8.86 The substance chlorine monoxide, ClO(g), is important in atmospheric processes that lead to depletion of the ozone layer. The ClO molecule has a dipole moment of 1.24 D and the Cl-O bond length is 1.60 Å. (a) Determine the magnitude of the charges on the Cl and O atoms in units of the electronic charge, e. (b) Based on the electronegativities of the elements, which atom would you expect to have a negative charge in the ClO molecule? (c) By using formal charges as a guide, propose the dominant Lewis structure for the molecule. Are the formal charges consistent with your answers to parts (a) and (b)? Can you reconcile any differences you find?
The Correct Answer and Explanation is :
Solution: Chlorine Monoxide (ClO) Analysis
(a) Determining the Magnitude of the Charges on Cl and O in Units of Electronic Charge (e)
The dipole moment ((\mu)) of ClO is given as 1.24 D (Debye). The bond length ((r)) is 1.60 Å (1.60 × 10⁻¹⁰ m). The relationship between dipole moment and charge separation is:
[
\mu = q \cdot r
]
where:
- ( \mu = 1.24 ) D ( = 1.24 \times 3.336 \times 10^{-30} ) C·m
- ( r = 1.60 \times 10^{-10} ) m
- ( q ) is the charge in Coulombs.
Solving for ( q ):
[
q = \frac{\mu}{r} = \frac{(1.24 \times 3.336 \times 10^{-30})}{(1.60 \times 10^{-10})}
]
[
q = 2.59 \times 10^{-20} \text{ C}
]
Since the elementary charge ( e ) is 1.602 × 10⁻¹⁹ C, the charge in units of ( e ) is:
[
q = \frac{2.59 \times 10^{-20}}{1.602 \times 10^{-19}} = 0.162 e
]
Thus, the magnitude of the charge on Cl and O is 0.162 e.
(b) Which Atom Carries the Negative Charge?
Electronegativity values:
- Chlorine (Cl): 3.16
- Oxygen (O): 3.44
Since oxygen has a higher electronegativity than chlorine, it will attract more electron density, meaning Oxygen (O) carries the negative charge while Chlorine (Cl) carries the positive charge.
(c) Dominant Lewis Structure and Formal Charges
The dominant Lewis structure for ClO consists of a single bond between Cl and O, with one lone pair on Cl and three lone pairs on O. A resonance structure with a Cl=O double bond is also possible.
To calculate formal charges:
[
\text{Formal Charge} = \text{Valence Electrons} – (\text{Lone Pair Electrons} + \frac{\text{Bonding Electrons}}{2})
]
For Cl (7 valence electrons):
- Cl in Cl–O (single bond): Formal charge = 7 – (4 + 1) = +2
- Cl in Cl=O (double bond): Formal charge = 7 – (2 + 2) = +1
For O (6 valence electrons):
- O in Cl–O: Formal charge = 6 – (6 + 1) = -1
- O in Cl=O: Formal charge = 6 – (4 + 2) = 0
The single-bond structure (Cl–O) with a negative formal charge on O and positive charge on Cl aligns best with part (b), but a Cl=O resonance structure also contributes.
Reconciling Formal Charges with (a) and (b):
- The dipole calculation suggests partial charge separation (±0.162 e), which is much lower than full ionic charges.
- The formal charge suggests Cl is positive and O is negative, which agrees with the electronegativity argument.
- The reality is that Cl–O bonding is covalent, but with some ionic character, explaining why the formal charge is larger than the calculated partial charges.
Thus, the dominant Lewis structure (Cl–O) with a negative charge on O is consistent with electronegativity trends and dipole moment calculations.