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To determine the formal charges in a molecule, we use the formula:
Formal Charge=(Valence Electrons)–(Nonbonding Electrons+12×Bonding Electrons)Formal \ Charge = (Valence \ Electrons) – (Nonbonding \ Electrons + \tfrac{1}{2} \times Bonding \ Electrons)Formal Charge=(Valence Electrons)–(Nonbonding Electrons+21×Bonding Electrons)
Carbon monoxide (CO) is a diatomic molecule with 10 total valence electrons (4 from carbon and 6 from oxygen). Its most stable Lewis structure involves a triple bond between carbon and oxygen, with one lone pair on carbon and one lone pair on oxygen.
Let’s calculate the formal charges step by step:
Carbon (C):
Valence electrons = 4
Nonbonding electrons = 2 (one lone pair)
Bonding electrons = 6 (three bonds with oxygen)
Formal Charge=4–(2+62)=4–(2+3)=–1Formal \ Charge = 4 – (2 + \tfrac{6}{2}) = 4 – (2 + 3) = –1Formal Charge=4–(2+26)=4–(2+3)=–1
Oxygen (O):
Valence electrons = 6
Nonbonding electrons = 2 (one lone pair)
Bonding electrons = 6 (three bonds with carbon)
Formal Charge=6–(2+62)=6–(2+3)=+1Formal \ Charge = 6 – (2 + \tfrac{6}{2}) = 6 – (2 + 3) = +1Formal Charge=6–(2+26)=6–(2+3)=+1
Thus, in CO:
Carbon carries a –1 charge
Oxygen carries a +1 charge
Although oxygen is more electronegative, the formal charge distribution shows carbon as slightly negative and oxygen as slightly positive. This unusual charge distribution explains why in many chemical reactions, carbon in CO acts as the electron donor (nucleophilic center), such as in metal–carbonyl complexes.
The net formal charge of the molecule is zero, which is consistent with its neutral character. However, the internal separation of charges contributes to CO’s bonding properties, dipole moment, and its ability to form strong coordinate bonds with transition metals.
Therefore, the correct formal charges in CO are: Carbon = –1 and Oxygen = +1 (Option C).
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