Which of the following best represents the Lewis dot structure for the phosphate ion, PO4^{3-}?

• A. A central P atom double-bonded to one O atom and single-bonded to three O atoms, with two O atoms carrying a single negative charge and all atoms having complete octets.
• B. A central P atom double-bonded to two O atoms and single-bonded to two O atoms, with one O atom carrying a single negative charge.
• C. A central P atom surrounded by four O atoms, each connected by a single bond, with three O atoms carrying a single negative charge and all atoms having complete octets.
• D. A central P atom single-bonded to four O atoms, with no formal charges shown and incomplete octets on the O atoms.

Answer: (C)

Verified step by step guidance

1. Determine the total number of valence electrons available for the phosphate ion, PO4^{3-}. Phosphorus (P) has 5 valence electrons, each oxygen (O) has 6 valence electrons, and the ion has an extra 3 electrons due to the 3- charge. Calculate the sum: \$5 + 4 \times 6 + 3$.
2. Draw a skeletal structure with phosphorus as the central atom bonded to four oxygen atoms. Connect each oxygen to phosphorus with a single bond initially.
3. Distribute the remaining valence electrons to complete the octets of the oxygen atoms first, placing lone pairs around each oxygen to satisfy the octet rule.
4. Check the formal charges on each atom by using the formula: \(\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2} \times \text{Bonding electrons})\). Adjust bonding (e.g., add double bonds) if necessary to minimize formal charges and achieve the most stable structure.
5. Confirm that the final Lewis structure has phosphorus bonded to four oxygen atoms with single bonds, three of the oxygen atoms carrying a single negative charge, and all atoms having complete octets, which matches the correct representation of the phosphate ion.