(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]

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.

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.

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.