Question

Write Lewis structures that obey the octet rule for the following species. Assign the formal charge for each central atom. a. \(\mathrm{POCl}_{3} \quad\) c. \(\mathrm{ClO}_{4}^{-} \quad\) \(\mathrm{e} \cdot \mathrm{SO}_{2} \mathrm{Cl}_{2} \quad\) g. \(\mathrm{ClO}_{3}^{-}\) b. \(\mathrm{SO}_{4}^{2-} \quad\) d. \(\mathrm{PO}_{4}^{3-} \quad\) f. \(\mathrm{XeO}_{4} \quad\) h. \(\mathrm{NO}_{4}^{3-}\)

Short answer

The following are the Lewis structures and formal charges for each central atom: a. \(\mathrm{POCl}_{3}\): P has a formal charge of +1. b. \(\mathrm{SO}_{4}^{2-}\): S has a formal charge of 0. c. \(\mathrm{ClO}_{4}^{-}\): Cl has a formal charge of -1. d. \(\mathrm{PO}_{4}^{3-}\): P has a formal charge of -1. e. \(\mathrm{SO}_{2} \mathrm{Cl}_{2}\): S has a formal charge of 0. f. \(\mathrm{XeO}_{4}\): Xe has a formal charge of 0. g. \(\mathrm{ClO}_{3}^{-}\): Cl has a formal charge of -1. h. \(\mathrm{NO}_{4}^{3-}\): N has a formal charge of -1.

Step-by-step solution

Draw Lewis structures for each species

To draw the Lewis structures for each species, start by identifying the central atom, then arrange the other atoms around it. Next, count the total number of valence electrons for all atoms, and then place lone pair electrons around the bonded atoms to achieve an octet configuration (or where it's allowed, a double or triple bond) for each atom. a. \(\mathrm{POCl}_{3}\) The central atom is P, and the total valence electrons are 5 (for P) + 3*7 (for Cl) = 26. The Lewis structure is: O || P-Cl |\ | Cl-Cl b. \(\mathrm{SO}_{4}^{2-}\) The central atom is S, and the total valence electrons are 6 (for S) + 4*6 (for O) + 2 (charge) = 32. The Lewis structure is: O- | S=O | O- c. \(\mathrm{ClO}_{4}^{-}\) The central atom is Cl, and the total valence electrons are 7 (for Cl) + 4*6 (for O) +1 (charge) = 32. The Lewis structure is: O- | Cl=O | O- d. \(\mathrm{PO}_{4}^{3-}\) The central atom is P, and the total valence electrons are 5 (for P) + 4*6 (for O) + 3 (charge) = 32. The Lewis structure is: O- | P=O | O- e. \(\mathrm{SO}_{2} \mathrm{Cl}_{2}\) The central atom is S, and the total valence electrons are 6 (for S) + 2*6 (for O) + 2*7 (for Cl) = 32. The Lewis structure is: O Cl || | S-Cl | O f. \(\mathrm{XeO}_{4}\) The central atom is Xe, and the total valence electrons are 8 (for Xe) + 4*6 (for O) = 32. The Lewis structure is: O || Xe=O || O g. \(\mathrm{ClO}_{3}^{-}\) The central atom is Cl, and the total valence electrons are 7 (for Cl) + 3*6 (for O) + 1 (charge) = 26. The Lewis structure is: O- | Cl=O | O- h. \(\mathrm{NO}_{4}^{3-}\) The central atom is N, and the total valence electrons are 5 (for N) + 4*6 (for O) + 3 (charge) = 32. The Lewis structure is: O- | N=O | O-

Calculate the formal charge of the central atom

To calculate the formal charge of the central atom in each species, use the formula: Formal Charge = Valence Electrons - (0.5 * Bonding Electrons) - Non-bonding Electrons. a. Formal Charge for P: 5 - (0.5 * 8) - 0 = +1 b. Formal Charge for S: 6 - (0.5 * 12) - 0 = 0 c. Formal Charge for Cl: 7 - (0.5 * 12) - 0 = -1 d. Formal Charge for P: 5 - (0.5 * 12) - 0 = -1 e. Formal Charge for S: 6 - (0.5 * 12) - 2 = 0 f. Formal Charge for Xe: 8 - (0.5 * 16) - 0 = 0 g. Formal Charge for Cl: 7 - (0.5 * 10) - 0 = -1 h. Formal Charge for N: 5 - (0.5 * 12) - 0 = -1 Summary: a. P in \(\mathrm{POCl}_{3}\) has a formal charge of +1 b. S in \(\mathrm{SO}_{4}^{2-}\) has a formal charge of 0 c. Cl in \(\mathrm{ClO}_{4}^{-}\) has a formal charge of -1 d. P in \(\mathrm{PO}_{4}^{3-}\) has a formal charge of -1 e. S in \(\mathrm{SO}_{2} \mathrm{Cl}_{2}\) has a formal charge of 0 f. Xe in \(\mathrm{XeO}_{4}\) has a formal charge of 0 g. Cl in \(\mathrm{ClO}_{3}^{-}\) has a formal charge of -1 h. N in \(\mathrm{NO}_{4}^{3-}\) has a formal charge of -1

Key concepts

Octet Rule

The octet rule is a fundamental concept in chemistry that helps predict the structure of a molecule. It states that atoms tend to form bonds in such a way that they have eight electrons in their valence shell, achieving a noble gas configuration. This rule is particularly important for main group elements, like carbon, nitrogen, and oxygen.
For example, in the case of the sulfate ion, \( ext{SO}_4^{2-} \), the sulfur atom forms double bonds with two oxygen atoms to fulfill the octet rule by sharing electrons. Each oxygen atom also fulfills its octet by forming bonds and hosting any additional lone pair of electrons necessary.
However, it's important to note that the octet rule is not universal. There are exceptions, especially among transition metals and elements in higher periods that can expand their valence shell beyond eight electrons. Examples from the exercise, like \( ext{XeO}_4 \), show xenon expanding beyond the octet since it's a noble gas that can accommodate more electrons due to its available d-orbitals.

Formal Charge

Understanding formal charge is crucial to determine the most stable Lewis structure for a molecule. The formal charge of an atom in a molecule is the hypothetical charge it would have if all bonding electrons were shared equally between the bonded atoms. To calculate formal charge, use the formula:
\[ \text{Formal Charge} = \text{Valence Electrons} - (0.5 \times \text{Bonding Electrons}) - \text{Non-bonding Electrons} \]
For instance, using this formula for the phosphate ion, \( ext{PO}_4^{3-} \), helps in identifying the most plausible distribution of electrons that results in the minimal formal charge, i.e., closest to zero. For phosphorus, after calculation, it has a formal charge of -1 in \( ext{PO}_4^{3-} \). The goal is usually to minimize the formal charges across the molecule to ensure stability, as this indicates a more effective sharing of electrons across the atoms involved.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in the formation of chemical bonds. Knowing the number of valence electrons helps predict how an atom will interact with others. These electrons determine the bonding capabilities and the overall neutral charge when isolated.
To find the Lewis structure effectively, it's essential first to count the total number of valence electrons present in the molecule. For example, in \( ext{POCl}_3 \), phosphorus has 5 valence electrons, and each chlorine atom has 7. By adding these with the oxygen's 6 electrons, the total tally reaches 32 valence electrons, which are then distributed to satisfy the octet rule.
In polyatomic ions like \( ext{ClO}_4^- \), it is key to account for the charge by adding one extra electron, thus the calculation includes this additional electron to ensure total valence matches the correct electronic structure intended for the ion.