General characteristics of coordinate covalent bond, valence bond approach, hybridization and VSEPR theory

General characteristics of coordinate-covalent bond

Covalent bond is formed when two atoms share a pair of electrons which results in the attraction of electrons by the nuclei of the atoms. Therefore, the atoms are held together. However, this may not always be the case, sometimes the shared pair of electrons come from the single atom which is called coordinate covalent bond.

Valence Bond Approach

This theory was proposed by Linus Pauling. According to this theory, atoms with unpaired electrons tend to combine with other atom with unpaired electrons. This results in pairing up of unpaired electrons and the stable electronic arrangement is attained. A covalent bond results from the pairing of electrons. According to Pauli' s exclusion principle that no two electrons in one atom can have all four quantum numbers same so the spins of two electrons must be opposite.

Directional Characteristics of covalent bond

When the overlapping of orbitals occurs through the internuclear axis, the extent of overlapping is maximum and the stronger bond is formed. Thus formed bond is called covalent bond. In oxygen atom there are six electrons in valence shell. The electronic configuration is 1s²2s²2px²2py¹2pz¹. Conventionally z-axis overlap along the internuclear axis to form \(\sigma\) bond in oxygen molecule. The 2py¹orbitals cannot fully overlap along internuclear axis rather these overlap sideways or laterally and form weaker bond in comparison to \(\sigma\) bond and is called \(\pi\) bond.

Hybridisation

The process of mixing of pure atomic orbitals on an atom of nearly equal energy to produce a new set of orbitals which are equal in number to the mixing orbitals and have identical energyand equivalent shape is known as hybridisation. The new orbitals so produced are called hybrid orbitals.

Silent features of hybridisation

1. Energy difference between the orbitals should be minimum.
2. Half as well as completely filled orbitals can take part in hybridisation.
3. The number of resulting hybrid orbitals are equal to the number of orbitals taking part in hybridisation.
4. The hybrid orbitals have equivalent energy and identical shape.
5. The hybrid orbitals can form stronger bond as compared to the pure atomic orbitals because they can overlap to greater extent.
6. The hybrid orbital has electron density concentrated on one side of the nucleus i.e. it has one lobe relateively larger than other.

Types of hybridisation

1. SP, diagonal (linear) hybridisation: one s atomic orbital intermix with one p-orbital.

2. SP², trigonal hybridisation: One s-orbital and two p-orbital combine to gibe a set of equivalent hybird orbitals with lobes directed to the corners of equiteral triangle.

3. SP³, tetrahedral hybridisation: One atomic orbital intermix with three p-orbitals of nearly equal energy to form four hybrid orbitals is knowm as sp³-hybridisation. The resulting hybrid orbitals are directed towards the four corner of regular tetrahedron with a bond angle 109°28′. For example formation of methane molecule (CH₄),

Outer electronic configuration carbon in ground state

= 2s², 2P_x^1,2P_y¹, 2P_z⁰

Outer electronic configuration of carbon in Excited state

= 2s¹,2P_x^1,2P_y¹, 2P_z¹

There are four orbitals having unpaired electrons. These 2s, 2p_x, 2p_yand 2p_zorbitals will first hybridise. Here these four hybridised orbitals overlap with the 1s-orbital of four hydrogen atoms through inter nuclear axes. As a resultσ-bonds are formed and the bond angle i s109°28′.

4. d²sp³or sp³d², octahedral hybridisation: When the dx²- y²and dz²orbitals are combined with an s orbital and a set of equivalent orbitals with lobes directed to the vertices of an octahedron can be formed.

5. dsp², square planar hybridisation: A dx²- y²orbital, an s orbital, and p_xand p_yorbitals combined to give a set of equivalent hybird orbitals with lobes directed to the corners of a square in the xy plane.

6. sp³d or dsp³trigonal bipyramidal hybridization: The orbitals s, p_x, p_y, p_z anddx²- y²combined to give a nonequivalent set of five hybrid orbitals directed to the vertices of a trigonal bipyramid.

7. sp³d³, pentagonal bipyramidal hybridisation: The orbitals s, p_x, p_y, p_zand three d-orbitals may be combined to give a set of seven hybrid. Five of them points towards the vertices of a regular pentagon while others are perpendicular to the plane.

VSEPR (Valence Shell Electron Pair Repulsion) theory:

This theory is a modified form of Sidgwick-Powell theory to predict and explain molecular shapes and bond angles more exactly by Gillespie and Nyholm. The postulates are as follows:

1. The shape of the molecule is determined by repulsions between all of the electron pairs present in the valence shell.
2. A lone pair of electron occupy more space in comparison to bond pair as lone pair is attracted to one nucleus while bond pair is shared by two nuclei. Hence the repulsion between two lone pairs is greater than repulsion between a lone pair and a bond pair, which in turn is greater than the repulsion between two bond pairs. Thus central metal atom with lone pairs are distorted from the ideal shape. L.P - L.P.> L.P.-b.P.> b.P.-b.P. (order of repulsion between electron pair)
3. The magnitude of repulsion between bonding bonding pairs of electrons depends on the elctronegativity difference between central atom and the other atoms. For example H-N-H bond angle in NH₃is 107°48′.
4. The molecules are arranged in shape that tend to minimize the electrostatic repulsion between electron pairs.
5. Repulsion between electron pairs at vertices greater than 115° apart can be neglected.

The prediction of the shape of the molecule on the basis of VSEPR theory

1. BF₃(Boron trifluoride) - In BF₃, B-atom contains three valence electrons and three fluoride atoms contributes three electrons each making a three pairs of electrons. Therefore, the shape of the molecle BF₃is trigonal planar according to the VSEPR theory.

   

2. SF₆(sulpur hexafluoride) - In SF₆, there are six valence electrons which form bond with the six F- atoms. Thus the central S-atom in SF₆has six pairs of electrons and the structure of SF₆is regular tetrahedral.

3. H₂O (Dihydrogen oxide) - The central O-atom in H₂O has eight outer electron. Therefore it has tetrahedron structure. The presence of two lone pairs in O-atom distort the bond angle 109°28′ to 104°27′ and has bent structure.

4. NH₃(Ammonia) - In the structure of NH₃, N has four electron pairs in outer shell, made up of three bond pairs and one lone pair. The shape is tetrahedral and the bond angle is distorted from 109°28′ to 107°48′ due to repulsion between bond pair and bond pair.

References

F.A. Cotton, G. Wilkilson and C. Gaus. Basic Inorganic Chemistry. 6th and 7th. John Wiley and Sons (Asia), Pvt., Ltd., 2007.

Lee, J.D. Concise Inorganic Chemistry. John Wiley and sons.Inc, 2007.

Chemwiki. n.d. <http://chemwiki.ucdavis.edu/Core/Theoretical_Chemistry/Chemical_Bonding/Valence_Bond_Theory/Hybridization>.