Question

Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule, and explain why they do not: (a) \(\mathrm{SO}_{3}^{2-}\), (b) \(\mathrm{AlH}_{3}\) (c) \(\mathrm{N}_{3}^{-}\) (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (e) \(\mathrm{SbF}_{5}\)

Short answer

The Lewis structures for the given ions or molecules are: (a) SO3^2-: \[\begin{array}{ c c c c } & \text{O} & \\ \text{O} & - & \text{S} & =\text{O} \\ \end{array}\] (does not obey octet rule, S has 10 electrons) (b) AlH3: \[\begin{array}{ c c c c } & \text{H} & \\ \text{H} & - & \text{Al} & -\text{H} \\ \end{array}\] (does not obey octet rule, Al has 6 electrons) (c) N3^-: \[\begin{array}{ c c c c } & & \text{N} & \\ & 3 \text{N} & 2 & \\ & & \text{N} & \\ \end{array}\] (obeys octet rule) (d) CH2Cl2: \[\begin{array}{ c c c } \text{H} & - & \text{C} \\ & | & \\ \text{H} & - & \text{Cl} \\ \end{array}\] (obeys octet rule) (e) SbF5: \[\begin{array}{ c c c } & & \text{F} \\ & 1 \text{Sb} & \\ 5 & | & \\ & & \text{F} \\ \text{F} & - & \text{F} \\ & & \text{F} \\ \end{array}\] (does not obey octet rule, Sb has 10 electrons)

Step-by-step solution

(a) SO3^2-

1. Add the valence electrons for each atom: S has 6, each O has 6, and we have two additional electrons due to the -2 charge. 2. Connect the atoms with single bonds. S is the central atom, and O atoms surround it. 3. Distribute the remaining electrons to complete the octet for each atom. 4. If necessary, add double or triple bonds to satisfy the octet rule. Lewis structure: \[\begin{array}{ c c c c } & \text{O} & \\ \text{O} & - & \text{S} & =\text{O} \\ \end{array}\] The molecule does not obey the octet rule and has an additional lone pair, making S have 10 electrons instead of 8.

(b) AlH3

1. Add the valence electrons for each atom: Al has 3 and each H has 1. 2. Connect the atoms with single bonds. Al is the central atom, and H atoms surround it. Lewis structure: \[\begin{array}{ c c c c } & \text{H} & \\ \text{H} & - & \text{Al} & -\text{H} \\ \end{array}\] The molecule does not obey the octet rule since Al only has 6 electrons in its outer shell.

(c) N3^-

1. Add the valence electrons: each N has 5, and we have one additional electron contributed by the -1 charge. 2. Connect the atoms with single bonds. 3. Distribute the remaining electrons to complete the octet for each atom. 4. If necessary, add double or triple bonds to satisfy the octet rule. Lewis structure: \[\begin{array}{ c c c c } & & \text{N} & \\ & 3 \text{N} & 2 & \\ & & \text{N} & \\ \end{array}\] The molecule obeys the octet rule.

(d) CH2Cl2

1. Add the valence electrons: C has 4, each H has 1, and each Cl has 7. 2. Connect the atoms with single bonds. C is the central atom surrounded by H and Cl atoms. 3. Distribute the remaining electrons to complete the octet for each atom. Lewis structure: \[\begin{array}{ c c c } \text{H} & - & \text{C} \\ & | & \\ \text{H} & - & \text{Cl} \\ \end{array}\] The molecule obeys the octet rule.

(e) SbF5

1. Add the valence electrons: Sb has 5, and each F has 7. 2. Connect the atoms with single bonds. Sb is the central atom surrounded by F atoms. 3. Distribute the remaining electrons to complete the octet for each atom. Lewis structure: \[\begin{array}{ c c c } & & \text{F} \\ & 1 \text{Sb} & \\ 5 & | & \\ & & \text{F} \\ \text{F} & - & \text{F} \\ & & \text{F} \\ \end{array}\] The molecule does not obey the octet rule since Sb has 10 electrons in its outer shell.

Key concepts

Octet Rule

The octet rule is a fundamental guideline in chemistry that refers to the tendency of atoms to prefer having eight electrons in their valence, or outermost, electron shell. This principle reflects the observed stability of the noble gases, which naturally have complete valence electron shells. Aspiring to achieve this stable configuration, atoms will often form bonds to either share, lose, or gain electrons until they are surrounded by a total of eight valence electrons.

However, there are exceptions to this rule, particularly when it comes to elements in periods three and beyond of the periodic table. These elements can have more than eight valence electrons due to the availability of d-orbitals, which allows them to accommodate additional electrons and form what are known as expanded octets. For example, in the Lewis structure of \(\mathrm{SO}_{3}^{2-}\), sulfur has a total of ten electrons, violating the octet rule, yet it is stable due to its ability to utilize d-orbitals.

On the flip side, some elements like boron or aluminum often form stable compounds with less than eight electrons surrounding them, as seen in the molecule \(\mathrm{AlH}_{3}\). In such cases, these elements settle for fewer than eight valence electrons because they form enough stable bonds to achieve a lower energy state, despite not having a full octet.

Valence Electrons

Valence electrons are the electrons that are located in the outermost electron shell of an atom. They are critically important because they participate in the formation of chemical bonds, allowing atoms to interact, share, or transfer electrons to achieve stability. The number of valence electrons an atom has will often dictate not only the types of chemical bonds it can form but also the overall reactivity of the atom.

When drawing Lewis structures, as in the textbook exercise examples like \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) and \(\mathrm{N}_{3}^{-}\), the first step is always to count the valence electrons available for bonding. This count includes the electrons that each atom brings naturally, plus any additional charges that may influence the total electronic count. For instance, negative ions have additional valence electrons corresponding to their charge, while positive ions have fewer. In the case of \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\), carbon has four valence electrons, and each hydrogen and chlorine add their respective valence electrons to complete the molecule’s bonding framework.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of the atoms that constitute a molecule. This geometry arises from the specific ways in which the different electron pairs, both bonding and lone pairs, arrange themselves around a central atom to minimize repulsion between them, a concept known as VSEPR (Valence Shell Electron Pair Repulsion) theory.

Understanding molecular geometry is key to deciphering molecule behavior, as it influences physical properties, reactivity, and function. For instance, \(\mathrm{SbF}_{5}\), as mentioned in the textbook solution, features an expanded octet and adopts a geometry that allows the central antimony atom to bond with five fluorine atoms. This geometry, which can be trigonal bipyramidal, allows all five fluorine atoms to be equidistant from the central antimony atom, thus minimizing electron pair repulsion and stabilizing the molecule.