1. Unit 2 (Chp. 8,9): Bonding & Molecular Geometry
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Unit 2 (Chp. 8,9): Bonding & Molecular Geometry
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

2. Energetics of Ionic Bonding
IE & EA & ??
(3rd factor)
+energy
electrostatic (Coulombic) attraction between Na+ ion & Cl– ion.
Forming attraction release or absorb energy?
attract
Coulomb’s law:
F = k
q1q2 d2
There must be a third piece to the puzzle.
–energy
F : force of attraction
q : charge
d : distance between

3. Lattice Energy: _________ charge, and ___________ size. increasing
energy released when a mole gaseous ions attract to form a solid ionic compound.
Lattice energy increases with:
_________ charge, and ___________ size.
increasing
decreasing
(q)
(d)
q , d
stronger attraction
E = 
q1q2 d
 energy

4. Covalent Bonding
share electrons
covalentbond

5. Electronegativity attraction Zeff , shield
-ability of an atom to attract electrons when bonded (sharing e–’s) to another atom.
attraction
Zeff , shield
most electronegative

6. Polarity of Covalent Bonds
the more electronegative F has more e– density. d+ d– F F H F
NONpolar (DEN ~ 0)
Polar (DEN > 0.4)
(shared equally)
(shared UNequally)
dipole:
partial charges (d+, d–) on atoms
that are sharing e–’s UNequally
dipole moment,  :
a measure of polarity with direction

7. Polar Covalent Bonds The greater the ∆EN, the more polar the bond. H IBr H Cl H F

8. H2O Molecular Polarity Overall dipole moment
a molecule is polar ONLY IF:
has polar bonds (DEN)
arranged asymmetrically
(do not cancel out)
H2O
Polar
Overall dipole moment

9. Drawing Lewis Structures
In 5 easy steps:
PCl3
Count (val e–’s)
Connect (bonds)
Octet (outer dots)
Octet (central dots)
(if necessary)
form Multiple Bonds to fill central octet

10. Drawing Lewis Structures
Draw a Lewis structure for CO2 –1 +1 or
Formal Charge
= (VE’s) – (NBE’s) –(½BE’s)
FC = (val e– ’s) – (dots) – (lines)
NOTE:
FC: charges IF all atoms shared e–’s equally (same electronegativity) IF 100% covalent.

11. Drawing Lewis Structures
Use FC to select the best Lewis structure ONLY when directed…
…the fewest & smallest charges.
…a “–” on the most electronegative atom. –2 +1
Which of the Lewis structures for NCO– is the most significant in terms of formal charge? –1 –1

12. Resonance Multiple resonance structures are required if…
…bonds can not be clearly identified between identical atoms due to delocalized e–’s.

13. Less Than Octet
For BF3:
A filled octet places a – on the B and a + on F.
A double bond between B and F is worse than leaving B with only 6 electrons.
Principle: Do NOT fill B’s octet if it gives a +FC on a more electronegative outer atom.

14. Expanded Octet phosphate ion (PO43–) OR
The better structure minimizes the FC’s.
Principle: ONLY expand octets for atoms in periods > 3 as needed when…
directed to consider Formal Charge.

15. Valence Shell Electron Pair Repulsion Theory (VSEPR)
“The best arrangement of a given number of electron domains is the one that minimizes repulsions among them.”

16. Molecular Geometry Must know basic shapes: AB2 linear
AB3 trigonal planar
AB4 tetrahedral
AB5 trigonal bipyramidal
AB6 octahedral
AB4
AB3E
AB2E2
all variations have 1 or more E’s (e– pairs)

17. Two Types of Orbital Overlap
Sigma () bond overlap forms:
head-to-head overlap
single bonds
Sigma
(single)

18. Two Types of Orbital Overlap
Pi () bond overlap forms:
side-to-side overlap.
double & triple bonds. Pi
(p orbitals)

19. Hybridization (of orbitals)
Consider carbon (C):
promote
hybridize
…4 degenerate
sp3 orbitals.
ground state

20. Multiple Bonds
C sp2 orbital forms  overlap with O sp2 orbital and with H s orbital.
unhybridized p orbitals of C & O form  bond overlap.

21. = MO Theory & Bond Order total # of bonds Bond Order =
# of bonding domains
What are the bond orders of these molecules: 2
1.5
resonance
delocalized p e–’s =