bonding theories in chemistry

Bonding TheoriesAdvanced Inorg. Chem.

Dr. Chris Sontag

University of PhayaoOct.2016

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After this lesson, we should understand:

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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 ?

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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 ?

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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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Electronegativity (EN)

The amount of EN difference determines the polarity of the bond:

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Allred-Rochow EN

Example: Flourine (r = 72 pm)Carbon (r = 77 pm)Calculate the AR electronegativity

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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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VALENCE ELECTRONS ANDLEWIS STRUCTURES

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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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Electron configurationdo NOT use core electrons !

Practise: write the configuration for:1. Fe2+

2. Pb3. W4. 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)

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Lewis MoleculesWrite the atom with the LOWEST EN in the middle !

Example(2) In molecules write all VE for each atom

Try yourself:

1. AlH3

2. LiH3. SiF4

4. C2H6

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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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Multiple Bonds and formal charges

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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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Transition metal compoundsExample:

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 +3Find the numbers for: KMnO4, MnO2, K4Fe(CN)6, Fe(CO)5

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

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VB Theory and molecular geometry

http://www.slideshare.net/Hoegler6/10-lecture 25

VB Theory = Hybridization of atomic orbitals

But NOT ALWAYS:

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Sigma - bonds 30

How many sigma-bonds can each atom form ?

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Try the same for FORMALDEHYDE HCHO32

Try the same for FORMALDEHYDE HCHO34

sp3d hybridization

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

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sp3d2 hybridization

Example: SF6 compared to SCl2 – both molecules are stable !

six

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Coordination Compounds

For TM ions the VE are counted ALL as d-electrons !

EMPTY metal orbitals are neededto 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

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MO THEORY

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

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sp-mixing – example B2 molecule which is a diradical:

https://en.wikipedia.org/wiki/Molecular_orbital_diagram 42

High energy difference-> small mixing

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Multi-atomic molecules: form GROUP ORBITALS

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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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Get LGO’s from group theory

http://plato.mercyhurst.edu/chemistry/kjircitano/inorgstudysheets/inorgstudyexamii.htm

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MO’s from oxygen AO’s and LGO’s

Bonding interactions: 2 Lone Pairs:

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

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Construct LGO’sGraphical 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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Example: NH3

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

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Coordination compounds

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

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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 ?

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Influence of ligand pi and pi* orbitals

http://wwwchem.uwimona.edu.jm/courses/LFT.html 54

bonding theories in chemistry

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