Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in \(\mathrm{O}_{3},(\mathbf{b})\) phosphorus in \(\mathrm{PF}_{6}^{-},(\mathbf{c})\) nitrogen in \(\mathrm{NO}_{2}\), (d) iodine in ICl \(_{3}\), (e) chlorine in \(\mathrm{HClO}_{4}\) (hydrogen is bonded to \(\mathrm{O}\) ).

To find the formal charge on the central oxygen atom in O3, first draw the Lewis structure of the molecule: O=O-O Now, calculate the total number of valence, non-bonding, and bonding electrons on the central oxygen atom. Oxygen has a total of 6 valence electrons. The central oxygen atom is bonded to another oxygen atom with a double bond and one single bond. Hence, it has 4 bonding electrons (2 in the double bond). There are no non-bonding electrons on the central oxygen atom (all electrons are involved in bonds). Apply the formal charge formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 × Bonding electrons) Formal charge = (6) - (0) – (1/2 × 4) Formal charge = 6 – 2 = +1 The formal charge on the central oxygen atom in O3 is +1.

First, draw the Lewis structure of PF6-. Phosphorus (P) is surrounded by 6 fluorine (F) atoms, with each P-F bond being a single bond. Since PF6- is an ion with a negative charge, add an extra electron to the Phosphorus valence electrons. Phosphorus has 5 valence electrons + 1 extra (due to ion charge) = 6 valence electrons. There are no non-bonding electrons on phosphorus, and phosphorus has 12 bonding electrons (6 single bonds to fluorine). Apply the formal charge formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 × Bonding electrons) Formal charge = (6) - (0) - (1/2 × 12) Formal charge = 6 - 6 = 0 The formal charge on phosphorus in PF6- ion is 0.

Draw the Lewis structure of NO2. Nitrogen is bonded to two oxygen atoms, both having one double bond. One of the nitrogen's valence electrons remains unbonded. Nitrogen has 5 valence electrons, 2 non-bonding electrons (the unbonded electron), and 8 bonding electrons (2 double bonds to oxygen). Apply the formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 × Bonding electrons) Formal charge = (5) - (2) - (1/2 × 8) Formal charge = 5 – 4 = +1 The formal charge on nitrogen in NO2 is +1.

Draw the Lewis structure of ICl3. Iodine is bonded to three chlorine atoms, all with single bonds, and have two lone pairs of electrons. Iodine has 7 valence electrons, 4 non-bonding electrons (two lone pairs), and 6 bonding electrons (3 single bonds to chlorine). Apply the formal charge formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 × Bonding electrons) Formal charge = (7) - (4) - (1/2 × 6) Formal charge = 7 – 4 - 3 = 0 The formal charge on iodine in ICl3 is 0.

Draw the Lewis structure of HClO4. Chlorine is bonded to four oxygen atoms, each having a single bond, and hydrogen is bonded to one of the oxygen atoms. The chlorine atom also has one lone pair of electrons. Chlorine has 7 valence electrons, 2 non-bonding electrons (one lone pair), and 8 bonding electrons (4 single bonds to oxygen). Apply the formal charge formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 × Bonding electrons) Formal charge = (7) - (2) - (1/2 × 8) Formal charge = 7 – 2 – 4 = +1 The formal charge on chlorine in HClO4 is +1.