Question

A pair of isoelectronic species for \(C\) and \(N\) exist with the formula \(\mathrm{X}_{2} \mathrm{O}_{4}\) in which there is an \(\mathrm{X}-\mathrm{X}\) bond. \(\mathrm{A}\) corresponding fluoride of boron also exists. Draw Lewis structures for these species and describe their shapes.

Short answer

The isoelectronic species for N2O4 molecule are Nitrogen and Carbon. The Lewis structures for N2O4 and BF3 show a double-bonded N-N in the center, doubly-bonded to 2 O atoms on each end and single-bonded B to three F atoms respectively. These structures reveal the shape of N2O4 to be bent with approximately 134 degrees bond angle, and BF3 to be a trigonal planar with 120 degrees bond angle.

Step-by-step solution

Identify the isoelectronic species

Given that we know these species to be isoelectronic with Carbon (C) and Nitrogen (N), we need to identify the species themselves. Given the formula X2O4, the isoelectronic species with Carbon (C) would be Nitrogen (N), and the species isoelectronic with Nitrogen (N) would be Carbon (C) due to the number of the valence electrons.

Draw the Lewis structures

Now that we have identified the isoelectronic species, which are N for Carbon (C) and C for Nitrogen (N), we can draw the Lewis structures. For the X2O4 molecule, it is a molecule of N2O4. Its Lewis structure consists of two N atoms double-bonded, or X-X bonded, in the center with two O atoms each doubly bonded to the N atoms on both ends. The corresponding Boron fluoride would be BF3. Its Lewis structure consists of one Boron (B) atom single-bonded to three Fluorine (F) atoms arranged around the B atom.

Describe the shapes

Based on the drawn Lewis structures, the N2O4 molecule shape can be described as a bent molecule with a bond angle of approximately 134 degrees. On the other hand, the BF3 molecule has a trigonal planar shape with bond angles of 120 degrees.

Key concepts

Isoelectronic Species

Isoelectronic species are molecules or ions that have the same number of electrons or the same electronic structure. This concept is crucial for understanding how atoms combine to form compounds with similar arrangements of electrons. In our exercise, we discussed isoelectronic species for carbon (C) and nitrogen (N) which share the formula \(\mathrm{X}_{2} \mathrm{O}_{4}\).

Understanding the term isoelectronic helps in predicting and visualizing the behavior of different elements when they bond. Thinking of atoms or ions that are isoelectronic is much like considering different families with the same number of members; they might look different on the outside (different sizes, or different locations on the periodic table), but internally, their 'family structure' or electron configuration is remarkably alike. This similarity accounts for many of their shared chemical properties.

The Lewis structures in the exercise use the 'family resemblance' of these isoelectronic species to demonstrate how they bond. With nitrogen (N) taking the place of carbon (C), and vice versa, we observe their ability to perform similar bonding roles despite being distinct elements.

Molecular Geometry

The molecular geometry, or the shape of a molecule, plays a vital role in determining the molecule's properties and reactivity. It is defined by the positions of the nuclei and the electron pairs around the central atom. To predict the shape, it's essential to understand the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron pairs around a central atom will configure themselves to minimize repulsion.

In the exercise, we examined two distinct molecules: N2O4 and BF3. For N2O4, we note that this molecule is bent due to the repulsion between the lone pairs and bonding pairs of electrons, resulting in a specific bond angle. Conversely, for the molecule BF3, the absence of lone pairs on the central boron atom leads to a trigonal planar structure with equal bond angles of 120 degrees.

Understanding molecule geometry not only enables us to visualize the shape of the molecule but also allows us to predict how molecules will interact with one another. It has far-reaching implications, affecting the boiling point, melting point, and solubility of the compounds.

Valence Electrons

Valence electrons are the electrons that reside in the outermost electron shell of an atom. These are the electrons involved in chemical bonding, and they play a critical role in determining how an element reacts chemically. The number of valence electrons can often be deduced from an element's group number on the periodic table, which makes it easier to understand and predict bonding patterns.

For instance, in the step-by-step solution provided, the valence electrons guide the formation of the Lewis structures. Both carbon (C) and nitrogen (N) have valence electrons that enable the formation of the \(\mathrm{X}_{2} \mathrm{O}_{4}\) molecule. It's also these electrons that determine the molecular geometry as discussed previously.

When drawing the Lewis structures, every bond or line represents a pair of valence electrons that are being shared between atoms, while dots represent unshared valence electrons. It is the arrangement of these valence electrons that lead to the formation of specific shapes of molecules, like the bent shape of N2O4 or the trigonal planar shape of BF3.