1. Essential Reading:
1. P. W. Atkins, Elements of Physical Chemistry, 4th Ed., Oxford University Press, 2007.
2. F. A. Carey, R. M. Guuliano, Organic Chemistry, Mcgraw-Hill, 6th edition, 2006.
3. J.D. Lee, Concise Inorganic Chemistry, 5th edition, Blackwell Publishing, 2008.
4. Fundamentals of molecular spectroscopy, C. N. Banwell, Tata McGraw-Hill Education, 1994.
Supplementary Reading:
1. J. Singh, L.D.S. Yadav, Advanced Organic Chemistry, PragatiPrakashan, 2009.
2. J. E. Huheey, E. A. Keiter and R. L. Keiter, Inorganic Chemistry, Principles of structure and
reactivity, Harper Collins, 1993.
3. Clayden, Greeves, Warren and Wothers, Organic Chemistry, Oxford, 2001.
4. B. R. Puri, L. R. Sharma, M. S. Pathania, Principles of physical Chemistry, ShobanLalNagin
Chand & Co., 2001.
Recommended Text Books
2. Ionic bond: Type of chemical bond that involves the electrostatic attraction
between oppositely charged ions. These ions represent atoms that have lost one or
more electrons (cations) and atoms that have gained one or more electrons (anions).
Here Sodium molecule is donating its 1 valence electron to the Chlorine
molecule. This creates a Sodium cation and a Chlorine anion. Notice that the net
charge of the compound is 0.
Basic Concept: Chemical Bonding
3. Some examples of ionic bonds and ionic compounds:
NaBr - sodium bromide NaF - sodium fluoride
KI - potassium iodide KCl - potassium chloride
CaCl2 - calcium chloride KBr - potassium bromide
Ionic bonding in sodium chloride
Formation of ionic bond in lithium fluoride
4. Covalent bond: A chemical bond that involves the sharing of electron
pairs between atoms. The stable balance of attractive and repulsive forces between
atoms when they share electrons is known as covalent bonding.
Here Phosphorous molecule is sharing its 3 unpaired electrons with 3 Chlorine
atoms. In the end product, all four of these molecules have 8 valence electrons and
satisfy the octet rule.
6. In chemistry, sigma bonds (σ bonds) are the strongest type of covalent chemical bond.
They are formed by head-on overlapping between atomic orbitals. Sigma bonding is
most clearly defined for diatomic molecules.
Sigma bond
7. Pi bonds (π bonds) are covalent chemical bonds where two lobes of one involved atomic
orbital overlap two lobes of the other involved atomic orbital. Each of these atomic
orbitals is zero at a shared nodal plane, passing through the two bonded nuclei.
π bond
8. Coordinate bond : A dipolar bond, more commonly known as a dative covalent
bond or coordinate bond is a kind of 2-center, 2-electron covalent bond in which the
two electrons derive from the same atom.
9. Metallic Bond: Metallic bonding constitutes the electrostatic attractive
forces between the delocalized electrons, called conduction electrons, gathered
in an electron cloud and the positively charged metal ions.
10. Valence Shell Electron Pair Repulsion (VSEPR) Theory
Valence shell electron pair repulsion (VSEPR) theory is a model in chemistry,
which is used for predicting the shapes of individual molecules.
The theory was suggested by Sidgwick and Powell in 1940 and was developed
by Gillespie and Nyholm in 1957. It is also called the Gillespie-Nyholm Theory
after the two main developers.
VSEPR theory is based on the idea that the geometry of a molecule or polyatomic
ion is determined primarily by repulsion among the pairs of electrons associated
with a central atom.
The pairs of electrons may be bonding or nonbonding (also called lone pairs).
Only valence electrons of the central atom influence the molecular shape in a
meaningful way.
11. VSEPR theory may be summarized as:
The shape of the molecule is determined by repulsions between all of the
electron pairs present in the valence shell.
A lone pair of electrons takes up more space around the central atom than a
bond pair. Three types of repulsion take place between the electrons of a
molecule:
The lone pair-lone pair repulsion (lp-lp)
The lone pair-bonding pair repulsion (lp-bp)
The bonding pair-bonding pair repulsion. (bp-bp)
12. The best spatial arrangement of the bonding pairs of electrons in
the valence orbitals is one in which the repulsions are minimized.
lp-lp> lp-bp> bp-bp
The magnitude of the repulsions between bonding pairs of electrons
depends on the electronegativity difference between central atom
and other atoms.
Double bonds cause more repulsion than single bonds, and triple
bonds cause more repulsion than a double bond.
13.
14. Predicted molecular shapes from Sidgwick- Powell Theory:
No. of electron
pairs in outer
shell
Arrangement of electron pairs Electron-pair
geometry
Bond angles
2
3
4
5
6
Linear
Trigonal Planar
Tetrahedral
Trigonal
bipyramid
Octahedral
180 0
120 0
109.50
90 0
120 0
90 0
15. Some examples using VSEPR Theory
SnCl2
Lewis model:
Shape : bent
lp-bp repulsions cause the Cl-Sn-Cl bond angle close to less than
120 0 (approx 950)
16. NH3
Lewis model:
Shape : Trigonal Pyramid
lp-bp repulsions cause the H-N-H angles to close to less than 109.5 o (107.3o).
17. H2O
Lewis model:
Shape : Bent
lp-bp repulsions cause the H-O-H angle to be lesser than 109.5 0 (104.50 )
18. ClF3
Lewis model:
Shape : T shape
Lone pairs occupy equatorial positions of trigonal bipyramid
lp-bp repulsions cause F-C-F angle to be lesser than 90 0
19. Limitations of VSEPR Theory
It fails to predict the shapes of isoelectronic species [ CH4 and NH4
+] and
transition metal compounds.
The model does not take relative size of substituents.
Atomic orbitals overlap cannot be explained by VSEPR theory.
The theory makes no predictions about the lengths of the bonds, which is
another aspect of the shape of a molecule.
20. Bent’s Rule
In a molecule, smaller bond angles are formed between electronegative ligands
since the central atom, to which the ligands are attached, tends to direct
bonding hybrid orbitals of greater p character towards its more electronegative
substituents.
Structure of water illustrating how the bond angle deviates from
the tetrahedral angle of 109.5°.
21. The carbon atoms are directing sp3, sp2, and sp orbitals towards the
hydrogen substituents. This simple system demonstrates that hybridised
atomic orbitals with higher p character will have a smaller angle
between them.
22. Limitations of Valence Bond Theory:
(i) It involves a number of assumptions.
(ii) It does not give quantitative interpretation of magnetic data.
(iii) It does not explain the color exhibited by coordination
compounds.
(iv) It does not give a quantitative interpretation of the
thermodynamic or kinetic stabilities of coordination compounds.
(v) It does not make exact predictions regarding the tetrahedral and
square planar structures of 4-coordinate complexes.
(vi) It does not distinguish between weak and strong ligands.
23. Molecular Orbital Theory
• Molecular orbitals result from the combination of atomic orbitals. Since orbitals are
wave functions, they can combine either constructively (forming a bonding
molecular orbital), or destructively (forming an antibonding molecular orbital).
• Consider the H2 molecule, for example. One of the molecular orbitals in this
molecule is constructed by adding the mathematical functions for the two 1s
atomic orbitals that come together to form this molecule. Another orbital is formed
by subtracting one of these functions from the other
24. Bonding Molecular Orbital Theory
The bonding orbital
results in increased electron
density between the two
nuclei, and is of lower energy
than the two separate atomic
orbitals.
25. Antibonding Molecular Orbital Theory
The antibonding
orbital results in a
node between the
two nuclei, and is of
greater energy than
the two separate
atomic orbitals.
26. Overlap of s & p Orbitals
+ -
+
+
+
+- -
- -
+
+
+
-
-
-
27. Sigma bonding orbitals
• From s orbitals on separate atoms
+ +
s orbital s orbital
+ ++ +
Sigma bonding
molecular orbital
28. Molecular Orbitals of the Second Energy Level
If we arbitrarily define the Z axis of the coordinate system for the O2 molecule
as the axis along which the bond forms, the 2pz orbitals on the adjacent atoms
will meet head-on to form a 2p bonding and a 2p* antibonding molecular orbital
29. Sigma bonding orbitals
• From p orbitals on separate atoms
p orbital p orbital
Sigma bonding
molecular orbital
30. The 2px orbitals on one atom interact with the 2px orbitals
on the other to form molecular orbitals that have a
different shape. These molecular orbitals are called pi (π )
orbitals because they look like p orbitals when viewed
along the bond.
31. Pi bonding orbitals
• P orbitals on separate atoms
Pi bonding
molecular orbital
