Bonding Theories in Chemistry

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

After this lesson, we should understand:
2

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

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

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

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

Allred-Rochow EN
Example:
Flourine (r = 72 pm)
Carbon (r = 77 pm)
Calculate the AR
electronegativity
10

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
11

VALENCE ELECTRONS AND
LEWIS STRUCTURES
12

The periodic table
is built up so that
elements with the
same number of VE
are in one column
= no. of VE
13

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

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)
15

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
17

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
19

Multiple Bonds and formal charges
20

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
21

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
22

VB THEORY
(DIAGRAMS FROM:
HTTP://WWW.SLIDESHARE.NET/HOEGLER6/10-LECTURE)
24

VB Theory and
molecular geometry
http://www.slideshare.net/Hoegler6/10-lecture 25

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

How many sigma-bonds can each atom form ?
31

Try the same for FORMALDEHYDE HCHO 32

Try the same for FORMALDEHYDE HCHO
34

sp3d hybridization
Example: PCl5 compared to PCl3 – both molecules are stable !
36

sp3d2 hybridization
Example: SF6 compared to SCl2 – both molecules are stable !
six
37

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

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
39

No orbital mixing
here – because the
energy difference
between N and O
is high
-> small s- and pz-
interaction
41

sp-mixing – example B2 molecule which is a diradical:
https://en.wikipedia.org/wiki/Molecular_orbital_diagram 42

High energy difference
-> small mixing
43

Multi-atomic molecules: form GROUP ORBITALS
44

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
45

Get LGO’s from group theory
http://plato.mercyhurst.edu/chemistry/kjircitano/inorgstudysheets/inorgstudyexamii.htm
46

MO’s from oxygen AO’s and LGO’s
Bonding interactions: 2 Lone Pairs:
47

Do the same exercise with NH3 (c3v symmetry)
Find the 3 group orbitals of the 3 H s-orbitals
48

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
49

Example: NH3
3 H-orbitals
-> 3 LGO’s :
(1) all same phase
(2) One node
(3) One node
50

Coordination compounds
Find 6 symmetry adapted ligand
combinations (SALC)
To fit with the metal s- p- and d-
orbitals
51

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

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 ?
53

Influence of ligand pi and pi* orbitals
http://wwwchem.uwimona.edu.jm/courses/LFT.html 54

Bonding Theories in Chemistry

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