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

2. Chemical Bonds (4 Types)
Ionic
electrostatic attraction of ions (charges)
Covalent (Molecular)
sharing of e–‘s
Covalent Network
shares e–‘s throughout
Metallic
metal atoms bonded by a sea of e–’s
Diamond
Quartz

3. Ionic Bonding

4. Energetics of Ionic Bonding
As we learned previously, it takes +495 kJ/mol to remove electrons from sodium.
ionization energy
electron affinity
We release –349 kJ/mol back by gaining electrons by chlorine.

5. Why is it so exothermic?
+495
–379
+106

6. Energetics of Ionic Bonding
There must be a third piece to the puzzle:
+energy
electrostatic attraction between Na+ cation and the Cl– anion.
Does getting closer release or absorb energy?
attract
There must be a third piece to the puzzle.
–energy

7. Lattice Energy This third piece of the puzzle is… lattice energy:
The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.
(OR energy released when ions attract/combine)
Coulomb’s law:
F : Force of Electrostatic Attraction
q : charge
d : distance between charges
F = k
q1q2 d2

8. Lattice Energy HW p. 337 #22 _________ charge, and ___________ size.
Lattice energy increases with:
_________ charge, and ___________ size.
increasing
decreasing
more q
less d
F = 
q1q2 d2

9. Covalent Bonding

10. Covalent Bonding share electrons
several electrostatic interactions in these bonds

11. Electronegativity
-ability of an atom to attract electrons when bonded with another atom.
most electronegative

12. Polar Covalent Bonds sharing electrons, but NOT always shared equally.
the more electronegative fluorine end of the molecule has more electron density.

13. Polar Covalent Bonds q– q+
When two atoms share electrons unequally, a bond dipole results.
The dipole moment, , of two charges separated by a distance, r, is:
 = qr q– q+

14. Polar Covalent Bonds
The greater the ∆EN, the more polar the bond.

15. CO2 Polarity We have discussed bond dipoles, but…
…polar bonds in a molecule do NOT demand a polar molecule.
CO2

16. Polarity Overall dipole moment sum of individual bond dipoles =
overall dipole moment for a molecule
Overall dipole moment

17. Polar Bonds—Polar Molecule?
HW Chp 8

18. Lewis Structures
Lewis structures are representations of molecules showing ALL bonding and nonbonding valence electrons.

19. Drawing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
For anions (–), add e– ’s.
For cations (+), subtract e– ’s.
PCl3
(7) = 26

20. Drawing Lewis Structures
The central atom is the least electronegative (never H).
Connect the outer atoms by single bonds.
Keep track of the electrons:
26  6 = 20

21. Drawing Lewis Structures
Fill the octets of the outer atoms.
Keep track of the electrons:
26  6 =  18 = 2

22. Drawing Lewis Structures
Fill the octet of the central atom.
Using Steps 1-4, draw a Lewis structure for HCN.
Keep track of the electrons:
26  6 =  18 = 2 2  2 = 0

23. Drawing Lewis Structures
If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.

24. 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)
Ox #’s describe charges IF electrons were lost or gained IF the compound was 100% ionic.
FC describes charges IF all atoms shared perfectly equally (same electronegativity).

25. Drawing Lewis Structures
Use FC to select the best Lewis structure ONLY when directed…
…the fewest charges.
…a negative on the most electronegative atom.

26. Drawing Lewis Structures
Draw Lewis structure for these molecules:
H2O
CHCl3
NH3
NH4+
F2 O2 N2 HF
VOLUNTEERS
please!
HW Chp 8 #46
Bonus: PCl5
Super Double Bonus: PCl6–

27. Resonance
Draw a Lewis structure for ozone, O3. O O O
Which?
Neither!

28. Resonance The observed structure of ozone:
…both O—O bonds are the same length.
…both outer O’s have a same e– density.

29. Resonance
Multiple resonance structures are required for equivalent structures when…
…one Lewis structure cannot accurately represent a molecule.

30. Resonance Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.
How many double bonds?
How many single bonds?
NONE!
NONE!

31. Resonance Consider HCO2–
the electrons that form the double bond can move among the two O’s and the C.
They are delocalized.

32. Resonance
In benzene (C6H6), a hexagon with a circle signifies the delocalized electrons in the ring.
HW Chp 8 #51ac

33. Exceptions to the Octet Rule
There are two types of ions or molecules that do not follow the octet rule:
Ions or molecules with less than an octet.
Ions or molecules with more than eight valence electrons (an expanded octet).

34. Less Than Octet Consider BF3:
A filled octet places a – on the B and a + on F.
Not an accurate picture of the true distribution of electrons in BF3 .

35. Less Than Octet
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

36. Less Than Octet
The lesson is: Do NOT fill the central atom’s octet if it gives a positive FC on a more electronegative outer atom.

37. Expanded Octet
Can expand the octet of atoms on the 3rd period or below if needed.
PCl5

38. Expanded Octet
phosphate ion (PO4–3)

39. Expanded Octet phosphate ion (PO4–3) OR
In terms of Formal Charge, the better structure is…
b/c it minimizes the FC’s.

40. Expanded Octet phosphate ion (PO4–3)
HW Chp 8 #62, 64
phosphate ion (PO4–3)
The lesson is: ONLY expand octets for period 3, 4, 5… OR when directed to consider Formal Charge.

41. Molecular Shapes shape affects reactivity
number of bonding and nonbonding e– pairs predicts the shape of a molecule

42. 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.”

43. What Determines the Shape of a Molecule?
e– pairs repel
assume the e– pairs are as far apart as possible, then predict the shape.

44. Electron Domains Each e– pair is one e– domain.
All (single, double, or triple) bonds, count as only one e– domain.
How many e– domains?

45. Electron-Domain Geometry
count e– domains in the Lewis structure to predict the e– domain geometry.
AB2 AB3 AB4 AB5 AB6

46. Molecular Geometry
e– domain geometry is not always the shape of the molecule.
molecular geometry (actual shape) is defined by the positions of only the atoms, not the nonbonding pairs. (B’s not E’s)

47. 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)

48. Molecular Geometry

50. Bond Angles & Electrons
Nonbonding pairs are larger than bonding pairs.
Therefore, their repulsions are greater; this decreases bond angles in a molecule.

51. Bond Angles & Multiple Bonds
Double and triple bonds place greater electron density than single bonds.
Therefore, they also affect bond angles.

52. Large Organic Molecules
It makes more sense to discuss the geometry about a particular atom rather than the geometry of the molecule as a whole.

53. Large Organic Molecules
WS 8b
(react at particular sites)

54. Overlap and Bonding When bonds/attractions form, energy is _________.
released
What repulsive forces?
What attractive forces?
Where is energy being released?
Where is energy being absorbed?

55. Valence Bond Theory
covalent bonds form by sharing electrons when orbitals on the two atoms overlap.

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

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

58. Single Bonds single bonds are  bonds  overlap is greater
stronger bond (stable).

59. Multiple Bonds  bonds are weaker (easier to break)
double & triple bonds:
1 strong  bond, & 1 or 2 weak  bonds

60. Hybrid Orbitals Consider beryllium (Be):
In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.
But if it absorbs a small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.

61. Hybrid Orbitals
combining the s and p orbitals yields two degenerate (same energy) orbitals as hybrids of the two original orbitals.
2 sp hybrid orbitals

62. Hybrid Orbitals
This 180 bond angle is consistent with the observed linear geometry of beryllium compounds.
BeF2

63. Hybrid Orbitals For beryllium (Be) we get… ground state promote
hybridize
2 degenerate sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

64. Hybrid Orbitals Using a similar model for boron (B) leads to…
ground state
promote
hybridize

65. Hybrid Orbitals
…3 degenerate sp2 orbitals.

66. Hybrid Orbitals
With carbon (C) we get…
ground state
promote
hybridize

67. …4 degenerate
sp3 orbitals.

68. 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.

69. Multiple Bonds The triple bond in ethyne, H–C≡C–H :
2 sp hybrid orbitals 1 C–C  bond a 1 C–H  bond, and 2 pairs of unhybridized p orbitals overlap to form 2  C–C bonds.

70. Molecular Orbital Theory
Valence bond theory and hybridization explain bond angles, lengths, and shapes like the 4 equal bonds in methane, CH as 4 equal sp3 hybrid orbitals rather.
But…some aspects of bonding are better explained by another model called Molecular Orbital (MO) Theory.
Molecular orbitals contain electrons associated with an entire molecule rather than the overlap of individual atomic orbitals of each atom.

71. Bond Order total # of bonds Bond Order = # of bonding domains
What is the bond order of these molecules: 2
1.5
1.3