Problem 135 - Chemistry The Science in Context

Question

Which of the following unstable nitrogen oxides, \(\mathrm{N}_{2} \mathrm{O}_{2}\) \(\mathrm{N}_{2} \mathrm{O}_{5},\) and \(\mathrm{N}_{2} \mathrm{O}_{3},\) are polar molecules? \(\left(\mathrm{N}_{2} \mathrm{O}_{2} \text { and } \mathrm{N}_{2} \mathrm{O}_{3}\right.\) have \(\left.\mathrm{N}-\mathrm{N} \text { bonds; } \mathrm{N}_{2} \mathrm{O}_{5} \text { does not. }\right)\)

Step-by-Step Solution

Answer

a. N2O2 and N2O3 b. N2O2 and N2O5 c. N2O3 and N2O5 d. All of them Answer: a. N2O2 and N2O3

Step 1: Determining the molecular shape of each molecule

For N2O2, it has a total of 6 valence electrons (5 from each nitrogen and 6 from each oxygen). It has a bent molecular shape with two N-O bonds coming off each nitrogen. For N2O5, it has a total of 22 valence electrons. It has a linear molecular shape with two N-O bonds and a central N-N bond. It also has a O=N-O bond on both sides. For N2O3, it has a total of 14 valence electrons. It has a linear molecular shape with one N-O bond, a central N-N bond and one N=O bond.

Step 2: Determining the electronegativity difference in the atoms

The electronegativity difference between nitrogen and oxygen atoms (∆EN = EN(O) - EN(N)) is around 0.5. This difference is small but enough to form polar covalent bonds in these molecules.

Step 3: Determining the polarity of the molecule

For N2O2, the bent molecular shape with two N-O polar bonds coming off each nitrogen results in an uneven distribution of electrons and makes the molecule polar. For N2O5, the linear molecular shape cancels out the individual bond dipole moments and creates a net dipole moment of zero. The molecule has no overall polarity and therefore is non-polar. For N2O3, the linear molecular shape cancels out the N-O covalent bond's dipole moments, but the N=O bond has an uneven distribution of electrons, making the molecule polar.

Step 4: Conclusion

Out of the given unstable nitrogen oxides (N2O2, N2O5, and N2O3), N2O2 and N2O3 are polar molecules.

Key Concepts

Molecular Shape

The molecular shape of a compound is pivotal in determining its polarity. It's all about how atoms are arranged and how the bonds form angles with each other. For example, in the molecule \(_2o_2\), the molecular shape is bent. This means that the atoms are not in a straight line. Instead, they form an angle, causing an uneven distribution of electrical charge across the molecule. This is what leads to a polar molecule where one end has a slight positive charge and the other a slight negative charge.

Contrast this with \(_2o_5\). This molecule is described as linear, meaning the atoms lie in a straight line. Because of this shape, any pull from individual dipole moments cancel each other out, resulting in a nonpolar molecule.

Understanding the concept of molecular shape helps predict whether a molecule is polar or nonpolar. The shapes can be linear, bent, trigonal planar, tetrahedral, among others. Each plays a significant role in the overall electron distribution and, consequently, the molecule's polarity.

Electronegativity

Electronegativity is how strongly an atom attracts electrons within a bond. Different elements have different electronegativities, which affects how electrons are shared in bonds. In the nitrogen oxides we're examining, nitrogen and oxygen are the atoms to focus on.

• Nitrogen has a lower electronegativity compared to oxygen.
• In a bond between nitrogen and oxygen, electrons will spend more time closer to the oxygen atom due to its higher electronegativity.

This difference in electronegativity (around 0.5) leads to polar covalent bonds. Even though this difference is not very large, it's enough to create a noticeable dipole.When considering molecules such as \(_2o_2\) and \(_2o_3\), where the molecule shape does not fully cancel out these differences, the compounds turn out to be polar. Meanwhile, in \(_2o_5\), the linear shape helps in cancelling these polar effects.

Dipole Moment

The dipole moment is a measure of the polarity of a molecule. It arises due to differences in electronegativity between bonded atoms and the resulting bond dipoles.In a molecule like \(_2o_2\), with a bent shape, the dipole moments do not cancel out, leading to a net dipole moment making it polar. This means there is a region of the molecule with a slightly negative charge and another with a slightly positive charge.

For \(_2o_5\), despite having polar bonds, the linear shape causes the dipole moments to cancel out, resulting in no net dipole moment, or a nonpolar molecule.Lastly, \(_2o_3\) with its somewhat skewed linear configuration, possesses areas where the dipole moments do not cancel out due to an especially polar N=O bond. This leads to the molecule being polar.

• A net dipole moment signifies an overall uneven distribution of electrons.
• No net dipole moment means the molecule has even electron distribution.

Understanding dipole moments thus is essential in determining the polarity and, ultimately, the behavior of molecules in chemical reactions.