This document provides an overview of bonding theories including molecular orbital theory and valence bond theory. It discusses concepts like electronegativity, formal charges, hybridization, sigma and pi bonds, and the formation of molecular orbitals from the combination of atomic and group orbitals. Examples are provided to illustrate these concepts for molecules and coordination complexes like NH3, H2O, and octahedral Co(NH3)6 2+. The document is intended to help readers understand bonding models and be able to analyze bonding in different substances.

1. Bonding Theories
Advanced Inorg. Chem.
Dr. Chris Sontag
University of Phayao
Oct.2016
1

2. After this lesson, we should understand:
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3. 3O3
1. Is this molecule stable ?
2. Does this molecule have a charge ?
3. Is this molecule linear or bent ?
4. Is the bond strength higher, the same
or lower than in O2 ?
3

4. 2 2H2O2
1. Is this molecule linear or bent ?
2. How many different kinds of electrons are in
this molecule ?
3. What is the oxidation number of O in this
molecule ?
4

5. CO
1. Is this molecule stable ?
2. Is it polar or non-polar ?
3. Is this molecule more or less reactive than CO2 ?
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6. 6

7. Electronegativity (EN)
The amount of EN difference determines the
polarity of the bond:
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8. 8

9. 9

10. Allred-Rochow EN
Example:
Flourine (r = 72 pm)
Carbon (r = 77 pm)
Calculate the AR
electronegativity
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11. Shielding of 2p electrons:
Flourine:
S = 6 * 0.35 + 2 * 0.85 = 3.8 => Z* = 9 – 3.8 = 5.2
EN = 3590 * 5.2/(72 2) + 0.744 = 4.35
Carbon:
S = 3 * 0.35 + 2 * 0.85 = 2.75 => Z* = 6 – 2.75 = 3.25
EN = 3590 * 3.25 / (77 2) ) + 0.744 = 2.71
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12. VALENCE ELECTRONS AND
LEWIS STRUCTURES
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13. The periodic table
is built up so that
elements with the
same number of VE
are in one column
= no. of VE
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14. Electron configuration
do NOT use core electrons !
Practise: write the configuration for:
1. Fe2+
2. Pb
3. W
4. Ti4+
Write only the
Valence
Electrons !
http://www.slideshare.net/Hoegler6/09-lecture14

15. Examples:
(1) Atoms and ions
Try yourself:
1. Al and Al3+
2. F and F-
3. K and K+
4. H and H-
Indicates that O
has 2 valences
(can make 2
bonds)
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16. 16

17. Lewis Molecules
Write the atom with the LOWEST EN in the middle !
Example
(2) In molecules write all VE
for each atom
Try yourself:
1. AlH3
2. LiH
3. SiF4
4. C2H6
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18. 18

19. Formal Charges
In some cases, VE cannot be arranged without creating charges
Each atom in a molecule has a “formal charge”:
Count the electrons that belong to this atom and compare to the
VE in the element
N has only
4 electrons,
1 is missing
O has 7
electrons, 1 more
than in oxygen
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20. Multiple Bonds and formal charges
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21. Formal charges and
Oxidation numbers
Oxidation number:
assign all electrons in bonds to
the atom with higher EN !
-> N has no el. -> ox no. +5
Formal Charge:
split all bonding el. Between the
atoms and count the remaining
-> N has 4 el. -> formal charge is +1
Find the same for: CO2 HCHO H3C-OH HCOOH
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22. Transition metal compounds
Example:
FeCl3 = ionic compound (“salt”)
But in water: Fe(3+) (H2O)6 + 3 Cl(-)
Lewis Formulas do not reflect the bonding in
coordination compounds:
Fe(3+) has 5 valence electrons, but forms 6 bonds !
Oxidation numbers:
can go from -1 to +7, normally +2 or +3
Find the numbers for: KMnO4, MnO2, K4Fe(CN)6, Fe(CO)5
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23. 23

24. VB THEORY
(DIAGRAMS FROM:
HTTP://WWW.SLIDESHARE.NET/HOEGLER6/10-LECTURE)
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25. VB Theory and
molecular geometry
http://www.slideshare.net/Hoegler6/10-lecture 25

26. 26

27. 27

28. VB Theory = Hybridization of atomic orbitals
But NOT ALWAYS:
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29. 29

30. Sigma - bonds 30

31. How many sigma-bonds can each atom form ?
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32. Try the same for FORMALDEHYDE HCHO 32

33. 33

34. Try the same for FORMALDEHYDE HCHO
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35. 35

36. sp3d hybridization
Example: PCl5 compared to PCl3 – both molecules are stable !
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37. sp3d2 hybridization
Example: SF6 compared to SCl2 – both molecules are stable !
six
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38. Coordination Compounds
For TM ions the VE are
counted ALL as d-electrons !
EMPTY metal orbitals are needed
to be filled with ligand-electrons !
We can form a d2 sp3 hybrid – called “inner shell” complex
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/valence.php 38

39. Here we cannot explain 6 ligands around the Ni(2+).
In this case we have to use the “outer” 4d orbitals to form a hybrid:
Use a sp3d2 “outer shell” complex
Explain the bonding in a [Fe(CN)6] 4- complex
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40. MO THEORY
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41. No orbital mixing
here – because the
energy difference
between N and O
is high
-> small s- and pz-
interaction
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42. sp-mixing – example B2 molecule which is a diradical:
https://en.wikipedia.org/wiki/Molecular_orbital_diagram 42

43. High energy difference
-> small mixing
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44. Multi-atomic molecules: form GROUP ORBITALS
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45. Formation of LGO
* ligand group orbitals *
H2O molecule has c2v symmetry
The 2 H-s-orbitals can be combined
to form 2 LGO’s:
One symmetric, another anti-
symmetric
A1 symmetry
B2 symmetric
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46. Get LGO’s from group theory
http://plato.mercyhurst.edu/chemistry/kjircitano/inorgstudysheets/inorgstudyexamii.htm
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47. MO’s from oxygen AO’s and LGO’s
Bonding interactions: 2 Lone Pairs:
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48. Do the same exercise with NH3 (c3v symmetry)
Find the 3 group orbitals of the 3 H s-orbitals
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49. Construct LGO’s
Graphical approach
(1)Arrange all ligand orbitals around the central atom
(2)First MO-combination: all are in the same phase
(3)Draw one node plane symmetrically = next energy level
(4)Draw two node planes symmetrically
LGO no. 1
LGO no. 2
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50. Example: NH3
3 H-orbitals
-> 3 LGO’s :
(1) all same phase
(2) One node
(3) One node
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51. Coordination compounds
Find 6 symmetry adapted ligand
combinations (SALC)
To fit with the metal s- p- and d-
orbitals
51

52. http://faculty.uml.edu/ndeluca/84.334/topics/topic6.htm 52

53. ML6 complex – Co(NH3)6 (2+)
Co(2+)
NH3 ligand binds by the lone pair
of ammonia:
Insert the
bonding ?
Insert the
electrons -
which are
bonding, non-
bonding, anti-
bonding ?
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54. Influence of ligand pi and pi* orbitals
http://wwwchem.uwimona.edu.jm/courses/LFT.html 54