This document provides an overview of chemical bonding concepts including the four main types of bonds (ionic, covalent, metallic, and polar covalent) and their properties. It also discusses electronegativity and bond polarity, Lewis structures, the octet rule, valence shell electron pair repulsion theory (VSEPR) and molecular geometry. Hybridization of orbitals is explained using examples of sp, sp2, and sp3 hybridization. The document concludes by noting some exceptions and limitations of the valence bond theory for more complex molecules and transition metals.

1. Chemical Bonds and Properties
What we will learn:
• 4 types of bonding and their properties
• Electronegativity and polar bonds
• Valence Bond Theory
• VSEPR
• Hybridization of Orbitals
• Sigma- and pi-bonds
• Oxidation Numbers
• Formal Charge
• Resonance

2. How to study effectively ?

4. O3
• What different kind of electrons are in this
molecule ?
• Is this molecule stable ?
• Does this molecule have a charge ?
• Is this molecule linear or bent ?
• Is the bond strength higher, the same or lower
than in O2 ?

5. H2O
• Is this molecule linear or bent ?
• How many different kinds of electrons are in
this molecule ?
• Why is this molecule a liquid at RT, but H2S is a
gas ?
• Why is this molecule more stable than
Hydrogenperoxide ?

6. CO
• Is this molecule stable ?
• Is this molecule more or less reactive than CO2 ?

7. • Is this molecule stable ?
• Is this molecule planar ?
• Why is this molecule a weak acid ?

8. Part 1
What is a chemical bond ?

9. Concept Map

10. After this lesson, we should understand:
• Valence electrons
• Covalent and ionic bond
• σ- and π-bonds, lone pairs
• Lewis Structure
• Electronegativity
• Dipole moment
• Oxidation number
• Formal Charge
• Octet Rule
• Hybridization
• Basic shapes of simple molecules

11. 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

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

13. 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

14. 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

15. 4 Types of Bond

16. Bond Polarity

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

19. Electronegativity Trends

20. Common Electronegativites
Highest value,
set to 4

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

23. Polar Bonds and Geometry

25. Which of these molecules have the
greatest dipole moment ?

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

29. Electron counting

30. (1) 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

31. 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 !

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

35. Sigma bonds

36. Pi Bonds

37. The VB Theory works well for diatomic molecules.
But for more complex molecules we need an extension !

38. VSEPR
Valence
Shell
Electron
Pair
Repulsion
Lewis 2D  3D structure

39. Most common geometrical structures

45. Draw 3D from Lewis Formula
• CO2 vs. SO2
• O2 vs. O3
• H2O vs. H2O2
• SOCl2 vs. HCHO

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

48. Important exception: Carbon Monoxide !

49. Part 3: Hybridization

50. Why another theory ?

51. Hybrid Orbitals
All orbitals in an atom involved in sigma-bonds
hybridize (mix) into orbitals of equal energy.
Pi-Bonds are still formed by p-orbitals.
Lone electron pairs count as “single bonds” and
are part of the hybridization !

52. sp3 Hybridization

54. sp2 Hybridization

56. sp Hybridization

57. Which hybrids form and when ?
+ lone pairs

58. 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

59. Inclusion of d-orbitals
Elements in row 3 and up, hybridization can
include also d-electrons.
Typical example: SF6 with 6 sigma bonds

60. sp3d

61. Special Case: Transition Metal compounds
Bonds in transition metal compounds are either
ionic like in FeCl3
but can be covalent as well as in Fe(CO)5
=> the VB Theory is not very suitable to explain
the bonding for transition metals !
This can be done with the “Crystal Field Theory”
(soon to come ….. )