SYNOPSES:
Chemical bond and molecule:
It is an attractive force, which holds various constituents like atoms, ions,
etc. together in different chemical molecules or ions.
Molecules can be of homo-nuclear atom or hetero-nuclear atom. For
instance,
Homo-nuclear diatomic molecules: H2, O2, Cl2 etc.
Hetero-nuclear polyatomic molecules: H2O, CO, H2S, HCl etc.
Kossel and Lewis Approach to Chemical Bonding:
In order to explain the formation of chemical bond in terms of electrons, the
scientists like Kossel and Lewis were among the first ones to provide some
logical explanation of valence, which was based on the inertness of noble
gases.
Ionic Bond:
The elements, having low ionization energy and those elements, having
high electron affinity form ionic bond easily.
In crystal state of ionic compounds, there is a systematic three dimensional
arrangement of cations and anions resulting to form a crystal structure.
Covalent Bond:
When two or more than two atoms share their valence electrons to
make pair of electrons, they are said to be joint by covalent bond.
According to number of shared electron pairs between two atoms, there are
single, double or triple bonds established.
On the basis of the type of sharing and type of atom of combining
elements, the covalent bond is divided into 3 parts:
(i) Polar covalent bond
(ii) Non–polar covalent bond
(iii) Co–ordinate covalent bond
Lewis Structure of Covalent Bond:
Lewis explained the covalent bond formation in compound on the basis of
electron configuration. He used dots to represent valence electrons.
Born-Haber cycle:
With the help of this cycle, using ionization energy, energy gained by
getting electron and other information, one can calculate crystal lattice
energy.
Resonance structures:
These are the characteristic structures of a molecule, which describe
delocalized electrons within certain molecules, where the bonding can not
expressed by one single bond.
Geometry of molecules:
Molecules have different shapes. The geometrical structures of different
molecules are:
(1) Linear
(2) Planar tri-gonal,
(3) Tetrahedral
(4) Square planar
(5) Tri-gonal bi-pyramidal
(6) Square pyramidal
(7) Octahedral
VSEPR principle:
In order to understand the shapes of molecules, the significant theory of:
“Valence Shell Electron Pair Repulsion - VSEPR” is proposed.
This theory also suggests that how a covalent bond is produced by
overlapping of atomic orbitals.
On the basis of VSEPR principle, the bond angles in some molecules are:
(i) Bonding in methane is 109°28’
(ii) Bonding in ammonia is 107°.
(iii) Bonding in water is 104°30’.
Polarity of bonds:
In hetero-nuclear diatomic molecule, the electron-pair of covalent bond
between the two atoms is shifted more towards more electronegative atom,
because of its larger effective nuclear charge and by this way, the polarity
is produced in the molecule.
Approach to covalent bond:
On the basis of quantum mechanics, the formation of a covalent bond is
described by two modern theories:
(i) Valence Bond Theory.
(ii) Molecular Orbital Theory.
Valence bond theory explains the formation and directional properties of
bonds in polyatomic molecules in terms of overlapping and hybridization of
atomic orbitals.
Molecular orbital theory suggests: In a molecule, the electrons are
present in the various molecular orbitals as the electrons present in the
various atomic orbitals of an atom.
Overlapping of atomic orbitals:
When two atoms come close to each other, there is overlapping of atomic
orbitals. This overlap may be positive, negative or zero depending upon the
properties of overlapping of atomic orbitals.
Depending upon the types of overlapping, there are mainly 2 types of bond
formed:
(i) Sigma (σ) bond
(ii) pi (π) bond.
Sigma (σ
σ) bond:
It involves the end to end (hand-on) overlap of bonding orbitals along the
internuclear axis. E. g., Axial overlapping of s–s, s–p, and p–p orbitals.
pi (π
π) bond:
It involves the sidewise overlapping of the atomic orbitals in such a way
that their axes remain parallel to each other and perpendicular to the
internuclear axis. E.g., Parallel overlapping of p–p orbitals.
The strength of a bond:
It depends upon the extent of overlapping, so that sigma (σ) bond is
stronger as compared to pi (π) bond.
Molecular orbital theory:The basic idea of the molecular orbital theory
is that atomic orbitals of individual atoms combine. They form two types of
molecular orbitals:
(i) Bonding Molecular Orbital (BMO)
(ii) Anti–Bonding Molecular Orbitals (ABMO).
Linear Combination of Atomic Orbitals (LCAO):
LCAO method is used to describe the combination of two atomic orbitals
and formation of two molecular orbitals by the linear combination (addition
or subtraction) of acceptable wave functions. These two MO are BMO and
ABMO.
Symbols σ and π are used to indicate BMO, while σ* and π* symbols are
used to indicate ABMO.
The relative energies of molecular orbitals:
The increasing order of molecular orbitals for molecules from H2 to N2:
σ1s < σ*1s < σ2s < σ*2s < (π2px = π2py) < σ2pz < (π*2px = π*2py) <
σ*2pz
The order is changed for molecules from O2 to Ne2 as under:
σ1s < σ*1s < σ2s < σ*2s < σ2pz < (π2px = π2py) < (π*2px = π*2py) <
σ*2pz
By this electronic configuration of a molecule, some important information
obtained, such as: stability of molecule bond order, magnetic character
etc.
Intermolecular attraction force:
This is the weak attraction force, which cannot be explained by any other
chemical attraction force, but ‘Van Der Waals attraction force’. This is
because the electrons on the surface of molecule experience some
attraction towards the nucleus of other molecule.
Hydrogen bond:
It is an attractive force, which binds hydrogen atom of 1 molecule with the
more electronegative atom (F, O or N) of another molecule.
There are two types of H-bond:
(i) Intra-molecular H–bond: It exists within the same molecule.
(ii)Intermolecular H–bond: It exists between the two different molecules of
same or different compounds.
The Metallic Bond:
In metallic crystal, the positively charged kernels remains in space as close
as possible and keep equal distance amongst them due to the attractive
force of delocalized electrons. It also can be explained by the theory of
electron sea model.
Hybridization:
It involves the combination of two or more than two orbitals, having very
less energy difference and then formation of same number of orbitals,
having similar shapes and energies.
PCl5 shows sp3d hybridization, while SF6 shows sp3d2 hybridization.