Some chemists believe that satisfaction of the octet rule should be the top criterion for choosing the dominant Lewis structure of a molecule or ion. Other chemists believe that achieving the best formal charges should be the top criterion. Consider the dihydrogen phosphate ion, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) , in which the \(\mathrm{H}\) atoms are bonded to \(\mathrm{O}\) atoms. (a) What is the predicted dominant Lewis structure if satisfying the octet rule is the top criterion? (b) What is the predicted dominant Lewis structure if achieving the best formal charges is the top criterion?

To draw the Lewis structure of \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\), we first need to count the total number of valence electrons. Hydrogen has 1 valence electron, phosphorus has 5, oxygen has 6, and there is an additional electron due to the negative charge. Total valence electrons = (2 Hydrogens × 1) + (1 Phosphorus × 5) + (4 Oxygens × 6) + 1 = 2 + 5 + 24 + 1 = 32

Place the least electronegative atom (phosphorus) in the center and surround it with the oxygen atoms. Then, place the hydrogen atoms bonded to two oxygen atoms. Connect them with single bonds. Now, assign the remaining valence electrons as lone pairs on the oxygen atoms.

The octet rule states that an atom in a molecule will be stable when it has 8 electrons in its outer shell (except for hydrogen, which only requires 2). In the skeleton structure, the phosphorus atom has only 4 electrons (from single bonds to each oxygen), so it requires double bonding with one oxygen to satisfy the octet rule. Additionally, hydrogen atoms are already stable with 2 electrons each. The other three oxygen atoms are also stable, having 8 electrons each in their outer shells (2 from the covalent bond and 6 from lone pairs). So, when satisfying the octet rule is the top criterion, the dominant Lewis structure involves a double bond between phosphorus and one of the oxygen atoms, and single bonds to the other oxygen atoms.

Now, let's calculate the formal charges of each atom in the structure obtained in step 3: Formal charge of Phosphorus = 5 (valence electrons) - 0.5(4 bonding electrons) - 4 (ownership of lone pair electrons) = 0 Formal charge of double-bonded Oxygen = 6 - 0.5(4) - 4 = 0 Formal charge of single-bonded Oxygen with Hydrogen = 6 - 0.5(2) - 6 = -1 Formal charge of single-bonded Oxygen without Hydrogen = 6 - 0.5(2) - 6 = -1 The total formal charge of the ion is (-1)+(-1)= -2 which is not equal to the ion charge (-1).

For achieving the best formal charges, we need to make sure that the sum of the formal charges is equal to the ion charge (-1). To do this, we can modify the structure obtained in step 3 by converting one of the single-bonded oxygen atoms without hydrogen into a double bond with phosphorus: Formal charge of Phosphorus = 5 - 0.5(6) - 2 = +1 Formal charge of one double-bonded Oxygen = 6 - 0.5(4) - 4 = 0 Formal charge of another double-bonded Oxygen = 6 - 0.5(4) - 4 = 0 Formal charge of single-bonded Oxygen with Hydrogen = 6 - 0.5(2) - 6 = -1 Formal charge of single-bonded Oxygen without Hydrogen = 6 - 0.5(2) - 6 = -1 The total formal charge is now (+1)+(-1)+(-1)= -1, which is equal to the ion charge. So, when achieving the best formal charges is the top criterion, the dominant Lewis structure involves two double bonds between phosphorus and two of the oxygen atoms, and single bonds to the other two oxygen atoms. In conclusion, the predicted dominant Lewis structure of \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) based on: (a) Satisfying the octet rule is a structure with one double bond between phosphorus and one oxygen atom, and single bonds to the other oxygen atoms. (b) Achieving the best formal charges is a structure with two double bonds between phosphorus and two oxygen atoms, and single bonds to the other two oxygen atoms.