1. Chapter 9 Chemical Bonding I: Lewis Theory

2. Determining the Number of Valence Electrons in an Atom
the column number on the Periodic Table will tell you how many valence electrons a main group atom has
Transition Elements all have 2 valence electrons; Why? 1A 2A 3A 4A 5A 6A 7A 8A Li Be B C N O F Ne
1 e-1
2 e-1
3 e-1
4 e-1
5 e-1
6 e-1
7 e-1
8 e-1

3. Lewis Symbols of Atoms aka electron dot symbols
use symbol of element to represent nucleus and inner electrons
use dots around the symbol to represent valence electrons
pair first two electrons for the s orbital
put one electron on each open side for p electrons
then pair rest of the p electrons

4. Lewis Symbols of Ions Li• Li+1
Cations have Lewis symbols without valence electrons
Lost in the cation formation
Anions have Lewis symbols with 8 valence electrons
Electrons gained in the formation of the anion
Li• Li+1

5. Stable Electron Arrangements And Ion Charge
Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas
Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas
The noble gas electron configuration must be very stable

6. Rules
when atoms bond, they tend to gain, lose, or share electrons to result in 8 valence electrons
ns2np6
noble gas configuration
Duet Rule: sharing of 2 electrons
E.g H2
H : H
Octet Rule: sharing of 8 electrons
Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule
E.g F2, O2

7. Exceptions many exceptions
H, Li, Be, B attain an electron configuration like He
He = 2 valence electrons
Li loses its one valence electron
H shares or gains one electron
though it commonly loses its one electron to become H+
Be loses 2 electrons to become Be2+
though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons
B loses 3 electrons to become B3+
though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons
expanded octets for elements in Period 3 or below
using empty valence d orbitals

8. Lewis Theory
the basis of Lewis Theory is that there are certain electron arrangements in the atom that are more stable
octet rule
bonding occurs so atoms attain a more stable electron configuration
more stable = lower potential energy
no attempt to quantify the energy as the calculation is extremely complex
Bonding pair: two of which are shared with other atoms
Lone pair or nonbonding pair: those that are not used for bonding

9. Electron-Dot Structures
Chapter 7: Covalent Bonds and Molecular Structure
4/26/2017
Electron-Dot Structures H • O •• H O ••
Think of this section as an introduction. It is much easier to write electron-dot structures using the rules listed in the next section.
Copyright © 2008 Pearson Prentice Hall, Inc.

10. Electron-Dot Structures
Chapter 7: Covalent Bonds and Molecular Structure
4/26/2017
Electron-Dot Structures
We have single, double, and triple bonds.
Copyright © 2008 Pearson Prentice Hall, Inc.

11. Covalent Bonding Predictions from Lewis Theory
Lewis theory allows us to predict the formulas of molecules
Lewis theory predicts that some combinations should be stable, while others should not
because the stable combinations result in “octets”
Lewis theory predicts in covalent bonding that the attractions between atoms are directional
the shared electrons are most stable between the bonding atoms
resulting in molecules rather than an array

12. Electronegativity measure of the pull an atom has on bonding electrons
increases across period (left to right) and
decreases down group (top to bottom)
fluorine is the most electronegative element
francium is the least electronegative element
the larger the difference in electronegativity, the more polar the bond
negative end toward more electronegative atom

13. Bond Polarity
covalent bonding between unlike atoms results in unequal sharing of the electrons
one atom pulls the electrons in the bond closer to its side
one end of the bond has larger electron density than the other
the result is a polar covalent bond
bond polarity
the end with the larger electron density gets a partial negative charge
the end that is electron deficient gets a partial positive charge

14. Polar Covalent Bonds: Electronegativity
Chapter 7: Covalent Bonds and Molecular Structure
4/26/2017
Polar Covalent Bonds: Electronegativity
NaCl
Cl2
HCl
Copyright © 2008 Pearson Prentice Hall, Inc.

15. Electronegativity and Bond Polarity
If difference in electronegativity between bonded atoms is 0, the bond is pure covalent
equal sharing
If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent
If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent
If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic
0.4
2.0
4.0 4%
51%
Percent Ionic Character
Electronegativity Difference
“100%”

16. Bond Dipole Moments 1 D = 3.34 x 10-30 C•m
the dipole moment is a quantitative way of describing the polarity of a bond
a dipole is a material with positively and negatively charged ends
measured
dipole moment, m, is a measure of bond polarity
it is directly proportional to the size of the partial charges and directly proportional to the distance between them
m = (q)(r)
r = radius
q = 1.6 x C
1 D = 3.34 x C•m

17. Polarity and Dipole Moment
a vector quantity from the center of the positive charge to the center of negative charge
Represents with an arrow

18. Example
Determine whether bond formed between the following pair is ionic, covalent, or polar covalent
N and O
Sr and F
N and Cl
E.g Draw the dipole moment for HF
HCl OF

19. Lewis Structures of Molecules
shows pattern of valence electron distribution in the molecule
useful for understanding the bonding in many compounds
allows us to predict shapes of molecules
allows us to predict properties of molecules and how they will interact together

20. Rules for writing Dots Lewis structures
Write the correct skeletal structure for molecule
Least electronegative atom will be in the center
Hydrogen will always be the terminal
Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule
If polyatomic ions, charges must be considered when calculating the total valence electrons
Distribute the electrons among the atoms, giving octets (or duet for hydrogen) to as many atoms as possible
If any atoms lack an octet, form double or triple bonds as necessary to give them octets.

21. Lewis Structures B C N O F use common bonding patterns
C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs
often Lewis structures with line bonds have the lone pairs left off
their presence is assumed from common bonding patterns
structures which result in bonding patterns different from common have formal charges B C N O F

22. FC = valence e- - nonbonding e- - ½ bonding e-
Formal Charge
during bonding, atoms may wind up with more or less electrons in order to fulfill octets - this results in atoms having a formal charge
FC = valence e- - nonbonding e- - ½ bonding e-
sum of all the formal charges in a molecule = 0
in an ion, total equals the charge

23. Examples
Draw a Lewis formula then assign formal charge for the following molecules and/or ions
HBr CO2
NH4+ SO32-

24. Resonance
when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures
the actual molecule is a combination of the resonance forms – a resonance hybrid
it does not resonate between the two forms, though we often draw it that way
look for multiple bonds or lone pairs

25. Rules of Resonance Structures
Resonance structures must have the same connectivity
only electron positions can change
Resonance structures must have the same number of electrons
Second row elements have a maximum of 8 electrons
bonding and nonbonding
third row can have expanded octet
Formal charges must total same
Better structures have fewer formal charges
Better structures have smaller formal charges
Better structures have − formal charge on more electronegative atom

26. Drawing resonance
Any compound for which more than one Lewis structure may be written is accurately described by no single structure. The actual structure is a resonance hydrid of them all (NOT “flipping back and forth” between resonance forms). The various structures are called contributing structure or resonance forms

27. Drawing Resonance Structures
draw first Lewis structure that maximizes octets
assign formal charges
move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge
if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond
if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. -1 -1 +1 -1 +1

28. Exceptions to the Octet Rule
expanded octets
elements with empty d orbitals can have more than 8 electrons
odd number electron species e.g., NO
will have 1 unpaired electron
free-radical
very reactive
incomplete octets
B, Al

29. Drawing Resonance Structures-1
draw first Lewis structure that maximizes octets
assign formal charges
move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge
if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond
if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. +2 -1

30. Examples
Identify Structures with Better or Equal Resonance Forms and Draw Them O3
NO2-
PO43-