Q.(i) Discuss the concept of hybridisation. What are its different types in a carbon atom.

(i) Concept of Hybridisation

Hybridisation is the mixing of atomic orbitals of similar energy (e.g., 2s and 2p) to form an equal number of new, equivalent hybrid orbitals with specific geometry. In carbon, the three common types are:

• sp³ — four hybrid orbitals, tetrahedral geometry (e.g., in alkanes)
• sp² — three hybrid orbitals, trigonal planar geometry (e.g., in alkenes)
• sp — two hybrid orbitals, linear geometry (e.g., in alkynes)

(ii) Hybridisation of starred carbon atoms

Count sigma bonds and lone pairs around each starred carbon:

• (a) *CH₂=CH—C(=O)—O—H Starred carbon in CH₂= is bonded to 3 atoms (2 H, 1 C) — sp²
• (b) CH₃—*CH₂—OH Starred carbon bonded to 4 atoms (2 H, 1 C, 1 O) — sp³ …

The Concept of Hybridisation

Hybridisation is a theoretical model that explains how atoms, especially carbon, form bonds with specific geometries. The central idea is that atomic orbitals (like $2s$ and $2p$) of similar energy mix together to create new, equivalent hybrid orbitals. These hybrid orbitals are directed in space to minimise repulsion, giving the molecule its shape.

For carbon, the ground state electron configuration is $1s^2 2s^2 2p_x^1 2p_y^1$. With only two unpaired electrons, carbon would form only two bonds — but we know it forms four. Hybridisation resolves this: one $2s$ electron is promoted to the empty $2p_z$ orbital, giving four unpaired electrons ($2s^1 2p_x^1 2p_y^1 2p_z^1$). These four orbitals then mix to form four equivalent $sp^3$ hybrid orbitals, each with 25% $s$-character and 75% $p$-character, arranged tetrahedrally at $109.5^\circ$.

The type of hybridisation depends on the number of sigma bonds and lone pairs around the atom. The key rule: count the number of sigma bonds and lone pairs — that number equals the number of hybrid orbitals needed.

A double bond consists of one sigma bond and one pi bond; a triple bond has one sigma and two pi bonds. Pi bonds use unhybridised $p$ orbitals, so they do not count toward the steric number.

Types of Hybridisation in Carbon

Carbon can exhibit three types:

1. $sp^3$ hybridisation: One $s$ and three $p$ orbitals mix to form four $sp^3$ orbitals. Each forms a sigma bond. Geometry: tetrahedral, bond angle $109.5^\circ$. Example: methane ($CH_4$).

2. $sp^2$ hybridisation: One $s$ and two $p$ orbitals mix to form three $sp^2$ orbitals. The remaining unhybridised $p$ orbital forms a pi bond. Geometry: trigonal planar, bond angle $120^\circ$. Example: ethene ($CH_2=CH_2$).

3. $sp$ hybridisation: One $s$ and one $p$ orbital mix to form two $sp$ orbitals. Two unhybridised $p$ orbitals form two pi bonds. Geometry: linear, bond angle $180^\circ$. Example: ethyne ($HC \equiv CH$).

Determining Hybridisation of Starred Carbons

Now we apply this to each molecule. For each starred carbon, count its sigma bonds (single bonds count as one sigma; double bonds have one sigma and one pi; triple bonds have one sigma and two pi). No lone pairs on carbon here.

(a) $^*CH_2 = CH — ^*C(=O) — O — H$

• First starred carbon ($^*CH_2$): It is part of a double bond with the next carbon. It forms three sigma bonds (two to H atoms, one to the adjacent C) and one pi bond (the second bond of the double bond). Steric number = 3 sigma bonds = 3. Hybridisation: $sp^2$.
• Second starred carbon ($^*C(=O)$): It has a double bond to oxygen (one sigma, one pi) and a single bond to the adjacent carbon (sigma) and a single bond to the O in the —O—H group (sigma). That's three sigma bonds total. Steric number = 3. Hybridisation: $sp^2$. …