Question

The following polyatomic anions involve covalent bonds between O atoms and the central nonmetal atom. Propose an acceptable Lewis structure for each. (a) \(\mathrm{SO}_{3}^{2-} ;\) (b) \(\mathrm{NO}_{2}^{-} ;\) (c) \(\mathrm{CO}_{3}^{2-} ;\) (d) \(\mathrm{HO}_{2}^{-}\).

Short answer

The Lewis structures for \(\mathrm{SO}_{3}^{2-}\), \(\mathrm{NO}_{2}^{-}\), \(\mathrm{CO}_{3}^{2-}\), and \(\mathrm{HO}_{2}^{-}\) can be drawn by total the valence electrons, adding any extra for the charge of the ion, and then arranging them to fill octets as much as possible.

Step-by-step solution

\(SO_{3}^{2-}\) Lewis structure

Sulfur (S) has 6 valence electrons and each oxygen (O) also has 6. Since there are three O atoms, the total number of valence electrons is \(6+3*6 = 24\). Adding two more for the charge of the ion, we have 26 electrons available to fill the octets of the atoms. Place the S in the center with a single bond to each of the O atoms, using up 6 electrons, and then distribute the remaining electrons (20 of them) as lone pairs on the O atoms. There should be 8 electrons (two lone pairs and the bond) around each O, and 8 around the S (which can expand its octet as needed).

\(NO_{2}^{-}\) Lewis structure

Nitrogen (N) has 5 valence electrons and each oxygen (O) has 6, for a total of \(5+2*6 = 17\), and adding one more for the charge of the ion gives 18 electrons. With N in the center, single bond each O and leave 10 electrons. Distribute these as lone pairs on the oxygens, and put any remaining (two) on the nitrogen. Note that one of the O has a full octet, but the N and the other O have only 7 electrons, so this species is a radical.

\(CO_{3}^{2-}\) Lewis structure

Carbon (C) has 4 valence electrons and each oxygen (O) has 6, so the total is \(4+3*6 = 22\), and adding two for the charge of the ion gives 24 electrons to use. Place the carbon in the center, single bond it to all three of the oxygens, and distribute the remaining electrons (18 in total) as lone pairs on the oxygens. There should be 8 electrons around each O, and the C is short of two electrons. To solve this issue, create a double bond with one O atom, giving a full octet to each atom and making each oxygen equivalent to the others.

\(HO_{2}^{-}\) Lewis structure

Hydrogen (H) has 1 valence electron, and each oxygen (O) has 6, for a total of \(1+2*6 = 13\), and adding one for the charge of the ion gives 14 electrons to use. Place one O in center, single bond to both H and the other O, and then distribute the remaining electrons (10 in total). Each O should have 8 electrons around it (a full octet), with any remaining electron (after filling the octets of the O atoms) going on the hydrogen.

Key concepts

Covalent Bonds

In the realm of chemistry, a covalent bond is a crucial type of chemical bond that involves the sharing of electron pairs between atoms. These shared electrons allow the atoms to stick together and form a molecule. Covalent bonds are formed when atoms share electrons to achieve stability and follow the octet rule, which implies having eight electrons in their outer shell, similar to that of noble gases.
When forming a covalent bond, the participating atoms typically have similar electronegativities, meaning the tendency to attract electrons towards themselves. This similarity avoids one atom pulling the shared electrons significantly more than the other, maintaining the bond's symmetry.
Lewis structures are a visual representation of these bonds, showing how the valence electrons are distributed around the bonded atoms.

• Shared electron pairs are shown as lines between atoms, indicating covalent bonds.
• Lone pairs, or non-bonding pairs, are shown as dots around individual atoms.

Understanding covalent bonds lays the groundwork for grasping the more complex interactions seen in polyatomic ions.

Polyatomic Ions

Polyatomic ions are a group of bonded atoms that carry a net charge, meaning that these clusters of atoms act as a single unit. Despite being bonded covalently within the ion, they are still capable of forming ionic bonds with other ions in compounds.
Each exercise example, like sulfate (\(SO_3^{2-}\)), nitrate (\(NO_2^{-}\)), and carbonate (\(CO_3^{2-}\)), showcases polyatomic ions by highlighting a central atom, typically a nonmetal, surrounded by oxygen atoms. These ions can exist freely in solutions and often participate in reactions to form salts.
The charge on a polyatomic ion, like the \(2-\) on carbonate, is essential for understanding its reactivity and role in larger compounds. It reflects the total number of electrons lost or gained by the ion.

• Understanding the distribution of electrons and charges helps predict how these ions behave in chemical reactions.
• The ability to draw accurate Lewis structures for such ions is crucial for visualizing their electron arrangements and predicting bonding behavior.

Polyatomic ions are an integral concept in chemistry, linking the ideas of covalent bonding and ionic interactions.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. These electrons play a vital role in bonding and chemical reactions, as they are the ones that atoms can lose, gain, or share to form stable configurations.
For instance, in the case of sulfate \(SO_3^{2-}\), sulfur has a total of six valence electrons. Oxygen also contributes six valence electrons from each atom, forming a total of 24 valence electrons when the additional two from the charge are added.
The ability to account for valence electrons is foundational when constructing Lewis structures because it helps in determining how atoms will bond and how many electrons will be shared or paired.

• This understanding helps in observing the octet rule and explaining the stability of molecules and ions.
• It also aids in identifying lone pairs, which can influence molecular shape and reactivity.

The concept of valence electrons is key for predicting how atoms will interact and the resultant chemical compounds.

Octet Rule

The octet rule is a basic principle in chemistry that states atoms tend to bond in a way that allows them to have eight electrons in their outer shell, mimicking the stable electron configuration of noble gases. This rule is a guiding principle for drawing Lewis structures, although there are exceptions.
Most atoms adhere to the octet rule while forming covalent bonds. For instance, when drawing the Lewis structure for carbonate \(CO_3^{2-}\), each oxygen atom aims to have eight electrons around it after bonding, while carbon achieves a complete octet by sharing electrons through double bonding with one of the oxygen atoms.
The octet rule helps elucidate why certain structures are more stable than others and why molecules form specifically the way they do.

• It's crucial to check each atom in the Lewis structure to ensure it achieves a complete octet unless exceptions (such as hydrogen or radicals) apply.
• It streamlines the prediction process for molecular shapes, reactivity, and polarity.

The octet rule is a cornerstone for understanding the stability and formation of molecules, guiding the proper construction of Lewis structures.