:
H: 2
Be: 4
B: 6
P: 10
S: 12
Expanded octets can be in the form of bonds or lone pairs of electrons
Avoid charges on individual atoms unless ion. The formal charge of the atom is the number of electrons before bonding – number of electrons surrounding the atom after bonding (bonds count as 1 electron).
Calculate the formal charge of each atom as written in lewis structure: CO2
Assign formal charge to entire molecule: sum of individual atoms’ formal charges. This number should add up to the original charge of the molecule.
Example: calculate formal charge of CO2
Lewis Structure Practice
Directions:
Name each molecule or ion.
Draw each Lewis structure.
Draw any resonance structures that are possible.
If there is an exception to the octet rule, circle the example and state what the exception is.
NH3
PO43-
SO42-
NO2- O3
NH4+
BH3
XeF4
ClO2
CN-
C2H4
C2H2

14. Steps for Drawing Lewis Structures Add up the total number of valence electrons available (EA). Be sure to take into account the total number of atoms present. Exception: if the compound is negatively charged, add 1 electron for each negative value Exception: if the compound is positively charged, remove 1 electron for each negative value Example: CH4 C= 4 VE H= 1VE = 4 + 4(1) = 8 EA Example: CO2 Example: SO42- Draw out a skeletal arrangement of lewis structure. Place the least electronegative atom in the center, surrounded by the remaining atoms. Draw lines to represent the bonds between the central atom and each surrounding atom. Exception: Hydrogen can never be a central atom because it is found in the first energy level and can only have 2 electrons at most. Draw the skeletal arrangement in the following examples CO H2SO4 (Just underline the central atom) NCl PF5 Count and draw in the number of electrons necessary (EN) for each to have a full octet. These are the electrons needed. Be sure to include the bonded electrons. Example: CO2 Subtract the number of valence electrons available (from step one) from the number of electrons needed (EN –EA). This is the number of valence electrons that you still need. We make up for these by adding double or triple bonds. If you get a negative number, you have extra electrons. These may become part of your expanded octet. Example: CO2 To add double or triple bonds: divide the number you get from step 4 by 2 (why? Because we are adding more bonds and each bond represents 2 electrons). The number you get is the number of bonds you need to add in the form of either double or triple bonds. Example: CO2