Chemistry · Chemical Bonding · NEET
Carbon has 4 valence electrons and oxygen has 6, so CO has 4 + 6 = 10 valence electrons total (5 pairs). You place these 5 pairs so both atoms get an octet. This is why CO ends up with a triple bond (3 shared pairs) plus one lone pair on C and one lone pair on O. Total = 6 bonding electrons + 4 lone-pair electrons = 10. This counting step is the first thing NEET expects you to get right.
With only 10 electrons, a double bond (C=O) would leave carbon with only 6 electrons around it, so carbon would NOT have an octet. To complete carbon's octet, one lone pair from oxygen is shared, making a third bond. So CO becomes :C≡O: — a triple bond. Both C and O then have 8 electrons each. Remember: a short, strong triple bond is why CO has one of the highest bond enthalpies known.
Add up: N = 5, each O = 6 (two oxygens = 12), and add 1 more electron for the negative charge. Total = 5 + 12 + 1 = 18 valence electrons = 9 pairs. The extra 1 electron comes from the -1 charge. Forgetting to add this charge electron is the single most common mistake NEET students make with ions.
NO2- carries an overall -1 charge because it has one extra electron. When you assign formal charges: nitrogen is 0, one oxygen (double bonded) is 0, and the single-bonded oxygen is -1. So the negative charge sits on the singly bonded oxygen. Because of resonance, this charge is actually shared equally by both oxygens in the real ion.
Formal charge = (valence electrons) - (lone-pair electrons) - ½(bonding electrons), written as FC = V - L - ½S. For the double-bonded O in NO2-: 6 - 4 - ½(4) = 0. For the single-bonded O: 6 - 6 - ½(2) = -1. For N: 5 - 2 - ½(6) = 0. The sum (0 + 0 - 1) = -1, which matches the ion charge — a good way to check your structure. NEET 2026 asked exactly this skill on the O3 molecule.
Yes, CO obeys the octet rule perfectly. Both carbon and oxygen have exactly 8 electrons around them. Carbon uses 3 bonds (6 electrons) plus 1 lone pair (2 electrons) = 8. Oxygen uses 3 bonds plus 1 lone pair = 8. The unusual part is that carbon looks trivalent here, but the octet is still complete because of the extra lone pair on carbon.
CO is a diatomic molecule, so it is simply linear (only two atoms). NO2- is bent (angular) with a bond angle near 115 degrees, because nitrogen has one lone pair that pushes the two oxygen atoms closer together. Knowing the shape helps in VSEPR questions, so pair this with VSEPR after you master the Lewis structure.
For the ozone (O3) Lewis structure shown — a central O atom (1) doubly bonded to one terminal O atom (2) and singly bonded to the other terminal O atom (3) — the correct formal charges on the oxygen atoms numbered 2, 1 and 3 respectively are:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
Yes, CO and N2 are isoelectronic — both have 10 valence electrons and a triple bond with one lone pair on each atom. That is why CO and N2 have very similar bond lengths and very high bond strengths. This isoelectronic link is a favourite exam point.
In one resonance structure of NO2-, nitrogen has 1 lone pair, the double-bonded oxygen has 2 lone pairs, and the single-bonded oxygen has 3 lone pairs — a total of 6 lone pairs. The single lone pair on nitrogen is what makes the ion bent.
NO2- has two equivalent oxygen atoms, so the double bond can be on either one, giving two resonance structures. CO has only two atoms and one fixed arrangement, so it has essentially one main structure (though the triple bond involves a coordinate bond).
Yes. The sum of all formal charges in a species must equal its overall charge: 0 for a neutral molecule like CO, and -1 for NO2-. This is a fast self-check — if your formal charges do not add up to the ion charge, your electron count is wrong.
