Ch 3 Lecture 1 Simple Bonding Theory
I.
Lewis Structures
A.
Lewis Bonding Model
1) Bonds between atoms consist of one or more shared pairs of electrons
2) Unshared electrons exist as lone pairs on atoms
3) Both bonding pairs and lone pairs influence:
a) Shape
b) Reactivity
4) Follows the Octet Rule
a) 8 valence electrons (s2p6) is particularly stable
b) Hydrogen prefers only 2 (1s2) valence electrons
5)
H
O
Examples:
H
O C O
H C C H
B.
Expanded Shells
1. Elements in the third or higher shells can have > 8 valence electrons
2. Empty d-orbitals are used to hold excess electrons (s2p6d10 = 18e- max)
3. Examples
C.
Resonance = ability to draw multiple Lewis structures for the same molecule
1. SO3
2.
Resonance structures are:
a. Taken as a set to represent the true structure
b. Interconverted by movement of e- only
c. Separated by double headed arrows
3.
Isoelectronic = resonance structures with identical electronic structures
a) SO3
b) CO32-
4.
Resonance structures may be electronically different as well: amides
5.
D.
O
O
H3C C NH2
H3C C NH2
More resonance structures means a lower energy for the compound.
Spreading the electrons over more atoms lets them occupy more space.
Polarity
1. Electronegativity = ability of an atom to attract shared electrons
a. Determines polarity of bonds
b. Will be discussed more later
c. Table 3-4 of your text lists electronegativities
d. Fluorine has the highest value
2.
Polar Bond = one between atoms of different electronegativities
a. Examples
H Cl
+ b.
3.
E.
C O
+

-

Use electronegativities to predict +/- parts of the molecule
Nonpolar bond = one between atoms of the same electronegativity
N N
Formal Charge: tool to evaluate resonance structures and to explain reactivity
1. Apparent charge on an atom based on its Lewis structure
2. Equation:
 valence e - on 
Formal Charge  
free atom
  unshared e -  bonds 

3.
An ion’s charge = sum of the formal charges of its atoms
4.
Uses:
a. Find best resonance structure by minimizing charge separation
b. Find best resonance structure by matching charge / electronegativity
5.
Examples:
a. Thiocyanate SCN-
b.
Cyanate OCN-
c.
Fulminate CNO-
d.
Ex. 3-1
6.
Expanded Shells can reduce Formal Charge
F.
Be and B Compounds
1. BeX2 and BX3 compounds have < 8 electrons around Be/B if single bonds
2. Be/B=X double bonds create large charge separation
3. Solids tend to have extended structures relieving charge separation/octet
4. Monomers are reactive as Lewis Acids (e- pair acceptors), even though
the small atom size allows less than an octet around Be/B
F
F
B
F
F
Be
F
II.
Valence Shell Electron Repulsion Theory (VSEPR)
A.
Using Lewis structures to predict molecular geometry
1) Electrons repel each other due to (-) charge
2) e- pairs exist due to quantum mechanical rules allowing shared orbitals
3) Molecules adopt geometries that maximize e- pairs separation
4) Steric Number (SN) = number of atoms and lone pairs around a central
atom and determines the molecules shape
a) CO2 SN = 2 Shape = linear
b) SO3 SN = 3 Shape = trigonal
B.
Procedure for using VSEPR
1) Find the Lewis structure for the molecule
2) Determine SN
3) Match SN to the appropriate geometry
C.
Figure 3-8 shows the expected geometries for molecules with no lone pairs on
the central atom
D.
1.
Origin of the square antiprism:
2.
SN = 5 and SN = 7 are not regular because there are no regular solids with
those numbers of corners
Lone Pair Repulsion
1. Lone pairs are as important as atom in VSEPR
2. Irregular structures result when the central
atom has lone pairs