In the attached Lewis structure of BrO3, how are every atom, bond, and lone pair positioned to complete the structure and what are the appropriate formal charges for each atom?

The Lewis structure of bromate (BrO₃^–) shows one bromine atom bonded to three oxygen atoms. In this structure, proper positioning of atoms, bonds, and lone pairs is crucial to reflect the molecule’s actual charge distribution.

To identify the formal charges, we need to follow a few steps:

1. Determine the number of valence electrons for each atom: Bromine (Br) has 7 valence electrons, and each Oxygen (O) has 6.
2. Count the number of bonds and lone pairs. In the BrO₃ structure, bromine is centrally located, forming one double bond with one oxygen and two single bonds with the other two oxygens. The oxygen atoms that are bound by single bonds also carry a negative formal charge, contributing to the overall -1 charge of the molecule.
3. Use the formal charge formula:
   Formal Charge = Valence Electrons – (Non-bonding Electrons + 0.5 * Bonding Electrons)

Calculating formal charges:

• Bromine (Br):
  • Valence Electrons = 7
  • Non-bonding Electrons = 0
  • Bonding Electrons = 8 (4 single bonds count as 1 each, and the double bond counts as 2)
  • Formal Charge = 7 – (0 + 0.5 * 8) = 7 – 4 = +3
• Double-bonded Oxygen (O):
  • Valence Electrons = 6
  • Non-bonding Electrons = 4 (two pairs)
  • Bonding Electrons = 4 (double bond)
  • Formal Charge = 6 – (4 + 0.5 * 4) = 6 – 6 = 0
• Single-bonded Oxygens (2x O):
  • Valence Electrons = 6
  • Non-bonding Electrons = 6 (three pairs)
  • Bonding Electrons = 2 (single bond)
  • Formal Charge = 6 – (6 + 0.5 * 2) = 6 – 7 = -1

In summary, the formal charges for BrO₃ are as follows: Br has a formal charge of +1, the double-bonded O has a formal charge of 0, and each single-bonded O has a formal charge of -1. This configuration allows each atom to satisfy its bonding preferences while also achieving an overall charge of -1 for the bromate ion.