A student proposes the following Lewis structure for the perchlorate (
) ion. [ :O: ]$^- \quad \quad || [:Cl-O-O:] \quad \quad || [ :O: ] Assign a formal charge to each atom in the student’s Lewis structure. atom \quad formal charge central O \quad 0 top O \quad 0 right O \quad 0 bottom O \quad 0 Cl \quad 0
The Correct Answer and Explanation is:
To determine the formal charges of each atom in the perchlorate ion, we use the formula:
Formal charge = valence electrons – non-bonding electrons – (bonding electrons / 2)
The perchlorate ion (ClO₄⁻) consists of a central chlorine atom bonded to four oxygen atoms. Chlorine has seven valence electrons, and oxygen has six valence electrons.
Assigning formal charges:
- Chlorine:
- Valence electrons: 7
- Non-bonding electrons: 0
- Bonding electrons: 8 (four single bonds)
- Formal charge = 7 – 0 – (8/2) = 7 – 0 – 4 = +3
- Each oxygen with a single bond:
- Valence electrons: 6
- Non-bonding electrons: 6 (three lone pairs)
- Bonding electrons: 2 (one single bond)
- Formal charge = 6 – 6 – (2/2) = 6 – 6 – 1 = -1
- Oxygen with a double bond:
- Valence electrons: 6
- Non-bonding electrons: 4 (two lone pairs)
- Bonding electrons: 4 (double bond)
- Formal charge = 6 – 4 – (4/2) = 6 – 4 – 2 = 0
The student’s Lewis structure assigns incorrect formal charges. The more accurate distribution should be three oxygen atoms with a -1 charge each and one oxygen with a zero formal charge. The central chlorine atom should ideally have a formal charge of zero, which can be achieved if it forms resonance structures where double bonds with oxygen redistribute charge. This stabilizes the perchlorate ion and aligns with experimental data showing minimal formal charges on atoms due to electron delocalization across the ion.
