Draw Lewis structures for the following polyatomic ions. a. \(\mathrm{OH}^{-}\) b. \(\mathrm{O}_{2}^{2-}\) c. \(\mathrm{NO}^{2-}\) d. \(\mathrm{NO}^{2+}\) e. \(\mathrm{AsO}_{4}^{3-}\)

Chemical bonding is the process that holds atoms together within molecules and compounds, forming the basis of all materials. Atoms bond in order to achieve a more stable electron configuration, which often means having a full outer electron shell, akin to that of a noble gas. This stability can be reached in several ways, giving rise to different types of bonds.

Ionic bonds form between atoms with significantly different electronegativities, where one atom donates its valence electrons to another, leading to the formation of ions. Covalent bonds, on the other hand, occur when atoms share pairs of electrons, thus filling their valence shell. A simple representation of these interactions is through Lewis structures, which depict the arrangement of electrons in a molecule or ion.

Considering the exercise provided, polyatomic ions such as \(\mathrm{OH}^{-}\), \(\mathrm{O}_{2}^{2-}\), and others are stable entities that have formed through chemical bonding, satisfying the octet rule wherever possible. The negative or positive charges indicate an excess or deficit of electrons, which impacts bonding and structure.

Valence electrons are the outermost electrons of an atom and are fundamental to the formation of chemical bonds. The number of valence electrons that an atom possesses determines how it will interact with other atoms. Atoms aim to achieve a valent shell resembling that of the noble gases, which are inherently stable due to their filled outer electron shells. This is commonly known as the 'octet rule', although there are exceptions.

In the Lewis structures exercise, valence electrons play an instrumental role. For example, oxygen (\(\mathrm{O}\)) typically has 6 valence electrons, while hydrogen (\(\mathrm{H}\)) has only 1. In forming the \(\mathrm{OH}^{-}\) polyatomic ion, the addition of the negative charge implies an extra valence electron. Similarly, arsenic (\(\mathrm{As}\)) in \(\mathrm{AsO}_{4}^{3-}\) will be surrounded by more valence electrons due to the charge of the ion. Lewis structures help visualize these valence electrons, which is crucial for understanding bonding and molecular geometry.

Polyatomic ions are charged species consisting of two or more atoms covalently bonded together, that function as a single ion. They bear either a positive (cationic) or negative (anionic) charge, not because the individual atoms are inherently charged, but because there are more or fewer electrons than protons in the combined entity. These ions are essential components in the formation of many compounds, including salts and certain molecular species.

Creating Lewis structures for polyatomic ions includes the step of adding or subtracting electrons corresponding to the ion's charge. For instance, the \(\mathrm{OH}^{-}\) ion is a simple polyatomic ion with one extra electron. In contrast, the \(\mathrm{AsO}_{4}^{3-}\) ion is more complex, featuring arsenic bonded to four oxygen atoms, and with three additional electrons added to the total count of valence electrons. In the exercise, understanding the concept of polyatomic ions is essential to accurately drawing their Lewis structures, which depict the bonding and electron arrangement.