Question

Two different compounds have the formula \(\mathrm{XeF}_{2} \mathrm{Cl}_{2}\) . Write Lewis structures for these two compounds, and describe how measurement of dipole moments might be used to distinguish between them.

Short answer

The Lewis structures for the two compounds with the formula XeF2Cl2 are: Compound 1: :F: | :Cl-Xe-Cl: | :F: Compound 2: :F: | :Cl:Xe:Cl: | :F: Measurement of dipole moments can be used to distinguish between these two compounds. Compound 1 has a net dipole moment of zero, making it a nonpolar molecule, whereas Compound 2 has a non-zero net dipole moment, making it a polar molecule. Techniques such as infrared or Raman spectroscopy can be used to differentiate between these compounds based on their interactions with electromagnetic radiation.

Step-by-step solution

Determine the Possible Lewis Structure for Compound 1

Firstly, we will create a possible Lewis structure of the first compound. Xe has a total of 8 valence electrons, F has 7, and Cl has 7. We can start by putting the Xe atom in the center and arranging the F and Cl atoms around it. In compound 1, let's consider that the two F atoms are positioned opposite to each other and the Cl atoms are also on opposite sides. The structure would look like: ``` F | Xe - Cl | F ``` Now, we will add lone pairs. Each F atom will have 3 lone pairs, and each Cl atom will have 2 lone pairs. The central Xe atom has no lone pairs left because it has 8 electrons and each bond contributes 2 electrons. The Lewis structure for compound 1 is: ``` :F: | :Cl-Xe-Cl: | :F: ```

Determine the Possible Lewis Structure for Compound 2

For the second compound, let's consider that the F and Cl atoms are positioned adjacently. The structure would look like: ``` F | Cl - Xe - Cl | F ``` Adding the lone pairs like before, we have: ``` :F: | :Cl:Xe:Cl: | :F: ```

Understand the Concept of Dipole Moments

A dipole moment is a measure of the charge distribution in a molecule. It is described as the product of the charge and the distance between the atoms. The dipole moment (\(\mu\)) can be calculated using the formula: \[\mu = Q \times d\] Where: - \(Q\) denotes the charge difference between the bonded atoms. - \(d\) denotes the distance between the bonded atoms. Dipole moments are vector quantities. In a molecule, the overall dipole moment is determined by the vector sum of the individual bond dipole moments. A molecule with a net dipole moment is polar, whereas a molecule with no net dipole moment (the bond dipoles cancel each other out) is nonpolar.

Determine Dipole Moments for both Compounds

In compound 1, the dipole moments of the opposite Xe-F and Xe-Cl bonds will cancel each other out, leading to an overall net dipole moment of zero. This compound is nonpolar. In compound 2, the bond dipole moments of the Xe-F and Xe-Cl bonds do not cancel each other out, leading to an overall non-zero net dipole moment. This compound is polar.

Explain how Measurement of Dipole Moments can be used to Distinguish Between Them

By measuring the dipole moments of the two compounds, we can distinguish between them. Since compound 1 has a net dipole moment of zero, it is nonpolar. On the other hand, compound 2 has a non-zero net dipole moment, making it a polar molecule. Using methods such as infrared or Raman spectroscopy, polar and nonpolar molecules can be distinguished based on their interactions with electromagnetic radiation. Therefore, by measuring the dipole moments and observing the behavior of the molecules in these techniques, it is possible to differentiate between compound 1 and compound 2.

Key concepts

Dipole Moments

Dipole moments are essential for understanding the distribution of electric charges in a molecule. They arise when there is a separation of positive and negative charges, which typically occurs in polar covalent bonds. The dipole moment is a vector quantity consisting of both magnitude and direction.
It can be calculated using the formula: \[\mu = Q \times d\]Here, \(\mu\) represents the dipole moment, \(Q\) is the difference in charge between atoms, and \(d\) is the distance between the bonded atoms.
In simpler terms, imagine a tug of war where atoms pull on a shared pair of electrons. The stronger the pull (difference in electronegativity), the larger the dipole moment.
For example, in a molecule like water (H₂O), there is a significant dipole moment because the oxygen atom is much more electronegative than the hydrogen atoms, creating a particular orientation of charges with distinct polar regions.
Measuring and analyzing dipole moments can provide insights into the molecular structure and help determine if a molecule is polar or nonpolar.

Molecular Polarity

Understanding molecular polarity involves looking at the shape and charge distribution of a molecule. It directly depends on the arrangement of bonds and the symmetry of the molecule.
A molecule is polar if it has an uneven distribution of electron density, resulting in distinct positive and negative ends, or poles. Nonpolar molecules, on the other hand, generally have their dipoles cancel out due to a symmetric arrangement.
When considering molecular polarity, it's important to consider both the shape and the constituent atoms' electronegativity. Even if bond dipoles are present due to electronegativity differences, the overall molecule can be nonpolar if the molecular geometry allows these dipoles to cancel each other.
For instance, carbon dioxide (CO₂) is a linear molecule where oxygen atoms on each side pull equally in opposite directions, leading to cancellation of the dipoles and a nonpolar molecule overall. In contrast, ammonia (NH₃) is polar because its pyramidal shape leads to an unequal distribution of charge.
Being aware of whether a compound is polar or nonpolar helps predict its interactions in a chemical system and its solubility in various solvents.

XeF2Cl2 Compounds

Compounds with the formula \(\mathrm{XeF_2Cl_2}\) offer a great example to study molecular polarity and dipole moments. Xenon (Xe), despite being a noble gas atom, can form compounds with fluorine (F) and chlorine (Cl) due to the availability of vacant p and d orbitals. This gives rise to diverse molecular geometries that influence molecular properties.
In the first possible compound structure, when \(\mathrm{F}\) and \(\mathrm{Cl}\) atoms are arranged opposite each other around \(\mathrm{Xe}\), the molecular shape allows for the cancellation of dipoles due to the symmetrical arrangement. As a result, the molecule has an overall net dipole moment of zero, making it nonpolar.
In the second structure, if \(\mathrm{F}\) and \(\mathrm{Cl}\) atoms are positioned adjacently, the asymmetrical shape leads to a net dipole moment as the dipoles do not completely cancel out. This makes the molecule polar.
These two structural isomers can actually be distinguished by measuring their dipole moments using techniques like spectroscopy. A nonpolar compound would not interact with an electric field in the same way as a polar one would. Learning about compounds like \(\mathrm{XeF_2Cl_2}\) provides an in-depth understanding of how atomic arrangement and molecular geometry affect chemical behavior.