Question

Write Lewis structures that obey the octet rule for each of the following molecules and ions. (In each case the first atom listed is the central atom.) a. \(\mathrm{POCl}_{3}, \mathrm{SO}_{4}^{2-}, \mathrm{XeO}_{4}, \mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{4}^{-}\) b. \(\mathrm{NF}_{3}, \mathrm{SO}_{3}^{2-}, \mathrm{PO}_{3}^{3-}, \mathrm{ClO}_{3}^{-}\) c. \(\mathrm{ClO}_{2}^{-}, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\) d. Considering your answers to parts a, b, and c, what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons?

Short answer

a. The Lewis structures for the species are as follows: $\mathrm{POCl}_{3}$: ``` Cl | O = P - Cl | Cl ``` $\mathrm{SO}_{4}^{2-}$: ``` O | O = S = O | O ``` $\mathrm{XeO}_{4}$: ``` O | O - Xe - O | O ``` $\mathrm{PO}_{4}^{3-}$: ``` O - | O - P - O 2- | O - ``` $\mathrm{ClO}_{4}^{-}$: ``` O - | O - Cl - O - | O - ``` b, c. Follow the same steps outlined in the given solution to find the Lewis structures for the remaining species. d. Species with the same number of atoms and valence electrons tend to have similar coordination numbers or bond multiplicities. However, charged atoms can change bond and lone pair arrangements.

Step-by-step solution

Count the valence electrons

For each species: - \(\mathrm{POCl}_{3}\): P(5) + 3 x Cl(7) = 26 valence electrons - \(\mathrm{SO}_{4}^{2-}\): S(6) + 4 x O(6) + 2(extra electrons) = 32 valence electrons - \(\mathrm{XeO}_{4}\): Xe(8) + 4 x O(6) = 32 valence electrons - \(\mathrm{PO}_{4}^{3-}\): P(5) + 4 x O(6) + 3(extra electrons) = 29 valence electrons - \(\mathrm{ClO}_{4}^{-}\): Cl(7) + 4 x O(6) + 1(extra electron) = 32 valence electrons

Place the central atom and draw single bonds to outer atoms

Here are the structures with single bonds: \(\mathrm{P-O-Cl, S-O, Xe-O, P-O, Cl-O}\)

Complete outer atom octets

Add lone pairs to the outer atoms to complete their octets: \(\mathrm{POCl}_{3}:\) ``` Cl | O = P - Cl | Cl ``` \(\mathrm{SO}_{4}^{2-}:\) ``` O | O = S = O | O ``` \(\mathrm{XeO}_{4}:\) ``` O | O - Xe - O | O ``` \(\mathrm{PO}_{4}^{3-}:\) ``` O | O - P - O | O ``` \(\mathrm{ClO}_{4}^{-}:\) ``` O | O - Cl - O | O ```

Form multiple bonds if necessary or add charges

For \({\mathrm{PO}_{4}}^{3-}\) and \({\mathrm{ClO}_{4}}^{-}\), add charges to the oxygen atoms: \(\mathrm{PO}_{4}^{3-}:\) ``` O - | O - P - O 2- | O - ``` \(\mathrm{ClO}_{4}^{-}:\) ``` O - | O - Cl - O - | O - ``` b. $\mathrm{NF}_{3}, \mathrm{SO}_{3}^{2-}, \mathrm{PO}_{3}^{3-}, \mathrm{ClO}_{3}^{-}$ You can repeat the process above for these species. c. \(\mathrm{ClO}_{2}^{-}, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\) Again, follow the same process for these species. d. Analyzing the patterns in the structures of these species: Species with the same number of atoms and valence electrons tend to have similar coordination numbers (number of bonded atoms for the central atom) or similar bond multiplicities (single, double, or triple bonds). However, the presence of charges on the atoms can change the arrangement of the bonds and lone pairs in the species.

Key concepts

Octet Rule

The octet rule is a chemical guideline that reflects the tendency of atoms to prefer having eight electrons in their valence shell, mirroring the electron configuration of noble gases which are known for their stability. When creating Lewis structures, this rule helps us depict how atoms bond and ensure that most atoms (especially those in the second row of the periodic table) end up with an octet. However, there are exceptions. For instance, some elements can have fewer or more than eight valence electrons; an example includes elements in expanded octet scenarios like phosphorus in \(\mathrm{POCl}_{3}\) or sulfur in \(\mathrm{SO}_{4}^{2-}\).

When dealing with ions such as \(\mathrm{PO}_{4}^{3-}\) and \(\mathrm{ClO}_{4}^{-}\), additional electrons that contribute to the overall negative charge also need to be considered to satisfy the octet rule. Through this, the octet rule serves as a foundation for predicting the stable arrangement of electrons in molecules and ions.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. They are critical in chemical bonding, as they are the ones involved in forming bonds between atoms. For example, phosphorus has five valence electrons and chlorine has seven. In \(\mathrm{POCl}_{3}\), we tally the valence electrons from each atom to determine how they might bond to satisfy their octet. Similarly, additional electrons for negatively charged ions are added, as seen in \(\mathrm{SO}_{4}^{2-}\) which has two extra valence electrons due to its charge.

Understanding the count of valence electrons is the first step in drawing Lewis structures, as observed in the textbook solution. The total count guides us in arranging atoms around the central one, forming initial bonds, and distributing remaining electrons to fulfill the atoms' need for a complete valence shell.

Molecular Geometry

Molecular geometry, informed by the number of bonds and lone pairs on an atom, defines the three-dimensional arrangement of atoms within a molecule. The geometry dictates the shape and sometimes the reactivity and properties of the substance. VSEPR theory, which stands for Valence Shell Electron Pair Repulsion, is a model used to predict geometry by assuming that electron pairs will arrange themselves to minimize repulsion.

For instance, \(\mathrm{POCl}_{3}\) adopts a trigonal pyramidal shape due to the lone pair on the phosphorus, whereas \(\mathrm{SO}_{4}^{2-}\) is tetrahedral as sulfur is bonded to four oxygen atoms with no lone pairs. It's fascinating how species like \(\mathrm{XeO}_{4}\) show that even noble gas atoms can partake in bonding, resulting in distinctive geometries contrary to their inert reputation.

Chemical Bonding

Chemical bonding is the force that holds atoms together in molecules. There are three primary types: ionic, covalent, and metallic. In the context of Lewis structures for molecules and polyatomic ions, we primarily deal with covalent bonds where atoms share pairs of electrons. But there's more to it: atoms may form single, double, or triple bonds depending on the need to satisfy the octet rule.

In our examples, the molecules \(\mathrm{PO}_{4}^{3-}\) and \(\mathrm{ClO}_{4}^{-}\) involve double bonds to ensure each oxygen atom achieves an octet. Furthermore, species containing the same number of atoms and valence electrons may exhibit similar bonding patterns, which aids in understanding and predicting the structure of unknown molecules.