Question

The type of hybridisation present in \(\mathrm{SO}_{2}\) and \(\mathrm{SO}_{3}\) is respectively (1) \(\mathrm{sp}, \mathrm{sp}^{2}\) (2) \(\mathrm{sp}^{2}, \mathrm{sp}^{2}\) (3) \(\mathrm{sp}^{2}, \mathrm{sp}^{3}\) (4) \(\mathrm{sp}, \mathrm{sp}^{3}\)

Short answer

Option (2): \(\text{sp}^2, \text{sp}^2\)

Step-by-step solution

Determine the hybridisation of \(\text{SO}_2\)

In \(\text{SO}_2\), sulfur forms two double bonds with oxygen atoms. The electronic geometry around the sulfur atom is such that it includes one lone pair of electrons, resulting in a total of three regions of electron density. This corresponds to \(\text{sp}^2\)-hybridisation.

Verify the hybridisation for \(\text{SO}_2\)

Each double bond counts as one region of electron density and the lone pair adds the third region. Therefore, the hybridisation around the sulfur atom in \(\text{SO}_2\) is indeed \(\text{sp}^2\).

Determine the hybridisation of \(\text{SO}_3\)

In \(\text{SO}_3\), sulfur forms three double bonds with three oxygen atoms. This results in three regions of electron density around the central sulfur atom. The absence of lone pairs on sulfur in this case means the geometry is trigonal planar, leading to \(\text{sp}^2\)-hybridisation.

Verify the hybridisation for \(\text{SO}_3\)

Counting the three double bonds as three regions of electron density and ensuring there are no lone pairs, the hybridisation of the sulfur atom in \(\text{SO}_3\) is confirmed to be \(\text{sp}^2\).

Match with given choices

The hybridisation types found for \(\text{SO}_2\) and \(\text{SO}_3\) are both \(\text{sp}^2\). Therefore, the correct option is (2) \(\text{sp}^2, \text{sp}^2\).

Key concepts

sp2 hybridisation

In chemistry, hybridisation is the concept of mixing atomic orbitals to form new hybrid orbitals. These hybrid orbitals can explain the observed molecular geometries more precisely. One common type of hybridisation is \(\text{sp}^2\)-hybridisation. This occurs when one s-orbital and two p-orbitals from the same atom combine. The resulting three hybrid orbitals are equivalent in energy and shape. They arrange themselves in a trigonal planar geometry, with an angle of 120° between them. This arrangement minimizes electron repulsion. In \(\text{sp}^2\)-hybridisation, each hybrid orbital forms a sigma bond with another atom, while any remaining p-orbital may form a pi bond.

electron density regions

Electron density regions around an atom indicate where electrons are most likely to be found. These regions are essential in determining the hybridisation and geometry of a molecule. Each lone pair, single bond, double bond, or triple bond counts as one region of electron density. When determining the hybridisation of an atom, you count the total number of these regions. For example, if there are three regions, \(\text{sp}^2\)-hybridisation is likely. This aspect helps in predicting molecular shapes accurately. The VSEPR (Valence Shell Electron Pair Repulsion) theory is instrumental in understanding how electron density regions influence the geometry of a molecule.

sulfur dioxide hybridisation

Sulfur dioxide (\text{SO}_2) is a molecule where the sulfur atom undergoes \(\text{sp}^2\)-hybridisation. In \(\text{SO}_2\), sulfur forms two double bonds with oxygen atoms. There is also one lone pair on sulfur. This results in a total of three regions of electron density. These regions adopt a trigonal planar configuration to minimize repulsion. The presence of one lone pair slightly distorts the geometry, making it a bent shape rather than perfectly planar. Despite this, the \(\text{sp}^2\)-hybridisation remains. Each \(\text{sp}^2\) orbital overlaps with an oxygen atom's orbital to form a sigma bond, while the remaining p-orbital forms a pi bond with each oxygen.

sulfur trioxide hybridisation

Sulfur trioxide (\text{SO}_3) showcases another instance of \(\text{sp}^2\)-hybridisation. Here, the sulfur atom forms three double bonds with three oxygen atoms. Unlike \(\text{SO}_2\), there are no lone pairs on the sulfur atom in \(\text{SO}_3\). This results in three regions of electron density, forming a trigonal planar geometry. The absence of lone pairs simplifies the geometry, making it perfectly trigonal planar with 120° bond angles. In this case, the sulfur atom's three \(\text{sp}^2\) hybrid orbitals each form a sigma bond with an oxygen atom. The p-orbital on sulfur forms pi bonds with the p-orbitals of the oxygen atoms, completing the double bonds.