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
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 ?
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
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
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 !
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
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 ….. )