Gilbert Newton Lewis (surrounded by pairs of electrons)
Hybrid orbitals, Lewis diagrams, the octet rule
Review
We can visualize a covalent bond as a balance between attraction and repulsion
A second way to view the formation of a bond is to consider orbital diagrams
Video (15 min - stop at HCl)
Now we can visualize atoms bonding because of the lower energy achieved via the stability of filled valence shells
The bottom line: atoms gain, lose, or share electrons to obtain noble gas electron configurations
Video: filling subshells
Video: filling subshells
Draw orbital diagrams for F + F, H + O, Li + F
Overlapping orbitals
Draw orbital diagrams for F + F, H + O, Li + F
F2
H2O
LiF is ionic (metal + non-metal)
Hybrid orbitals
Two overlapping orbitals form what is known as a hybrid or molecular orbital
Just as in a s,p,d, or f orbital the electrons can be anywhere in the orbital (even though the electron has started out in one atom, at times, it may be closer to the other nucleus)
Each hybrid orbital has a specific shape (described in chapter 8)
You do not need to know shapes
You need to know that hybrid orbitals exist and that they are formed from overlapping orbitals
Lewis diagrams
Read Pg. 229. Draw Lewis dot diagrams for Ne, Sb, Rb, F. How many variations of the Lewis diagram for P can be drawn?

Ne
Sb
Rb
F
P
P
P
P
Lewis diagrams follow the octet rule: atoms when forming ions, or bonding to other atoms in compounds have 8 outer electrons
Q - How can the octet rule be explained?
A - s (2 e–) and p (6 e–) orbitals are filled
Ionic bonding
Recall: Ionic bonding involves 3 steps:
1) loss of e-, 2) gain of e-, 3) +ve, -ve attract
3)
This can be represented via Lewis diagrams…
Read remainder of 7.2 (pg. 230)
Diagram the reaction between Li + Cl and Mg + O (PE 3)
The octet rule (ionic compounds)
Draw Li + Cl and Mg + O (PE 3, pg. 230)
Note also that the charge on an ion can be determined by the number of places removed from a noble gas (Ca, N, Al?)
Ca2+, N3-, Al3+
See pg. 230 for example of Ca + Cl
Covalent bonding
Covalent bonds can also be shown via Lewis diagrams - E.g draw Lewis diagrams showing the combination of 1) H+Cl, 2) C+Cl, 3) H+O, 4) Mg+F, 5) N+H, 6) Do PE 4 (pg. 234)

HCl
CCl4
H2O
MgF2 - Ionic
NH3
Note bonds can also be drawn with a dash to represent two electrons (read 234-5 upto 7.5)
The octet rule (covalent compounds)
(Read 7.4 up to PE4 (pg. 232 - 4). Do PE4.)
Read remainder of 7.4
Although the octet rule works for most compounds, PCl5 and SF6 are examples of exceptions (see pg. 235 for structures)
Also, it can be difficult to draw the correct bonds for atoms with multiple bonds
Some clear rules have been established for drawing Lewis structures
We will see that the rules dictate molecular structure, which dictates shape, which can influence chemical properties of a compound
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