This document discusses chemical bonding and molecular structure, detailing various types of bonds such as ionic, covalent, polar covalent, and metallic bonds. It introduces valence bond theory and hybridization concepts while highlighting resonance and molecular orbital theory, showcasing how atomic orbitals combine to form molecular orbitals. The document also addresses the properties and behaviors of molecules and the role of electronegativity and charge distribution in bonding.

1. Molecular Structure and Bonding
Dr.Christoph
Phayao University June 2014

3. Part 1
What is a chemical bond ?

4. Ionic Bond
Normally between a metal and a non-metal:
They exchange electrons and become ions
(charged atoms) which attract each other by
electrostatic force.
A pair of ions does not stay alone but form crystals

5. Covalent Bond
Two non-metals share
(valence) electrons:
(Remark: Transition metals can form covalent bonds also !)

6. Polar Covalent Bond
Two non-metals share electrons unevenly because
of electronegativity difference.
Electrons are closer to one atom than the other.
This results on partially negative and positive charges on the atoms

7. Metallic Bond
Metal atoms share all
their valence electrons,
which freely move
between all atoms which
form a network.
Therefore all metals can conduct electricity and look shiny

8. Bond Polarity

9. Polar Bonds
Uneven sharing of electrons due to
differences in Electronegativity
The “pull”
an atom has
for electrons

11. Electronegativity Trends

12. Common Electronegativites
Highest value,
set to 4

14. Polar Molecules
Electrons are not equally shared in a bond,
which can lead to a dipolmoment of the whole molecule

15. Polar Bonds and Geometry

16. Which bond type ?
(exception:
Transition
metals !)

17. Electron counting

18. Formal Charge
Split all bonds in the middle
=> “real” charge on atoms
(2) Octet Rule
Count all bonding electrons for
one atom
=> 8 is most stable
(3) Oxidation Number
Give all bonding electrons to
the more electronegative atom

19. Special Cases
“Extended
octet”
Especially P and
S can use d-
orbitals to make
more than 3
resp. 2 bonds !
6 VE:
Especially
common for
B and Al !

20. Part 2: Valence Bond Theory
(VB)
“Valence Electrons are located in
bonds and lone pairs”

22. Sigma bonds

23. Pi Bonds

24. “Resonance”
Write the resonance formula for OZONE !
Does the molecule have a charge ?

26. Important exception: Carbon Monoxide !

27. Homework(2)
Draw Lewis structure(s) and find
formal charges (all atoms)
and hybridization (central atom) in:
1. NO3 (-)
2. PO4 (3-)
3. CH3 Cl
4. CH2Cl2
5. SO2
6. SO3
7. CO3 (2-)
8. H2O2
9. N2O
10. Cl O2

28. ***** Break *****

29. Prentice Hall © 2003 Chapter 9
• Atomic orbitals can mix or hybridize in order to adopt an
appropriate geometry for bonding.
• Hybridization is determined by the electron domain
geometry.
sp Hybrid Orbitals
• Consider the BeF2 molecule (experimentally known to
exist):
Hybrid Orbitals

30. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.
atomic
orbitals
hybrid
orbitals
orbital box diagrams

31. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued).
orbital box diagrams with orbital contours

33. Figure 11.3 The sp2 hybrid orbitals in BF3.

34. sp2 and sp3
Hybrid Orbitals

35. Figure 11.4 The sp3 hybrid orbitals in CH4.

36. Figure 11.5 The sp3 hybrid orbitals in NH3.

37. Figure 11.5 continued The sp3 hybrid orbitals in H2O.

38. Including d-orbitals
3d orbitals can be filled as well
=> Al acts as Lewis acid
=> P and S have “hypervalence”

39. Figure 11.6 The sp3d hybrid orbitals in PCl5.

40. Figure 11.7 The sp3d2 hybrid orbitals in SF6.

41. Acid or Base ?
Compare AlCl3 and PCl3 ?
Which acts as acid and which as base – and why ?
Why is FeCl3 a strong Lewis acid ?

43. SOLUTION:
PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO.
PLAN: Use the Lewis structures to ascertain the arrangement of groups and
shape at each central atom. Postulate the hybrid orbitals taking note of
the multiple bonds and their orbital overlaps.
H3C
C
CH3
O
sp3 hybridized
sp3 hybridized
C
C
C
O
H
H
HHH
H
sp2 hybridized
bonds
bond
C
C
C
O
sp3
sp3
sp3
sp3
sp3
sp3
sp3
sp3
sp2 sp2
sp2
sp2
sp2sp2
H
H
H
H
H
H

44. Tasks
• Draw the Lewis Structures and the Hybrid
Orbitals for Ethane, Ethene and Ethyne
(mark the hybrid orbitals)
• Which hybridization has the central atom in:
H2O, O2, NH3, NH4+, N in pyridine, O in THF,
S in SOCl2, C in HCHO compared to CO

45. Chemical Reactivity
From the hybrid orbitals we can estimate if a
molecule acts as Lewis acid or base
(if there is an electrophilic or nucleophilic center)
Consider the “empty” pz orbital of C in HCHO vs. the
“filled” sp orbital of C in CO
-> in the first case, it acts as Lewis acid, in the second
as base !

46. ***** Break *****

47. VSEPR

48. VSEPR Theory
Clip:
http://www.youtube.com/watch?v=nxebQZUVvTg

49. MO Theory

50. The Central Themes of MO
Theory
A molecule is viewed on a quantum mechanical level as a collection of nuclei
surrounded by delocalized molecular orbitals.
Atomic wave functions are summed to obtain molecular wave functions.
If wave functions reinforce each other, a bonding MO is formed (region of
high electron density exists between the nuclei).
If wave functions cancel each other, an antibonding MO is formed (a node of
zero electron density occurs between the nuclei).

51. Amplitudes of wave
functions added
Figure 11.14
An analogy between light waves and atomic wave functions.
Amplitudes of
wave functions
subtracted.

52. Prentice Hall © 2003 Chapter 9
Molecular Orbitals
• Molecular orbitals:
• each contain a maximum of two electrons
• have definite energies
• can be visualized with contour diagrams
• are distributed over the whole molecule
(not only in between 2 atoms)
• When two AOs overlap, two MOs form.

53. Prentice Hall © 2003 Chapter 9
Molecular Orbitals
The Hydrogen Molecule

54. Prentice Hall © 2003 Chapter 9
Figure 11.15 The MO diagram for H2.
Energy
MO
of H2
*1s
1s
AO
of H
1s
AO
of H
1s
H2 bond order
= 1/2(2-0) = 1
Filling molecular orbitals with electrons follows the
same concept as filling atomic orbitals.

56. Prentice Hall © 2003 Chapter 9
Electron Configurations and Molecular
Properties
• Two types of magnetic behavior:
• paramagnetism (unpaired electrons in molecule):
strong attraction between magnetic field and
molecule;
• diamagnetism (no unpaired electrons in molecule):
weak repulsion between magnetic field and
molecule.
• Magnetic behavior is detected by determining the mass
of a sample in the presence and absence of magnetic
field:

57. Diatomic molecules

58. The energy level
is the lower, the
higher the EN of
the atom is !

60. Naming of MO’s: example O2 molecule
“g” = symmetric to C axis
“u” = anti-symmetric

61. Diatomic molecules
Consider the EN of each atom – the higher the EN,
the lower is the energy of the orbitals !
The highest filled MO is called “HOMO”, the lowest
unoccupied MO “LUMO”
-> check example CO
http://firstyear.chem.usyd.edu.au/calculators/
mo_diagrams.shtml

62. Example CO
HOMO
LUMO
“lone
pair” on C

63. Chemical Reactivity
Important are the HOMO and LUMO (“frontier orbitals”)
http://www.meta-synthesis.com/webbook/12_lab/lab.html

64. GroupOrbitals

65. Construction of Group Orbitals – example H2O

66. Interaction 1: in-phase H orbitals

67. Interaction 2: out-of-phase H orbitals

68. Indicate different MO types: (bonding, non-bonding. anti-bonding)

69. Combination of 3 H
orbitals to
3 group orbitals
BH3 molecule

70. Compare HOMO/LUMO to BH3 !
=> what is an acid / base ?

71. Homework (3)
Number Molecule
1 CN
2 CN(-)
3 BC
4 BN
5 BO
6 BF
7 CF
8 NO
9 NO (+)
10 NO (-)
Number Molecule
11 NF
12 OF
13 CH4
14 BH3
15 SbF6
16 XeF2
17 XeF4
18 XeF6
http://firstyear.chem.usyd.edu.au/calculators/mo_diagrams.shtml