Question

(a) Draw the Lewis structures for at least four species that have the general formula $$[: \mathrm{X} \equiv \mathrm{Y}:]^{n}$$ where \(X\) and Y may be the same or different, and \(n\) may have a value from +1 to \(-2 .\) (b) Which of the compounds is likely to be the strongest Bronsted base? Explain. [Sections 22.1, 22.7, and 22.9]

Short answer

Based on the analysis of electronegativity and charge, the compound $$[\mathrm{:N} \equiv \mathrm{N}:]^{-2}$$ is likely to be the strongest Bronsted base. This is because nitrogen is less electronegative than oxygen and the double negative charge makes it more inclined to donate its lone pair of electrons. The other species either have more electronegative atoms involved or more stable bonds that prevent them from donating electrons as easily.

Step-by-step solution

Draw the Lewis structures for four species

Choose four different species and draw their Lewis structures. We have selected the following four species: 1. $$[:\mathrm{N} \equiv \mathrm{N}:]^{-2}$$ 2. $$[:\mathrm{C} \equiv \mathrm{N}:]^{-}$$ 3. $$[:\mathrm{O} \equiv \mathrm{O}:]^{+}$$ 4. $$[:\mathrm{N} \equiv \mathrm{O}:]^{0}$$ Now, we will draw the Lewis structures for these species: 1. $$[\mathrm{:N} \equiv \mathrm{N}:]^{-2}$$: There are two nitrogen atoms, each with five valence electrons, and one extra electron for the double negative charge, so a total of 11 valence electrons. 2. $$[\mathrm{:C} \equiv \mathrm{N}:]^{-}$$: There is one carbon atom with four valence electrons and one nitrogen atom with five valence electrons, and one extra electron for the negative charge, so a total of 10 valence electrons. 3. $$[\mathrm{:O} \equiv \mathrm{O}:]^{+}$$: There are two oxygen atoms, each with six valence electrons, and one less electron for the positive charge, so a total of 11 valence electrons. 4. $$[\mathrm{:N} \equiv \mathrm{O}:]^{0}$$: There is one nitrogen atom with five valence electrons and one oxygen atom with six valence electrons, so a total of 11 valence electrons.

Determine the strongest Bronsted base

Now, we need to determine which of these compounds is likely to be the strongest Bronsted base. To do this, we will consider the electronegativity of the atoms involved and the charge on the species, as these properties are related to the basicity of the species. 1. $$[\mathrm{:N} \equiv \mathrm{N}:]^{-2}$$: Nitrogen is less electronegative than oxygen and has a double negative charge, which suggests it may be a strong Bronsted base. Nitrogen is more likely to donate its lone pair of electrons. 2. $$[\mathrm{:C} \equiv \mathrm{N}:]^{-}$$: Carbon is less electronegative than nitrogen and the species has a negative charge, so this species may also be a strong Bronsted base. However, the carbon-nitrogen triple bond is very stable, which may affect its willingness to donate electrons. 3. $$[\mathrm{:O} \equiv \mathrm{O}:]^{+}$$: Oxygen is more electronegative and the species has a positive charge, so this species would be a weak Bronsted base. It is less likely to donate its lone pair of electrons. 4. $$[\mathrm{:N} \equiv \mathrm{O}:]^{0}$$: The electronegativities of nitrogen and oxygen are close, but oxygen is still more electronegative. This species has no charge, so it has less inclination to donate or accept electrons.

Conclusion

Based on the analysis of electronegativity and charge, the compound $$[\mathrm{:N} \equiv \mathrm{N}:]^{-2}$$ is likely to be the strongest Bronsted base. This is because nitrogen is less electronegative than oxygen and the double negative charge makes it more inclined to donate its lone pair of electrons. The other species either have more electronegative atoms involved or more stable bonds that prevent them from donating electrons as easily.

Key concepts

Bronsted Bases

A Bronsted base is a substance that is capable of accepting a proton ( ext{H}^{+} ext{ion}). Understanding Bronsted bases is essential within the context of Lewis structures because their ability to accept protons depends on the availability of lone pairs of electrons.

• To determine the strength of a Bronsted base, we often consider factors such as the electronegativity of the atoms involved and the charge on the ion.
• Less electronegative atoms tend to hold onto electrons less tightly, making them better at sharing electrons with protons and thus stronger bases.
• Negatively charged species tend to be better Bronsted bases, as they have extra electrons that they can donate or share easily.

In our example, the species $$[: ext{N} ext{≡} ext{N}:]^{-2}$$, is the strongest Bronsted base. The two nitrogen atoms provide lone pairs, and the double negative charge enhances its proton-accepting ability.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom and are critical in determining how atoms bond with one another. In Lewis structures, valence electrons are represented as dots.

• In the species $$[: ext{N} ext{≡} ext{N}:]^{-2}$$, each nitrogen atom has five valence electrons, plus an additional two electrons for the $-2$ charge, giving a total of 12 valence electrons to be shared.
• Similarly, carbon contributes four valence electrons, while oxygen contributes six, impacting the structure and stability of compounds.
• The number of valence electrons helps us predict the bonding capacity of an atom; for example, nitrogen typically forms three bonds, while oxygen can generally form two.

Understanding and counting valence electrons allows chemists to accurately draw Lewis structures and predict molecular geometry.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons. In the context of these chemical species, it plays a major role in determining the compound's behavior as a Bronsted base.

• Oxygen is more electronegative than nitrogen, hence it attracts electrons more strongly, which affects the willingness of species containing oxygen to donate electron pairs.
• For instance, in $$[: ext{O} ext{≡} ext{O}:]^{+}$$, the high electronegativity of oxygen reduces its potential to act as a Bronsted base, despite its lone electron pairs.
• Conversely, nitrogen's moderate electronegativity makes it more capable of sharing electrons compared to oxygen, contributing to the strength of $$[: ext{N} ext{≡} ext{N}:]^{-2}$$ as a base.

Thus, electronegativity helps in predicting how atoms will behave in a compound, especially in proton donation or acceptance.

Triple Bonds

Triple bonds are a type of chemical bond where three pairs of electrons are shared between two atoms, significantly contributing to the stability and strength of a molecule.

• In Lewis structures, a triple bond is depicted by a line comprising three pairs of dots or a solid line connecting the atoms, such as in $$[: ext{C} ext{≡} ext{N}:]^{-}$$ and $$[: ext{N} ext{≡} ext{N}:]^{-2}$$.
• Triple bonds are shorter and stronger than single or double bonds, which affects the chemical reactivity of the species.
• The stability imparted by triple bonds can sometimes reduce the ability of an atom to donate electron pairs, influencing its basicity, as seen in the lower basicity of $$[: ext{C} ext{≡} ext{N}:]^{-}$$ despite its negative charge and lone pairs.

Knowing about triple bonds helps us understand molecular interactions and predict physical properties like bond length and bond strength.