Question

The nitrogen atoms in \(N_{2}\) participate in multiple bonding, whereas those in hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4}\), do not. (a) Draw Lewis structures for both molecules. (b) What is the hybridization of the nitrogen atoms in each molecule? (c) Which molecule has the stronger \(\mathrm{N}-\mathrm{N}\) bond?

Short answer

(a) Lewis structures: N₂ has a triple bond between the nitrogen atoms (N≡N) each with one lone pair of electrons, while N₂H₄ has single bonds between nitrogen atoms (N-N) with each nitrogen bonded to two hydrogens and having one lone pair of electrons (H-N-N-H, with hydrogens single-bonded on the sides of nitrogen atoms). (b) Hybridization: For N₂, each nitrogen atom has an sp hybridization due to two electron groups (one triple bond and one lone pair). In N₂H₄, each nitrogen atom has sp³ hybridization with four electron groups (one single bond with another nitrogen, two single bonds with hydrogens, and one lone pair). (c) The N-N bond in N₂ is stronger due to the triple bond between the nitrogen atoms, while in N₂H₄, the N-N bond is a single bond.

Step-by-step solution

Drawing Lewis Structures

We will begin with drawing the Lewis structures for both molecules. We need to remember the valence electrons for each atom and distribute them accordingly. 1) N₂: Nitrogen has 5 valence electrons. In the N₂ molecule we have two nitrogen atoms, hence a total of 10 valence electrons. These electrons can be distributed as follows: N≡N Here, we have a triple bond between the two nitrogen atoms and each nitrogen atom has one lone pair of electrons, making a total of 10 valence electrons. 2) N₂H₄ (Hydrazine): In hydrazine, we have two nitrogen atoms and four hydrogen atoms. Hydrogen has 1 valence electron each, giving a total of 4 valence electrons for 4 hydrogen atoms. Nitrogen has 5 valence electrons, so two nitrogen atoms together have 10 valence electrons. Therefore, N₂H₄ has a total of 14 valence electrons. These electrons can be distributed as follows: H | N-N | | H H In this structure, the nitrogen atoms are single-bonded to one another and each nitrogen atom has two single bonds with two hydrogen atoms. Each nitrogen also has one lone pair of electrons, adding up to the total of 14 valence electrons.

Determine Hybridization of Nitrogen Atoms

To determine the hybridization, we will count the number of electron groups around each nitrogen atom (where "electron group" means either a single, double or triple bond, or a lone pair) and use the result to identify the hybridization in each case. 1) N₂: Each nitrogen atom has one triple bond (1 electron group) and one lone pair (1 electron group). Therefore, there are 2 electron groups around each nitrogen atom. This configuration corresponds to sp hybridization. 2) N₂H₄ (Hydrazine): Each nitrogen atom has one single bond with its neighbor nitrogen atom (1 electron group), two single bonds with two hydrogen atoms (2 electron groups), and one lone pair (1 electron group). Therefore, there are 4 electron groups around each nitrogen atom. This configuration corresponds to sp³ hybridization.

Compare N-N bond strength

To compare the strength of the N-N bonds in both molecules, we can use the information about their bonding: 1) N₂: The N-N bond in N₂ is a triple bond, which consists of 1 sigma bond and 2 pi bonds. 2) N₂H₄ (Hydrazine): In hydrazine, the N-N bond is a single sigma bond. Triple bonds are stronger than single bonds. Thus, the N-N bond in the N₂ molecule is stronger than the N-N bond in the N₂H₄ (hydrazine) molecule.

Key concepts

Lewis Structures

Lewis structures are a useful way to visualize the distribution of valence electrons in molecules. They provide insight into how atoms bond and the shape of a molecule. Let's examine the Lewis structures for nitrogen (\(N_2\)) and hydrazine (\(N_2H_4\)).

- For \(N_2\), each nitrogen atom has 5 valence electrons, making a total of 10 for the molecule. The Lewis structure shows a triple bond between the nitrogen atoms, which means there are three shared pairs of electrons. Each nitrogen also has one lone pair of electrons. This structure completes each nitrogen atom's valence shell, achieving the stable configuration of 8 electrons.

- In hydrazine (\(N_2H_4\)), there are two nitrogen atoms and four hydrogen atoms. Hydrogens contribute 4 valence electrons (1 each), and nitrogen offers 10. Therefore, hydrazine has 14 electrons available for bonding. The Lewis structure depicts each nitrogen single bonded to another nitrogen atom and also bonded to two hydrogen atoms, with lone pairs remaining on each nitrogen. This setup results in a single bond between the two nitrogen atoms, differing significantly from the triple bond structure in \(N_2\).

Hybridization

The concept of hybridization helps us understand the bonding and geometry around central atoms in a molecule. It involves the mixing of atomic orbitals to form new hybrid orbitals which can then bond with other atoms.

- In \(N_2\), hybridization of the nitrogen atoms occurs as \(sp\). Each nitrogen atom in \(N_2\) forms one triple bond and has a lone pair, totaling two electron groups. The \(sp\) hybridization is a result of mixing one s orbital with one p orbital, forming two identical \(sp\) hybrid orbitals, ideal for linear geometry, explaining the strong linear triple bonds.

- In contrast, hydrazine (\(N_2H_4\)) displays \(sp^3\) hybridization for its nitrogen atoms. Each nitrogen forms single bonds with one another, two hydrogen atoms, and has one lone pair, totaling four electron groups. This \(sp^3\) hybridization results from mixing one s orbital with three p orbitals, giving four equivalent orbitals that adapt to a tetrahedral shape, which matches the three-dimensional structure of the molecule.

Bond Strength

Bond strength refers to how strongly atoms are held together. It is crucial in determining a molecule's stability and reactivity.

- In a nitrogen molecule (\(N_2\)), the \(N-N\) bond consists of a triple bond. This bond type comprises one sigma bond and two pi bonds. Triple bonds are the strongest of covalent bonds because more electrons are shared between the atoms, creating a stronger attraction.

- On the other hand, hydrazine (\(N_2H_4\)) contains a single \(N-N\) bond, which incorporates just one sigma bond with no pi components. Single bonds are typically weaker than multiple bonds because they involve fewer shared electrons. Consequently, the \(N-N\) bond in \(N_2\) is significantly stronger than in \(N_2H_4\).

Understanding these differences in bond types and strengths sheds light on why some molecules are more stable or reactive than others. The \(N_2\) molecule, with its robust triple bond, is much more resilient than hydrazine, which has weaker single bonds.