Question

Nitromethane has the formula \(\mathrm{CH}_{3} \mathrm{NO}_{2}\), with the \(\mathrm{N}\) bonded to the \(\mathrm{C}\) and without \(\mathrm{O}-\mathrm{O}\) bonds. Draw its two most important contributing structures. a. What is the hybridization of the C, and how many hybrid orbitals are in the molecule? b. What is the shortest bond? c. Between which two atoms is the strongest bond found? d. Predict whether the HCH bond angles are greater or less than \(109.5^{\circ}\) and justify your prediction.

Short answer

Carbon is sp3 hybridized with 8 hybrid orbitals in the molecule, the nitrogen-oxygen double bond is the shortest and the strongest, and the HCH bond angles are predicted to be slightly less than 109.5 degrees due to steric hindrance from the NO}_2{)} group.

Step-by-step solution

Determine the Lewis Structures

Identify the valence electrons for each atom in nitromethane (CH}_3NO}_2{)}). Carbon has 4, nitrogen has 5, each hydrogen has 1, and each oxygen has 6. To draw the Lewis structures, connect each hydrogen to carbon, then connect nitrogen to carbon and add two oxygens to nitrogen. Distribute the remaining electrons to satisfy the octet rule for nitrogen and oxygen. The most important contributing structures have nitrogen with one single and one double bond to the two oxygens, and a single bond to carbon. There will be resonance between the nitrogen-oxygen bonds.

Determine the Hybridization of Carbon

Carbon is bonded to three hydrogens and one nitrogen. Thus it must use four orbitals for its bonds, which means it is sp3 hybridized.

Count Total Hybrid Orbitals in the Molecule

In nitromethane, every atom except hydrogen is hybridized. The carbon atom provides 4 sp3 orbitals, nitrogen provides 4 sp2 or sp3 (depending on resonance structure), for a total of 8 hybrid orbitals.

Identify the Shortest Bond

Typically, the double bond is shorter than the single bond. Therefore, the nitrogen-oxygen double bond will be the shortest.

Determine the Strongest Bond

Double bonds are generally stronger than single bonds because they have more electron density between the atoms. In nitromethane, the nitrogen-oxygen double bond is the strongest.

Predict the HCH Bond Angles

Since the carbon atom is sp3 hybridized, the ideal bond angles would be 109.5 degrees. However, due to the presence of the NO}_2{)} group which can cause steric hindrance, the HCH bond angles might be slightly less than 109.5 degrees.

Key concepts

Resonance Structures

Understanding resonance structures is key to grasping the behavior of molecules like nitromethane ((CH}_3NO}_2{)}). Resonance structures depict different ways electrons can be distributed within a molecule. They are not real, individual structures but rather a blend of multiple possibilities that together represent the molecule's actual electron distribution. For nitromethane, the resonance involves the nitrogen and oxygen atoms.

When depicting these structures, one might draw nitrogen with a single bond to one oxygen and a double bond to the other. Another structure will flip this arrangement. These drawings demonstrate that the actual molecule is a hybrid of these extremes, with the bond order to each oxygen being somewhere between a single and a double bond. As a result, the nitrogen-oxygen bonds are of equal strength and length, which is characteristic of resonance in a molecule.

For students, visualizing this concept can be challenging. Simplifying, think of resonance as a 'sharing' arrangement where electrons are not localized but rather are spread out or 'delocalized' over multiple atoms, stabilizing the molecule overall.

sp3 Hybridization

The concept of sp3 hybridization reveals how atoms like carbon in nitromethane can form four equivalent bonds despite having different orbitals available for bonding. In the context of nitromethane, the carbon atom must form bonds with three hydrogen atoms and one nitrogen atom. This atom has one s and three p orbitals in its outer shell, which mix to form four equivalent sp3 hybrid orbitals.

This hybridization process results in four orbitals that are equal in energy and oriented tetrahedrally to minimize repulsion between them. It's essential to understand that these are 'mixed' orbitals, combining characteristics of both s and p orbitals. This hybridization allows carbon to form four sigma (σ) bonds, one with each adjacent atom. The term sp3 derives from one s orbital merging with the three available p orbitals.

Remember, in layman's terms, hybridization is like making a new recipe by mixing different ingredients to create something better suited for a specific purpose—in this case, four identical orbitals that allow carbon to bond effectively with its neighbors.

Bond Angles

Bond angles play a significant role in defining the shape and structure of molecules. In an ideal sp3 hybridized molecule like methane (CH}_4{)}), the bond angles are 109.5 degrees, which is the angle that gives the atoms the most space from each other due to repulsion forces. This is the tetrahedral angle.

However, in molecules like nitromethane where substituents are not all the same, such as the NO}_2{)} group attached to the sp3 hybridized carbon, these angles can deviate from the ideal. The NO}_2{)} group introduces electron-withdrawing effects and steric hindrance—physical crowding by larger substituents—which can reduce the bond angle slightly from the ideal tetrahedral angle.

In summary, while the HCH bond angles in nitromethane may still be close, they are likely to be somewhat less than the perfect 109.5 degrees due to the presence of the NO}_2{)} group. This deviation from the ideal angle can lead to varying reactivity and physical properties, so understanding this concept is important for predicting the behavior of organic molecules.