(a) Draw the Lewis structures for at least four species that have the general formula $$ [: X=Y:]^{n} $$ where \(\mathrm{X}\) and \(\mathrm{Y}\) 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\(]\)

The concept of a Bronsted base is central to understanding many chemical reactions, particularly acid-base reactions. According to Bronsted's theory, a base is any substance that can accept a proton (a hydrogen ion, H⁺). This is an important concept to grasp because it helps explain how chemical species interact with each other in a solution.

Analyzing the exercise, we know that the strongest Bronsted base is the one that is most likely to accept a proton. Factors influencing a molecule's ability to act as a base include its charge and the presence of electronegative atoms, such as oxygen, which are keen on attracting protons due to their partial negative charge. Among the listed species, NO^- stands out as the strongest Bronsted base because it has a negative charge, making it more likely to attract and accept a hydrogen ion.

Chemical bonding is the force that holds atoms together in molecules and compounds. At the heart of chemical bonding is the concept of sharing or transferring electrons. There are several types of bonds, but the one we are concerned with here is the covalent bond, where atoms share pairs of valence electrons.

In Lewis structures, which reveal the nature of chemical bonding in molecules, a single line represents a shared pair of electrons, or a single bond. In the molecules from the exercise, like CO or NN, double lines between atoms indicate double bonds, where two pairs of electrons are shared. Determining how these electrons are shared is crucial for predicting molecular properties, reactivity, and how a molecule will interact with others, such as how it behaves as a Bronsted base.

The concept of valence electrons is fundamental in understanding chemical bonding and reactivity. These are the electrons that reside in the outermost shell of an atom and participate in the formation of chemical bonds. The number of valence electrons determines an atom's bonding behavior.

For example, in the step-by-step solution provided, carbon in CO has four valence electrons, which it uses to form multiple bonds with oxygen. Similarly, the negative charge in NO^- indicates an extra valence electron that influences its ability to act as a Bronsted base. Recognizing the role valence electrons play is essential for drawing accurate Lewis structures and predicting the properties of a molecule or ion.