Chemical Bonding
Chemical Bonding
Chemical Bonding
Structures
Dikki
Institute
Kossel Lewis
Approach
Lewis approach
He pictured the atom in terms of a positively
Lewis symbol
of valence electrons.
This number of valence electrons help to
calculate the common or group valence of
that element.
The group valence of the element is
generally either equal to the number of
dots in Lewis symbol or 8 minus the
number of dots.
Other important facts given by Kossel are In periodic table, the highly electronegative halogens
Example
Na
Na+ + e-
[Ne] 3s1
[Ne]
Cl-
or
[Ar]
NaCl
Octet rule
According to electronic theory of chemical
Covalent Bonding
Covalent bond is the sharing of the
VALENCE
Br
Br
Covalent bond
The bond formed by atoms when they share
Formal charge
Aformal charge(FC) is the charge
exceptions to it.
1)the incomplete octet of the
central atom
In some compounds, the
number of electrons
surrounding the central
atom is less than eight.
Examples- LiCl, BeH2 and
BCl3
2) Odd-electron molecule
in molecules with an odd number of
electrons like nitric oxide, NO and nitrogen
dioxide, the octet rule is not satisfied for all
atoms.
of the molecules.
6) It does not explain the relative stability of
the molecules being totally silent about the
energy of a molecule.
Bond Parameters
Bond parameters
Bond length
Bond angle
Bond enthalpy
It is defined as the amount of energy required to break one
mole of bonds of a particular type between two atoms in a
gaseous state.
The unit of bond enthalpy is kJ per mol.
ExampleThe H-H bond enthalpy in hydrogen molecule is 435.8
kJ/mol.
H2 -------> H(g)+H(g)
aH=498 kJ/mol
Bond order
In Lewis description of covalent bond, the
Resonance
Resonance structures
A single Lewis structure may be inadequate
can be represented by
two structures as shown
in the figure.
For O3 these two
structures constitute the
canonical structures or
resonance structures and
their hybrid shown by the
third figure represents the
structure of o3 more
accurately. this is called
resonance hybrid.
Resonance Energy
It is the difference between the actual bond
energy of the molecule and that of the most
stable of the resonating structures (having
least energy).
Polarity Of Bonds
Polarity of Bonds
The existence of a hundred percent ionic or covalent bond
represents an ideal situation. In reality no bond or a
compound is either completely covalent or ionic. Even in
case of covalent bond between two hydrogen atoms, there
is some ionic character.
When covalent bond is formed between two similar atoms, for
example in H2, O2, Cl2, N2or F2, the shared pair of electrons
is equally attracted by the two atoms. As a result electron
pair is situated exactly between the two identical nuclei.
The bond so formed is called non polar covalent bond.
Contrary to this in case of a heteronuclear molecule like HF,
the shared electron pair between the two atoms gets
displaced more towards fluorine since the electronegativity
of fluorine (Unit 3) is far greater than that of hydrogen. The
resultant covalent bond is a polar covalent bond.
As a result of polarisation, the molecule possesses the dipole
moment (depicted below) which can be defined as the
product of the magnitude of the charge and the distance
between the centres of positive and negative charge. It is
usually designated by a Greek letter ;. Mathematically, it
is expressed as follows :
VSEPR Theory
follows:
The shape of a molecule depends upon the
Bond pair (bp) > Bond pair (bp) Bond pair (bp)
Nyholm and Gillespie (1957) refined the VSEPR
ORBITAL OVERLAP
CONCEPT
ORBITAL OVERLAP
CONCEPT
If orbitals of 2 atoms are mixed with each other partially during bond
the same, i.e., one is positive and the other negative, then it is a
type of negative overlap.
3.Zero overlap: If overlap of orbitals present in 2 different planes
TYPES OF OVERLAPS
RESULTING IN SIGMA BOND
FORMATION
Due to lateral/sideways
overlap of P-P orbitals
present in the same
plane, Pi bond is
formed.
COMPARISION
It is a strong bond.
Electron cloud is
SIGMA
PI
HYBRIDIZATION
The intermixing of different atomic orbitals of approximately
equal energy levels to produce hybrid orbitals before bond
formation is called as HYBRIDIZATION.
Here arrangement of hybrid orbitals are such that there is
minimum repulsion in between the hybrid orbitals.
No. of orbitals mixed=No. of hybrid orbitals produced.
DIFFERENT TYPES OF HYBRIDISATIONS:
Sp Hybridization
NO. of hybrid orbitals produced=2
Structure = LINEAR
Bond angle = 180 degree ..e.g. BeF
sp HYBRIDIDSATION
SP2 Hybridization:
No. of hybrid orbitals produced = 3
Arrangement of these orbitals = TRIGONAL PLANAR
Bond angle=120 degree..ex. BF, etc.
SP3 Hybridization:
No. of hybrid orbitals = 4
Arrangement= TETRAHEDRAL
BOND ANGLE..
IN CH = 109.28 degree
IN NH = 107.3 degree.
Here, in methane, sigma bond is formed between H and C atom due to
SP2
Hybridization
SP3
Hybridization
SP3d Hybridization:
NO. of hybrid orbitals produced = 5
Structure = TRIGONAL BIPYRAMIDAL
BOND ANGLES:
Equatorial = 120 degree.
Axial = 90 degree.
If lone pair of electron is present at central atom, its position is always
equatorial.Ex.PCl5
Length of axial bonds is longer than that of equatorial bonds because
of minimum repulsion.
IT may have different shapes according to no. of lone pairs it has:
1 lone pair seesaw shapee.g. SF
2 lone pairs bent T shapeE.g. BrF
3 Lone pairs Linear shape
SP3d2 Hybridization:
No. of hybrid orbitals produced = 6
Arrangement = OCTAHEDRALe.g. SF
( lone pairs = 0)
( lone pairs = 1)
(lone pairs = 2)
SP3d3 Hybridization:
No. of hybrid orbitals = 7
Arrangement = PENTAGONAL BIPYRAMIDAL
BOND ANGLES:
Axial with equatorial = 90 degree.
Equatorial to equatorial = 72 degree.
SP3d2 Hybridization
Sp3d3 Hybridization
RULES REGARDING
HYBRIDIZATION
Only orbitals of approximately same energy levels can
take part.
No. of orbitals mixed = No. of hybrid orbitals produced.
Most hybrid orbitals are similar but not always identical
in shape. They may differ from one another in their
orientation in space.
The electron waves in hybrid orbitals repel each other
and this tend to the farthest apart.
Hybrid orbitals can form only sigma bonds.
Depending on the number and the nature of the orbitals
undergoing hybridization, various types of hybrid orbitals
directing towards the corners of specified geometrical
figures come into existence.
Molecular Orbital
Theory
MOT
OF ATOMIC ORBITALS
For atomic orbitals to combine, resulting in
Designations of Molecular
Orbitals
Just as atomic orbitals are designated as s, p, d,
fetcmolecular orbitals of diatomic molecules are
named(sigma) ,
(pi) ,(delta)etc.
MOLECULAR ORBITALS
The molecular orbitals which are cylindrically
symmetrical around inter-nuclear axis are
called- molecular orbitals. The molecular
orbital formed by the addition of 1s orbitals is
designated as1s and the molecular orbital
formed by subtraction of 1s orbitals is designated
as* 1s . Similarly combination of 2s orbital
results in the formation of two
2 s - molecular orbitals designated as2s
and* 2s
O2
DETERMINATION OF SHAPE
OF MOLECULES
1.Determine the number of electrons in the
molecule. We get the number of electrons per atom
from their atomic number on the periodic table.
(Remember to determine the total number of
electrons, not just the valence electrons.)
2.Fill the molecular orbitals from bottom to top
until all the electrons are added. Describe the
electrons with arrows. Put two arrows in each
molecular orbital, with the first arrow pointing up
and the second pointing down.
3. Orbitals of equal energy are half filled with
parallel spin before they begin to pair up.
Examples
1. The molecular orbital diagram for a diatomic hydrogen molecule, H2, is
stable.
The two unpaired electrons show that O 2is paramagnetic
BONDING IN HOMONUCLEAR
DIATOMIC MOLECULES
Diatomic molecules are molecules composed only
of two atoms, of either the same or different
chemical elements. The prefix di- is of Greek origin,
meaning two. Common diatomic molecules are
hydrogen (H2), nitrogen (N2), oxygen (O2), and carbon
monoxide
(CO).
Seven
elements
exist
as
homonuclear
diatomic
molecules
at
room
temperature: H2, N2, O2, F2, Cl2, Br2, and I2. Many
elements and chemical compounds aside from these
form diatomic molecules when evaporated. The
noble gases do not form diatomic molecules: this can
be explained using molecular orbital theory (see
molecular orbital diagram).
H MOLECULE
2
INTRODUCTION
1.Two H atoms in their ground state
configuration come together and form a
single bond. The bond formation stabilizes
both atoms and, therefore, is lower in energy
than the atomic orbitals. This is also
observed in Valence Bond Theory, which
implies that each H atom in H2shares its
electron with one another, so that both can
achieve the stable configuration of He.
2.On top of that, MO Theory allows one to
compute the amount of energy released from
a bond formation and a distance between
two bonded atoms as well as predict the
magnetic property of a molecule (or a
substance). For H2, the bond strength is -432
kJ/Mol, and the bond length is 74 angstrom
(or 74 pm). H2 is a diamagnetic molecule
because the electrons paired up; therefore, it
is not attracted by a magnetic field.
Bond order in H2
Bond order = 1/2 (#e- in bonding MO - #e- in
antibonding MO)
For H2, bond order = 1/2 (2-0) = 1, which means
H2 has only one bond. The antibonding orbital is
empty. Thus, H2 is a stable molecule.
Again, in the MO, there is no unpaired electron,
so H2 is diamagnetic
Hydrogen Bond
HYDROGEN BOND
In compounds of hydrogen with strongly
electronegative elements, such as fluorine,
oxygen and nitrogen, electron pair shared
between the two atoms lie far away from the
hydrogen atom. As a result, the hydrogen
atom becomes highly electropositive with
respect to the other atom. This phenomenon
of charge separation in the case of hydrogen
fluoride is represented as. Such a molecule
is said to bepolar. The molecule behaves as
a dipolebecause one end carries a positive
charge and the other end a negative charge.
The electrostatic force of attraction between
such molecules should be very strong. This is
becausethepositive end of one moleculeis
attracted by thenegative end of the other
molecule. Thus, two or molecules may
associate together to form larger cluster of
molecules. This is illustrated below for the
described as (HF)n.
It may be noted that hydrogen atom is
described as (H2O)n
Intramolecular hydrogen
bonding
This type of bonding involves electrostatic
forces of attraction between hydrogen and
electronegative elementboth present in the
same moleculeof the substance. Examples
o-nitrophenol and salicylaldehyde.
p-Nitrophenol , on account of large distance
Water
Density in solid state(ice) is less than that in
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