Question

Use the Lewis structure model to deduce the type of nitrogen-nitrogen bond present in (a) \(\mathrm{N}_{2} \mathrm{H}_{4},\) (b) \(\mathrm{N}_{2} \mathrm{F}_{4}\) (c) \(\mathrm{N}_{2} \mathrm{F}_{2}\) and (d) \(\left[\mathrm{N}_{2} \mathrm{H}_{5}\right]^{+}\)

Short answer

N₂H₄ and N₂F₄ have single N–N bonds; N₂F₂ has a double N=N bond; [N₂H₅]⁺ has a single N–N bond.

Step-by-step solution

Determine Valence Electrons

To determine the type of nitrogen-nitrogen bond, first, calculate the total number of valence electrons available for each molecule. - **N₂H₄**: Nitrogen has 5 valence electrons each, and hydrogen has 1. Thus, for N₂H₄, we have \( 2 \times 5 + 4 \times 1 = 14 \) valence electrons.- **N₂F₄**: Nitrogen has 5 valence electrons each, and fluorine has 7. Thus, for N₂F₄, we have \( 2 \times 5 + 4 \times 7 = 34 \) valence electrons.- **N₂F₂**: Similarly, for N₂F₂, we have \( 2 \times 5 + 2 \times 7 = 24 \) valence electrons.- **[N₂H₅]⁺**: Nitrogen has 5 valence electrons each, and hydrogen has 1. Additionally, due to the positive charge, subtract 1 electron. Thus, we have \( 2 \times 5 + 5 \times 1 - 1 = 14 \) valence electrons.

Sketch Basic Lewis Structures

Based on total valence electrons, draw the basic Lewis structures for each molecule: - **N₂H₄**: Connect two nitrogen atoms with a single bond (N–N) and satisfy each nitrogen with three hydrogen atoms each. - **N₂F₄**: Connect two nitrogen atoms with a single bond (N–N) and attach two fluorine atoms to each nitrogen atom. - **N₂F₂**: Connect the two nitrogen atoms with a double bond (N=N) and attach one fluorine atom to each nitrogen atom. - **[N₂H₅]⁺**: Connect the two nitrogen atoms with a single bond (N–N) and attach a set of hydrogens, distributed evenly to give each nitrogen atom a single bonded hydrogen, one nitrogen will have three hydrogens forming the cationic structure.

Analyze Bond Types in Structures

Observe the Lewis structures to identify nitrogen-nitrogen bond orders based on shared electron pairs: - **N₂H₄**: Displays a single N–N bond with shared lone electron pairs creating a stable single bond. - **N₂F₄**: Has a single N–N bond where each nitrogen is saturated with its other bonded atoms. - **N₂F₂**: Has a double N=N bond to meet the octet rule with shared electrons. - **[N₂H₅]⁺**: Contains a single N–N bond as all valence electrons are paired with hydrogen while maintaining a positive charge.

Key concepts

Understanding Nitrogen-Nitrogen Bonds

Nitrogen-nitrogen bonds are the connections between two nitrogen atoms in a molecule. These bonds can be classified based on how many pairs of electrons are shared between the atoms. There are three main types:

• Single bond (N–N): One pair of electrons is shared. This is the weakest type of nitrogen-nitrogen bond and typically found in compounds like hydrazine (N₂H₄).
• Double bond (N=N): Two pairs of electrons are shared, providing a stronger bond. This is found in molecules like difluorodiazine (N₂F₂).
• Triple bond (N≡N): This bond involves three shared pairs of electrons and is the strongest, as seen in nitrogen gas (N₂). However, it does not appear in the given exercise compounds.

Each bond type affects the molecule's stability and reactivity. Recognizing these bonds is crucial for understanding molecular behavior and interactions.

The Role of Valence Electrons

Valence electrons are the outermost electrons of an atom and play a fundamental role in bond formation. When drawing Lewis structures, the first step is to determine the total number of valence electrons available for bonding. This is done by adding the valence electrons of each atom and considering any charge present.

• For nitrogen, there are 5 valence electrons per atom.
• For hydrogen, there is 1 valence electron per atom.
• Fluorine has 7 valence electrons per atom.

In the given compounds, calculating the available valence electrons helps predict possible molecular structures:

• N₂H₄ has 14 valence electrons.
• N₂F₄ has 34 valence electrons.
• N₂F₂ has 24 valence electrons.
• [N₂H₅]⁺ has 14 valence electrons after accounting for its positive charge.

Understanding valence electrons aids in anticipating bond formation and molecular polarity.

Exploring Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule. It is influenced by the types of bonds and the number of lone pairs around the central atom. Here are some typical geometries resulting from the compounds in the exercise:

• N₂H₄ (hydrazine) is described as having a "bent" or "angular" geometry due to its single N–N bond and surrounding hydrogen atoms.
• N₂F₄ tends to have a "staggered" or "twisted" conformation to minimize repulsion between large fluorine atoms.
• N₂F₂, with its double N=N bond, typically exhibits a planar geometry, maintaining equal spacing amongst atoms.
• [N₂H₅]⁺, considering the positive charge and single N–N bond, forms a somewhat asymmetrical geometry with the attached hydrogens.

These shapes are crucial to predicting how a molecule might interact with others, and can influence properties like solubility and boiling points.

Deciphering Bond Orders

The bond order refers to the number of chemical bonds between a pair of atoms. In nitrogen-containing compounds, it tells us how strongly two nitrogen atoms are linked.

• A single bond has a bond order of 1, as there's only one pair of shared electrons, seen in N₂H₄ and N₂F₄.
• A double bond, with a bond order of 2, indicates two pairs of shared electrons, as observed in N₂F₂.

Higher bond orders generally correspond to shorter and stronger bonds. By evaluating these bond orders, chemists can deduce a wealth of information about molecular stability, reactivity, and overall chemical behavior. This insight is vital for applications in synthesis and materials science.