Question

State the hybridization of the atom in boldface type in each of the following molecules: (a) \(\mathrm{SF}_{6}\); (b) \(\mathrm{O}_{3} \mathrm{Cl}-\mathrm{O}-\mathrm{ClO}_{3}\); (c) \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COOH}\) (glycine); (d) \(\mathrm{OC}\left(\mathrm{NH}_{2}\right)_{2}\) (urea).

Short answer

The hybridizations are: (a) SF6 is sp3d2 hybridized, (b) the central Cl in O3Cl-ClO3 is sp3 hybridized, (c) N in glycine is sp3 hybridized, C in COOH is sp2 hybridized, (d) C in urea is sp2 hybridized and N in urea is sp3 hybridized.

Step-by-step solution

Determine the Hybridization of Sulfur in SF6

To find the hybridization of sulfur in SF6, count the number of bonds and lone pairs around the sulfur atom. Sulfur forms 6 bonds to fluorine atoms and has no lone pairs. Using the formula for hybridization, sp^n, where n is the number of sigma bonds plus the number of lone pairs, we get sp^3d^2. Therefore, sulfur in SF6 is sp3d2 hybridized.

Determine the Hybridization of the Central Chlorine in O3Cl-ClO3

For the central chlorine atom in O3Cl-ClO3, there are 3 sigma bonds and 1 lone pair. Using the hybridization formula sp^n, we find sp^3. Thus, the hybridization is sp3.

Determine the Hybridization of Nitrogen in Glycine

In glycine, the amino nitrogen has 3 sigma bonds (two to hydrogen and one to carbon) and 1 lone pair. The hybridization is thus sp^3, indicating sp3 hybridization.

Determine the Hybridization of Carbon in the COOH Group

The carbon in the COOH group is double-bonded to one oxygen (one sigma and one pi bond) and single-bonded to an OH group and to nitrogen, making a total of 3 sigma bonds and no lone pairs. This results in sp^2 hybridization.

Determine the Hybridization of Carbon and Nitrogen in Urea

In urea, the carbon is double-bonded to an oxygen (one sigma and one pi bond) and single-bonded to two nitrogen atoms, resulting in three sigma bonds. The carbon atom is sp^2 hybridized. The nitrogen atoms in urea each have two sigma bonds and one lone pair, indicating sp^3 hybridization for both nitrogen atoms in urea.

Key concepts

Chemical Bonding

Chemical bonding is the force that holds atoms together in a molecule. Atoms bond by either transferring electrons (ionic bonding) or sharing electrons (covalent bonding). In the example of sulfur hexafluoride (\textbf{\textsf{SF}}\(_6\)), the sulfur and fluorine atoms form covalent bonds. Covalent bonding occurs when pairs of electrons are shared by atoms, and the strength of the bond depends on the attraction between the atoms' nuclei and the shared electrons.

Ionic bonding, on the other hand, involves the transfer of electrons from one atom to another, resulting in positively and negatively charged ions which attract each other. For the molecules in question, the covalent bonds are of particular significance because they result in specific hybridizations and molecular geometries.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It's crucial for predicting both the physical properties and reactivity of the molecule. For sulfur hexafluoride (\textbf{\textsf{SF}}\(_6\)), with its sp\(^3\)d\(^2\) hybridization, the molecular geometry is octahedral. This means the six fluorine atoms are equally spaced around the central sulfur atom, forming a shape similar to two pyramids base to base.

In glycine, the central carbon atom in the \textbf{\textsf{COOH}} group is connected to three other atoms and has no lone pairs, which leads to a trigonal planar geometry due to its sp\(^2\) hybridization. Understanding the hybridization helps us predict the molecular geometry, which is essential in the study of molecules and their interactions, including complex molecules like proteins and DNA.

Valence Bond Theory

Valence Bond Theory explains how atoms bond together to form molecules. It posits that covalent bonds are formed when the atomic orbitals of two atoms overlap and that the electrons in these overlapping orbitals are shared between atoms. The concept of hybridization comes from this theory. It describes how atomic orbitals mix to form new hybrid orbitals that make sufficient overlap to form covalent bonds.

For instance, the nitrogen in glycine (\textbf{\textsf{H\(_2\)N-CH\(_2\)-COOH}}), by using its sp\(^3\) hybrid orbitals, forms a tetrahedral geometry, contributing to the molecule's three-dimensional structure. Valence Bond Theory also acknowledges the existence of unpaired electrons and lone pairs in the determination of molecular shape and reactivity. By applying Valence Bond Theory, we can infer the hybridization states of molecules, as seen in the provided examples, which help us understand the intricate details of molecular bonding.