Chemical Bonding & Molecular Structure

Chemical Bonding & Molecular Structure

Content

Kossel - Lewis approach to chemical bonding

Octet rule

Covalent bond

Ionic bond

Bond parameters

VSEPR's theory

Valence bond theory


Matter is made up of one or different type of elements. Under normal conditions no other element exists as independent atom in nature, except noble gases. The attractive forces

which holds various constituents (atoms, ions etc.) together in different chemical species is called a chemical bond.


In the formation of a molecule, only the outer shell electron take part in chemical combination and are known as valence electrons. The inner electrons are well protected and are

generally not involved in the combination process.


The bond formed, as a result of electrostatic attraction between the positive and negative ions as the electrovalent bond.

A covalent bond is formed by the sharing of a pair of electrons between the two atoms. In this shared pair of electrons, each atom contributes one electron.

If one electron is shared between the two atoms ---- single covalent bond.

Two electron pairs are shared ------- double covalent bond

Three electron pairs shared ------ triple bond

The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and

the number of electrons assigned to that atom in the Lewis structure.




Most stable Lewis structure is one in which the atoms have a lowest formal charge.

Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. Bond lengths are measured by spectroscopic, X - ray diffraction and

electron diffraction techniques. The covalent radius is measured approximately as the radius of an atom core which is in contact with the core of an adjacent atom in a bounded

situation. The Vander Waals radius represents the overall size of the atom which includes its valence shell in a non - bonded situation.


Bond angle is defined as an angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complexion.

Bond angle is expressed in degree which gives some idea regarding the distribution of orbitals around the central atom in a molecule/complex ion and hence it helps us in determining the shape.


Bond enthalpy is defined as the energy required to break a particular bond in one mole of gaseous molecule.

Bond order is defined as the number of covalent bonds present between the two atoms of a molecule.



Isoelectronic molecules and ions have identical bond orders e.g. F₂ and O₂^2- have bond order 1. N₂, CO and NO⁺ have bond order 3.



Polarity of bonds - In case of a heteronuclear molecule like HF, the shared electron pair between the two atoms gets displaced more towards fluorine since the electro negativity of

fluorine is far greater than that of hydrogen. The result covalent bond is a polar covalent bond.

As a result of polarisation, the molecule possesses the dipole moment which can be defined as the product of the magnitude of the charge and the distance between the centres of

positive and negative charge.

Dipole moment µ = charge Q x distance of separation (r)

Dipole moment is usually expressed in Debye units (D) the conversion factor is -

1D = 3.33564 x 10^-30 cm


In case of polyatomic molecules the dipole moment not only depend upon the dipole moments of bonds known as bond dipoles but also on the spatial arrangement of various bonds

in the molecule. In such case the dipole moment of a molecule is the vector sum of the dipole moments of various bonds.

Percentage of Ionic Character

Every ionic compound having some percentage of covalent character according to Fajan's rule. The percentage of ionic character in a compound having some covalent character can

be calculated by the following equation. The percent ionic character

=



The valence shell electron pair repulsion (VSEPR) theory-

The valence shell electron repulsion theory was developed by Sidgwick and Powell and is used to predict the shape of covalent molecule. Basic assumptions are as follows:-

Electron pairs (EP) in the valence shell of a central atom repel each other and to minimise repulsion, these electron pair should be kept as far apart as possible.

a triple bond or double bond is considered as one electron pair.

When a molecule is represented by two or more structure, then the VSEPR theory can be applied to any structure.


Valence bond theory - valence bond theory was introduced by Heitler and London (1927) and developed further by Pauling and others. A discussion on the valence bond theory is

based on the knowledge of atomic orbitals, electronic configurations of elements, the overall criteria of atomic orbitals, the hybridization of atomic orbitals and the principle of - variation

and superposition.

According to valence bond theory-

A covalent bond is formed by the overlapping of atomic orbitals from the valence shell. The overlap means that the two orbitals have same common region in space and can

accommodate only two electron with opposite spin.

Overlapping of AO can takes place in two different ways.

Overlapping of AO along the internuclear axis or end-to-end or head on overlapping of AO.

Overlapping of AO perpendicular to the axis or sideways or parallel overlapping of AO. This give pi () bond.


Hybridization -

Hybridization is a mixing of atomic orbitals from the same atom having comparable energy to form hybrid orbitals which are degenerate.

Number of hybrid orbitals formed equals to the number of atomic orbitals mixed.


VSEPR (Valence Shell Electron Pair Repulsion Theory)

This theory starts from the general principle that valence shell electrons occupy essentially localised orbitals. Mutual interaction among the electrons orient the orbitals in space to an

equilibrium position where repulsion becomes minimum. The extent of repulsive interaction then follows the order.

Lone pair - lone pair > lone pair - bond pair> bond pair - bond pair


A lone pair is concentrated around the central atom while a bond pair is pulled out between two bonded atoms. As such repulsion becomes greater when a lone pair is involved.

Let's take an example to illustrate this theory. CH₄ contains no lone pairs. The bond pair - bond pair interactions brings about the most stable equilibrium bond angle of 109°28¢, the

angle predicted from sp3 hybridisation.


NH₃ contains one lone pair while H₂O has two. The lone pair - bond pair and

lone pair - lone pair repulsions tend to reduce the bond angles. The bond pair-bond pair interaction tends to open the angles. But as the interaction is weaker than the lone pair - bond

pair repulsion, so the ultimate bond angle is smaller. Bond angle NH₃ = 107° and bond angle H₂O = 104.5° H₂O has a smaller bond angle because there are 2 lone pairs which

reduces it further.



Rule for Determination of Total Number of Hybrid Orbitals

Detect the central atom along with the peripheral atoms.

Count the valence electron of the central atom and the peripheral atoms.

Divide the above value by 8. Then the quotient gives the number of s bonds and the remainder gives the non-bonded electrons. So number of lone pair equals to (non bonded electrons)/2

The number of s bonds and the lone pair gives the total number of hybrid orbitals.

Note: Whenever there are lone pairs in TBP geometry they should be placed in equatorial position so that repulsion are minimum.


Co-ordinate covalent bond is a covalent bond between the atoms A and B in which both the electrons comes from one atom and none from the second. The atom which accepts

e-pair is called Lewis acid and one which denotes e-pair is called Lewis base.



The molecule which accept e-pair, the hybridization of central atom is increased by one atomic orbital (sp sp², sp² sp³, sp³ sp³d). one which denotes e - pair

hybridization, remains same.

Note:- According to valence bond theory alone of electron does not take part in bonding.



Molecular orbital theory

The main postulates of molecular orbital theory are given below:

When atoms are join to form a molecule, new regions of high electron probability i.e. molecular orbitals are established.

The molecular orbitals are found by the combination of atomic orbitals of roughly the same energy and proper symmetry.

The number of molecular orbitals formed is equal to the number of combining atomic orbitals. The combination of the two atomic orbitals produces two MO bonding MO and
antibonding MO

The energy of the bonding MO is less than the corresponding antibonding MO hence it is more stable.

In simple homonuclear diatomic molecules the order of MO's based on increasing energy is



If the molecule contains unpaired electrons in MO's it will be paramagnetic but if all the electrons are paired up then the molecule will be diamagnetic.

Bond order = no. of occupying bonding MO's - no. of e-s occupying antibonding MO's