Discuss in brief sp3 hybridisation. Explain the formation of methane and ethane.
The four C – H bonds are directed towards the four corners of a regular tetrahedron. So methane has a tetrahedral structure. Each H – C – H bond angle is of 109°.28’. Each C – H bond length is 109 pm (1.09 Å).
(ii) Molecular orbital picture of ethane: In ethane molecule, both carbon atoms are in the sp3 hybrid state. In its formation, one hybrid orbital of one carbon atom overlaps with one sp3 hybrid orbital of a second carbon atom along the internuclear axis to form a sigma (σ) C – C bond. The remaining three sp3 hybrid orbitals of each carbon atom overlap with 1 s orbital of hydrogen atom axially to form six sigma C – H bonds.
The length of C-C bond in ethane is 154 pm (or 1- 54 Å) and that of each C - H bond is 109 pm (or 1-09 Å).
With the help of VSEPR theory, explain the shape of: (i) NH3 (ii) H2O.
distortion in the geometry of the molecule. The lone pairs repel the bond pairs more effectively resulting in the decrease of H – O – H angle from 109.5° to 104.5°. The water molecule may be regarded as bent or angular or V-shaped.
On the basis of VSEPR theory, discuss the geometry of the following covalent molecules: (i) BeF2 (ii) BF3 (iii) CH4.
On the basis of VSEPR theory, predict the shapes of given molecules: PCl5 and SF6.
(i) PCl5 molecule. The electronic configuration of central phosphorus atom is
It has five valence electrons. All the five electrons are mutually shared with the electrons of five chlorine atoms to form five P - Cl bonds as shown.
Hence, P atom is surrounded by five shared pairs of electrons. These repel each other and take up such positions and there is no further repulsion between them. The most favourable arrangement is trigonal bipyramidal.
(ii) SF6 molecule: The electronic configuration of central sulphur atom is
It has six valence electrons.
All the six valence electrons are mutually shared by the electrons of six fluorine atoms to form six S – F bonds.
Hence, S atom is surrounded by six shared pairs of electrons (six bond pairs). These repel each other and try to remain as far apart as possible so that there is no further repulsion between them. Under such conditions, the most favourable arrangement is octahedral.
State and explain hybridisation.
The phenomenon of intermixing of atomic orbitals of slightly different enthalpies of an atom so as to redistribute their enthalpies to form the same number of new orbitals of equivalent enthalpies and identical shapes is called hybridization. The new orbitals, thus formed are called hybrid orbitals or hybridised orbitals.
Explanation : In order to understand hybridisation, let us take an example of carbon (Z= 6). Its ground state electronic configuration is,
Since it has two half filled orbitals, therefore, the valency of the carbon atom should be 2. But actually, carbon atom always exhibits a valency of four (tetravalent). To achieve this, an electron is promoted from 2s filled orbital to the vacant higher enthalpy 2p orbital. This is called excited state of a carbon atom.
In the excited state of carbon s and p, orbitals have different enthalpies. Consequently, four bonds of carbon must be of two types. Three of the bonds should be of one type (s - p bonds) while fourth bond should be a different type (s - s bond). However, experimental evidence indicates that all the four bonds in case of CH4 (methane) are equivalent. To explain the equivalence of all the four bonds in case of methane, the concept of hybridisation is used i.e. all the four orbitals in the valence shell of carbon may get mixed, redistribute enthalpies and give orbitals of new enthalpy and shape. These equivalent orbitals are called hybrid orbitals.