Grade 12 Chemistry Note

Chemical Bonding and Shape of Molecules

Hybridization:

The structures of different molecules can be explained on the basis of hybridization. For e.g., in case of carbon, the ground state electronic configuration is 1s2 2s2 2p1x 1y

 

To explain the tetravalency of carbon, it was proposed that one of the electrons from 2s filled orbital is promoted to the 2p empty orbital (2pz), which is in a higher energy state. Thus, four half-filled orbitals form in the valence shell this accounts for the bonding capacity of four carbon atoms. This state is known as excited state and the configuration of carbon in the excited state is:


The above configuration reveals that all the four bonds formed by carbon will not be identical. For e.g., in the formation of CH4 molecule, one C-H bond will be formed by the overlapping of 2s-orbital of C and 1s-orbital of H whereas the other three C-H bonds will be formed by the overlapping of 2p-orbitals of C and 1s-orbital of H. Therefore, all the bonds will not be equivalent.

But actually, in most of the carbon compounds, such as methane (CH4), carbon tetrachloride (CCl4) etc., all the four bonds of carbon atom are equivalent. The equivalent character of the bonds can be explained with the help of hybridization.

Hybridization may be defined as the phenomenon of intermixing of the orbitals of slightly different energies so as to redistribute their energies and to give new set of orbitals of equivalent energy and shape. The new orbitals formed as a result of hybridization are called hybrid or hybridized orbitals. Thus, to form four equivalent bonds, one 2s and three 2p-orbitals of carbon hybridize and form four new orbitals. Such orbitals are called sp3 hybrid orbitals.

 

The important characteristics of hybridization are listed below:

(i) The number of hybridized orbitals formed is equal to the number of orbitals that get hybridized.

(ii) The hybridized orbitals are always equivalent in energy and shape.

(iii) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.

(iv)The hybrid orbitals are directed in space in some preferred directions to have stable arrangements.

Therefore, the type of hybridization gives the geometry of the molecule. Depending upon the different combinations of s- and three p-orbitals, three types of hybridizations are known.

sp hybridization:                                       

This involves the mixing of one s- and one p-orbital forming two sp-hybrid orbitals. The two sp-hybrid orbitals are oriented in a linear arrangement and bond angle is 180°. For e.g., BeF2 involves sp-hybridization

sp2 hybridization:

In this case, one s- and two p-orbitals hybridize to form three sp2 hybrid orbitals. These three sp2 hybrid orbitals are oriented in a trigonal planar arrangement. For e.g., in BH3 boron atom undergoes sp2hybridization and therefore, BH3 has trigonal planar geometry and HBH bond angle is 120o.

 

sp3 hybridization:

In this case, one s- and three p-orbitals hybridize to form four sp3 hybrid orbitals. These four sp3-hybrid orbitals are oriented in a tetrahedral arrangement. The common example of molecule involving sp3-hybridisation is methane (CH4). Therefore, CH4 has tetrahedral geometry and HCH bond angle is 109.5o.

(i) sp hybridization

(ii) sp2 hybridization

(iii) sp3 hybridization

 

CONCEPT OF SIGMA BOND AND PI BOND

Sigma Bond (s)

When the overlap of orbital’s of two atoms takes place along the line joining the two nuclei (orbital axis) then the covalent bond formed is called sigma (s) bond. These bonds can be formed due to’s-s', 's-p' or 'p-p' overlap along the orbital axis. Free rotation around a sigma bond is always possible.


 

Fig: (iv)-

Formation of sigma bond due to various overlapping Pi bond (p)

When the two atoms overlap due to the sideways overlap of their 'p' orbitals, the covalent bond is called as pi(p) bond. In a pi bond the electron density is concentrated in the region perpendicular to the bond axis.

 

Characteristics of the pi (p) bonds                                          

-The pi bonds are weak because the orbtial overlap is partial.

-For a complete sideways overlap the 'p' orbitals should be parallel to each other. This is possible when all the atoms of the molecule are in the same plane i.e., there is no rotation of one part of the molecule relative to the other about the pi (p) bond.

 

 
Fig: (v) Rotation about a double bond

 

In molecules containing double bond, which has a pi bond, there is no free rotation and such molecules exist in isomeric forms of 'cis' and 'trans'. In the 'cis' form the similar atoms lie on the same side of a plane placed along the internuclear axis while in the 'trans' form the similar atoms lie on the opposite sides.

The electrons in the pi (p) bond are placed above and below the plane of the bonding atoms and so they are more exposed. They are more susceptible to attack by electron seeking or oxidizing agents. Hence, they are the most reactive centers in unsaturated (multiple bonded) compounds.

Comparative properties of sigma and pi bonds

Sigma (σ) bond

Pi (π) bond

Formed due to the axial overlap of two orbitals (‘s-s’, ‘s-p’or’p-p’).

Formed by the lateral (sideways) overlap of two ‘p’ orbitals.

Only one sigma bond exists between two atoms.

There can be more than one pi bonds between the two atoms.

The electron density is maximum and cylindrically symmetrical about the bond axis.

The electron density is high along the direction at right angles to the bond axis.

Free rotation about the sigma bond is possible.

Free rotation about the pi bond is not possible.

This bond can be independently formed, i.e., without the formation of a pi bond.

The pi bond is formed after the sigma bond has been formed,

Sigma bond is relatively strong.

Pi bond is a weak bond.

   

Valence shell Electron Pair Repulsion (VSEPR) theory:

The VSEPR theory, proposed by R.J.Gillespie and R.S. Nyholmm in 1957, is based on the repulsions between the electron-pairs in the valence-shell of the atoms in the molecule. It was developed to predict the shapes of the molecules in which the atoms are bonded together with single bonds only.


The main postulates of VSEPR theory are:

1. The shape of the molecule is determined by both the total number of electron pairs (bonding and non-bonding) around the molecules central atom and the orientation of these electron pairs in the space around the central atom.

2. In order to minimize the repulsion forces between them, electron pairs around the      molecules central atom, tend to stay as far away from each other as possible

3. Electron pairs around the molecule's central atom can be shared or can be lone pairs. The 'shared pairs' of electrons are also called bond pairs of electrons. The presence of lone pair(s) of electrons on the central atom causes some distortions in the expected regular shape of the molecule.

The strength of repulsions between different electron pairs follows the order:

Lone pair - Lone pair > Lone pair - Shared pair > Shared pair - Shared pair.

 

Prediction of molecular geometry on the basis of VSEPR theory:

 

Molecule with two bond pairs: In a molecule having two bond pairs of electrons around its central atom, the bond pairs are located on the opposite sides (at an angle of 180o) so that the repulsion between them is minimum. Such molecules are therefore linear. Some molecules, which show linear geometry are: BeF2 (beryllium fluoride), BeCl2(beryllium chloride), BeH2 (beryllium hydride), ZnCl2 (zinc chloride), and HgCl2(mercuric chloride)

Molecules with three bond pairs:

In a molecule having three bond pairs of electrons around its central atom, the electron pairs form an equilateral triangular arrangement around the central atom. These molecules have trigonal planar (or triangularplanar) shape and the three bond pairs are at 120°C with respect of each other.

In a molecule of the type AB3, the three bond pairs of electrons are located around A in a triangular arrangement and the molecule AB3, has a triangular planar geometry. Some molecules that show triangular planar geometry are BCl3, BF3, etc.

 

 

 

borontrifluoride is a trigonal planar molecule

 

Molecules with four bond pairs:

Molecule having four bond pairs of electrons around the central atom, arrange their electrons tetrahedrally. These molecules have tetrahedral shapes and the four bond pairs are at an angle of 109°28' with respect to each other. Some molecules which show tetrahedral geometry are CH4, CCl4, NH4+, SiH4 etc.

 

 

Molecules with five bond pairs:

Five bond pairs orient themselves around the central atom in a trigonalbipyramidal way. A molecule having five bond pairs around its central atom has a triangular bipyramidal shape. Three bond pairs are arranged in an equatorial triangular plane and are oriented at an angle of 120° with respect to each other. The other two bond pairs are opposite to each other, and at right angles to the triangular plane formed by the three bond pairs. Some other molecules, which show trigonalbipyramidal geometry are; PCl5, PF5, SbCl5.

For example in a molecule of the type AB5, the five bond pairs are distributed in a trigonalbipyramidal around the central atom 'A'. Therefore, the molecules of the type AB5 are trigonalbipyramidal in shape.

 


 
AB5, five bond pairs are oriented around A in a trigonalbipyramidalshape.Similarly,PCl5 also has a trigonalbipyramidal shape.
 

Molecules with six bond pairs:

Six bond pairs in a molecule are distributed octahedrally around the central atom. A molecule having six bond pairs around its central atom has an octahedral shape. In a molecule of the type AB6, the six 'B' atoms are placed octahedrally around 'A'. Thus, the molecules of the type AB6 are octahedral. The molecule SF6has an octahedral geometry

 

 

AB6 type molecules are octahedral in shape. SF6 molecule has an

octahedral geometry.

Shapes of the molecules having bond pairs and lone pairs of electrons

The pairs of electrons in the valence shell of an atom, which are not involved in bonding, are called lone pair of electrons. Well known instances of these types of molecules are: the oxygen atom in water molecule H2O, has two lone pairs of electrons; the nitrogen atom in ammonia molecule NH3 has one lone pair of electrons. Mentioned below are a few illustrative examples of these types of molecules.

 

Molecules with two bond pairs and two lone pairs:

The four electron pairs (two bond pairs + two lone pairs) are distributed tetrahedrally around the central atom as shown below. The two lone pairs on the central atom repel the bond pair slightly inwards due to greater lone pair-bond pair repulsion. As a result, the bond angle in such a molecule is less than the tetrahedral value of 109°28'. The presence of only two bonds in the molecules gives a bent (V-shaped) structure.

Examples include H2S, H2O, F2O and SCl2.

 

H2S: Bent (V-shaped) structure. S has 6 electrons in its outermost shell. 2H-atoms contribute 2 electrons during bonding. Thus, there are 8 electrons or 4 electron pairs around S. This gives a tetrahedral distribution of electron pairs around S. The two corners of the tetrahedron are occupied by H-atoms and the other two by the lone-pairs of electrons. Thus, H2S is a bent structure molecule

 

Molecules with four bond pairs and two lone pairs:

The four bond pairs are distributed in a planar distribution. The two lone pairs are in a direction at right angles to this plane. This gives a square planar shape to such molecules.

Examples include ICl-4, XeF4 and [Ni(CN) 4]2.

 


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