Molecular orbital theory provides an approach to calculate molecular orbitals through a variational method. This involves taking linear combinations of atomic orbitals to form molecular orbitals. Electrons occupy these molecular orbitals according to certain rules. The molecular orbital theory can explain properties such as why some molecules are paramagnetic that valence bond theory cannot. Calculating molecular orbitals variationally involves using trial wave functions in the Schrodinger equation to find the lowest possible energy state.

2. 1. Lewis Structure 2.Approach to
calculate Molecular Orbits
3.Donor Acceptor Properties

3.  Molecular Electronics.
1. What is molecular electronics.
2. History.
3. About molecular electronics.
4. Chemical structure of an acceptor-bridge-donor molecule.
 Lewis Structure.
1. Single bond and Lewis structure.
2. Rules for drawing Lewis structure.
3. Resonance structures of Methyl Nitrite.
4. Another examples of Resonance structures.

4. 
MOLECULAR
ELECTRONICS

5.  Molecular electronics lies at the intersection of chemistry
with nano-electronics.
 Molecular electronics is the study and application of
molecular building blocks for the fabrication of electronic
components. It is an interdisciplinary area that spans physics
, chemistry, and materials science.
 The unifying feature is use of molecular building blocks to
fabricate electronic components.
 Due to the prospect of size reduction in electronics offered
by molecular-level control of properties, molecular
electronics has generated much excitement.
 It provides a potential means to extend Moore’s Law beyond
the foreseen limits of small-scale conventional silicon
integrated circuits.

6.  The first time in history molecular electronics are mentioned
was in 1956 by the German physicist Arthur von Hippel, who
suggested a bottom up procedure of developing electronics
from atoms and molecules rather than using prefabricated
materials, an idea he named molecular engineering.
 However the first breakthrough in the field is considered by
many the article by rather and Aviram in 1974.
 In this article named Molecular Rectifiers, they presented a
theoretical calculations of transport through a modified
charge transfer molecule with donor acceptor groups that
would allow transport only in one direction, essentially like a
semiconductor diode.
 This was a breakthrough that inspired many years of research

7.  The principle of above research is that biological
systems can give useful paradigm for developing
electronic and computational devices at the
molecular level.
 The approach involves the design and synthesis of
dyads, triads and other super molecular species
using the techniques of organic chemistry.
 In order to studied newly prepared molecule time
resolved laser spectroscopy, NMR spectroscopy.

8. For electronic application, molecular structures
has four major advantages.
 Size.
 Assembly & Recognition.
 Dynamical stereochemistry.
 Synthetic tailor ability.

9.  (a) Chemical structure
of an acceptor-bridge-
donor molecule
proposed as a
molecular rectifier.
The electron acceptor
is
tetracyanoquinodimeth
ane (TCNQ) and the
donor is
tetrathiofulvalene
(TTF). The bridge
consists of three
parallel chains of

11.  Lewis structures. of 𝐻2and 𝐶𝑙2molecules:
 Slightly complicated ammonia molecule:
 Three of the valence electrons are shared with the three
hydrogen atoms to make
 three single bonds. The remaining electrons on the
nitrogen form a localized concentration of negative

12. DOUBLE BOND TRIPLE BOND
 Inorganic Example
 Organic Example
 Inorganic Example:
 Organic Example:

13.  Once we have determined the number of total valance
electrons we can start distributing them throughout the
molecule.
 When we represents electrons they will be in pairs (since an
orbital holds 2 electrons).
 Electron pairs can be represented with 2 dots or a solid line.
 In covalent compounds atoms share electrons to form bonds
in order to achieve stable noble gas electron configuration. In
ionic compounds electrons are transferred from one atom to
another to achieve stable noble gas electron configuration.

14.  Step 1: Count total number of valance electron
𝑃𝐶𝐿3 5 + 3 × 7 = 26
 Step 2:Choose the least electronegative atom
and put it at the center of the structure,
connecting the other atoms by single bonds.
 Step 3:Complete octets for the outer atoms—
remember that the bonds count for two
electrons each:
 Step 4:Complete the octet for the central atom

15.  Step 5:If you run out of electrons before you form an
octet on the central atom, keep forming multiple bonds
until you do
 Formal charge is the charge calculated for an atom in
a Lewis structure on the basis of an equal sharing of
bonded electron pairs
 Step 6: Now assign formal charges. The formal charge is
the valence number minus (the number of electrons in
lone pairs plus half the number of electrons in bonds).
For example, the formal charges on each atom in
6 − 7 = −1, 4 − 4 = 0, and 6 − 5 = +1.
 The formal charge in

16.  Step 7: Calculate formal charges.
 Example: This structure has formal charges; is
less stable Lewis structure
 Same atomic positions
Differ in electron position

17.  Electrons in molecules are often delocalized
between two or more atoms.
Electrons in a single Lewis structure are
assigned to specific atoms-a single Lewis structure
is insufficient to show electron delocalization.
Composite of resonance forms more accurately
depicts electron distribution.

18.  In Benzene each carbon satisfies the octet rule
and each hydrogen shares two electrons.
 In this case, the new states are delocalized over
the entire benzene ring. The bonding is described
as aromatic and not shown as alternating double
and single bonds.
 Two Lewis structure for benzene and the
resultant aromatic structure are drawn:

19.  What is MOT?
 Probably the second and most applicable for
bonding and structure of molecule.
 Put forward by HUND and MULLIKEN.
 Later modified by JONES and COULSON.
 Drawbacks of VBT-
 Fail to explain paramagnetic nature of
 Resonance plays a major role in VBT but no role in
MOT.
 VBT did not give any weightage to ionic structure
but MOT did not.
 It didn’t say about Anti bonding orbital thus failed
to say about spectral lines.

20. MO theory suggests that atomic orbitals of
different atom combine to create Molecular
Orbitals.
Electrons in these molecular orbitals belong to
the molecules as whole.
This contracts to VB theory which suggests that
electrons are shared by simple overlap atomic
orbital’s or hybridized atomic orbitals’.
Molecular orbital can be constructed from linear
combination of atomic orbital’s
MO= LCAO

21.  When two AOs mix, two Mos will be produced.
 Each orbital can have a total of two electrons
( pauli principal)
 Lowest energy orbitals are filled first (Aufbau principal)
 Unpaired electrons have parallel spin (Hund’s rule)
 Bond order = ½ ( bonding electrons-antibonding
electrons)

22.  Rules for linear combination:
 Atomic orbital’s must be roughly of the same energy.
 The orbital must overlap one another as much as
possible- atoms must be close enough for effective
overlap.
 In order to produce bonding and antibonding Mos ,
either the symmetry of two atomic orbital must remain
unchanged when rotated about the internuclear line or
both atomic orbital’s must change symmetry in
identical manner.

23.  The wave function for the molecular orbitals
can be approximated by taking linear
combination of atomic orbitals.

25. INTRODUCTION OF MOT:
In 1927 hietler and London proposed the valance
Bond theory .Valence Bond Theory fails to answer
certain questions like Why He2 molecule does not
exist and why O2 is paramagnetic? Therefore in 1932
F. Hood and Robert S. Mulliken came up with
theory known as Molecular Orbital Theory to
explain questions like above. According to
Molecular Orbital Theory individual atomic orbitals
combine to form molecular orbitals, as the electrons
of an atom are present in various atomic orbitals and
are associated with several nuclei.

26. Magnetic Behavior:
If all the molecular orbitals in species are spin
paired, the substance is diamagnetic. But if one or
more molecular orbitals are singly occupied it is
paramagnetic. For Example, if we look at CO
Molecule, it is diamagnetic as all the electron in CO
are paired

27.  When we discuss about MOT of hydrogen molecule then
we know that in valance shell of hydrogen atom have one
one electron,
 When they combine together they form sigma bond.
 There is no any electrons are in Antibonding molecular
orbital.
 Both two electrons are occupy bonding molecular orbital.
Hydrogen molecule is formed sigma bond by
combination of S-S orbital
MOT OF HYDROGEN

28. The Variational Approach to Calculating
Molecular Orbitals
 In this section, we will review the variational
approach to building up molecular orbitals from
atomic orbitals. At the end of the process, we will draw
diagrams of the energy levels associated with the
various types of orbitals. Knowing how these energy
levels are filled with electrons will often be all that we
need in order to understand the electronic properties
of the molecules. The variational formulae are derived
as follows: Multiplying the time independent
Schrödinger equation (Equation 2.31) from the left by
ψ∗ and integrating both sides yields the following
result for the energy of a stationary state:

29. where we have deliberately not used the fact that
eigenstates of the Schrödinger equation satisfy ψn|ψm
= δnm. If we did not know the eigenstates for a system,
we might, nonetheless, guess a trial function for a
particular state, φ. 8.3 The variational approach to
calculating molecular orbitals 269 If we use Equation
8.1 to calculate an “energy,” Eφ, we would always get a
value greater than (or, at best, equal to) E0 because E0
is the ground state energy, the lowest allowed by
quantum mechanics, given the intrinsic quantum
mechanical kinetic energy of the electron.