Question

The molecule formamide, \(\mathrm{HCONH}_{2}\), has the approximate bond angles \(\mathrm{H}-\mathrm{C}-\mathrm{O}, 123^{\circ} ; \mathrm{H}-\mathrm{C}-\mathrm{N}, 113^{\circ}\) \(\mathrm{N}-\mathrm{C}-\mathrm{O}, 124^{\circ} ; \mathrm{C}-\mathrm{N}-\mathrm{H}, 119^{\circ} ; \mathrm{H}-\mathrm{N}-\mathrm{H}, 119^{\circ}\) The \(\mathrm{C}-\mathrm{N}\) bond length is \(138 \mathrm{pm}\). Two Lewis structures can be written for this molecule, with the true structure being a resonance hybrid of the two. Propose a hybridization and bonding scheme for each structure.

Short answer

The bonding scheme for the first structure involves an \(sp^2\) hybridization at the Carbon (C) forming a sigma bond with Nitrogen (N) which is \(sp^3\) hybridized in this case and a sigma bond and pi bond with oxygen which is \(sp^2\) hybridized. In case of the second structure, the hybridization at Carbon (C) is still \(sp^2\), but it forms a sigma bond with the \(sp^2\) hybridized Nitrogen (N) and sigma bond with the \(sp^3\) hybridized oxygen.

Step-by-step solution

Draw the Lewis Structures

Lewis structure for Formamide, \(HCONH_{2}\) can function in two ways: with oxygen double bonded to carbon and nitrogen single bonded to carbon, or with nitrogen double bonded to carbon and the oxygen single bonded to carbon. In both cases, each hydrogen atom forms a single bond with either carbon or nitrogen. In Lewis notation: 1. Structure 1: \(H-C=O / \backslash NH_{2}\) 2. Structure 2: \(H-C-N=O / \backslash H\)

Proposed hybridization scheme for Structure 1

In this case, the carbon atom is bonded to oxygen through a double bond and to nitrogen via a single bond. This suggests that the hybridization at Carbon (C) is \(sp^2\), with one of the \(sp^2\) orbitals forming a sigma bond with the nitrogen's \(sp^3\) hybrid orbital (N has an \(sp^3\) because it forms three sigma bonds and one lone pair), and two \(sp^2\) orbitals and the remaining p orbital overlaping with oxygen's orbitals to form a sigma and two pi bonds respectively. Thus, for Oxygen (O), the hybridization is \(sp^2\) (two sigma bonds and one lone pair), while for Nitrogen (N), it is \(sp^3\).

Proposed hybridization scheme for Structure 2

In the second structure, the carbon atom is bonded to nitrogen through a double bond and to oxygen via a single bond. This suggests that the hybridization at Carbon (C) is \(sp^2\), much like in Structure 1. For Nitrogen (N), one of the \(sp^2\) orbitals forms a sigma bond with the carbon's \(sp^2\) hybrid orbital, and two \(sp^2\) orbitals along with the remaining p orbital overlap with the hydrogen's orbitals to form sigma bonds, hence the Nitrogen (N) is also \(sp^2\) hybridized. In case of oxygen, it forms two sigma bonds (with carbon and hydrogen) and therefore is \(sp^3\) hybridized.

Key concepts

Lewis Structures

When drawing the Lewis structures for formamide (\(\mathrm{HCONH}_{2}\)), we're essentially outlining potential ways the molecule's atoms can be bonded. There are two main possibilities often considered:

• In the first structure, the oxygen forms a double bond with carbon (\(\mathrm{C=O}\)), while nitrogen forms a single bond with carbon (\(\mathrm{C-NH}_2\)). Each hydrogen forms single bonds with either carbon or nitrogen.
• In the second structure, a different configuration arises: nitrogen forms a double bond with carbon (\(\mathrm{C=N}\)), while oxygen forms a single bond with carbon (\(\mathrm{C-OH}\)).

In both structures, we visualize electrons as dots around atoms to show possible electron pairings. Lewis structures provide a visual way to understand which electrons participate in bonding and which act as lone pairs.

Bonding Scheme

In formamide, each Lewis structure suggests a distinct bonding scheme, which refers to how atoms share electrons through overlapping orbitals. Let's break this down:
For Structure 1, where oxygen and carbon form a double bond, the carbon atom is using its\(sp^2\) hybrid orbitals. These arrange in a flat, trigonal shape, making room for oxygen's\(sp^2\) orbitals to overlap with carbon's, forming a sigma bond alongside a double bond with electrons occupying an unhybridized p orbital.

• Oxygen, with its two lone pairs, also exhibits\(sp^2\) hybridization.
• Nitrogen forms single bonds using\(sp^3\) hybridization due to its lone pair and three bound neighbors.

Switching to Structure 2, nitrogen now forms a double bond with carbon:

• Nitrogen's\(sp^2\) orbitals participate in the creation of a sigma and pi bond.
• Oxygen shifts to an\(sp^3\) hybridization since it predominantly forms single bonds in this scenario.

Overall, these bonding schemes describe how atoms hybridize their orbitals to create stable molecules.

Resonance Hybrid

The term "resonance hybrid" refers to the actual structure of a molecule, which is a blend of multiple valid Lewis structures. In formamide's case, the two Lewis structures depict electron placement differently, but neither gives the full picture of the atomic connectivity or electron distribution by itself.
Resonance hybrids stabilize a molecule by allowing electrons to delocalize, or spread out, over more atoms rather than being restricted to a particular bond. For formamide:

• The presence of a resonance hybrid generally lowers the overall energy of the molecule, contributing to its stability.
• The true structure is a "blend" that averages physical and chemical properties predicted by both individual resonance forms.

This distribution leads to bond characteristics like partial double bond qualities between the atoms involved, despite some bonds being single within traditional Lewis depictions.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. For formamide, the molecule displays different bond angles due to its geometry.
The geometry can be understood by looking at bond angles:

• \(\mathrm{H-C-O}\) angles are approximately\(123^{\circ}\), indicating a trigonal planar arrangement around carbon in both structural forms.
• In structure 1, carbon experiences\(sp^2\) hybridization, leading to a similar planar shape with reflected angles around oxygen (double bonded).
• Nitrogen's involvement brings angles like\(119^{\circ}\) in the\(\mathrm{C-N-H}\) bonding, supporting a bent pyramid shape.

Each atom's position results from these combined hybridizations, aiding in understanding molecular shape and predicting reactiveness based on electron distribution.