Q.Lewis dot structure of CO, NO2⁻ and CO3²⁻ are I, II and III respectively (see figure). Which of the above structure(s) is/are wrong?

Checking the electron count for CO, $NO_2^-$ and $CO_3^{2-}$ shows each structure described is a standard, chemically valid Lewis dot structure (using the common convention of showing only the overall ionic charge rather than individual formal charges).

I — CO: Total valence electrons $= 4(C) + 6(O) = 10$. A triple bond uses 6 electrons, leaving 4 electrons = one lone pair on C and one on O, exactly as shown. This is the accepted Lewis structure of carbon monoxide (giving formal charges $-1$ on C, $+1$ on O, consistent with its known triple-bond character).

II — $NO_2^-$: Total valence electrons $= 5(N) + 2(6)(O) + 1 = 18$. One N=O double bond + one N–O single bond uses 6 electrons; the remaining 12 electrons distribute as lone pairs to complete octets on both oxygens (and a lone pair on N, which also accounts for the ion's bent shape) — consistent with the standard resonance form of nitrite drawn in the figure.

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N2 must show a lone pair on both nitrogens; the option with a lone pair on only one N is the incorrect Lewis structure. Answer: (c).

Check each structure:

• (a) O=C=O with two lone pairs on each oxygen: correct Lewis structure of CO2 (carbon has 4 bonds, each O has 2 lone pairs).
• (b) N=O with an unpaired electron on N and lone pairs on O: nitric oxide (NO) is an odd-electron molecule, so this is acceptable. …

NH4Cl contains both covalent bonds (within the NH4+ ion) and an ionic bond (between NH4+ and Cl-).

In NH4Cl, the four N-H bonds within the ammonium ion (NH4+) are covalent (electron sharing), while the bond between the NH4+ cation and the Cl- anion is ionic (el …

Sulfur (Group 16, 6 valence electrons) forms one single bond to each hydrogen, using 4 of its 6 electrons for 2 lone pairs.

Step 1: Count total valence electrons.

• S: 6 valence electrons
• 2 H: 1 electron each = 2
• Total = 8 electrons = 4 electron pairs.

Step 2: Sulfur is the central atom (less electronegative than... well H is bonded terminally); place the 2 H atoms around S and form one single (2-electron) bond to each:

$H-S-H$

This uses 2 bond pairs (4 electrons).

Step 3: The remaining $8 - 4 = 4$ electrons are placed as 2 lone pairs on sulfur (satisfying the octet: 2 bond pairs + 2 lone pairs = 8 electrons around S).

Lewis structure:

$H-\overset{\displaystyle \cdot\cdot}{\underset{\displaystyle \cdot\cdot}{S}}-H$

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CO3^2- Lewis structure: C double-bonded to one O and single-bonded to two O⁻ atoms, with the negative charge and double-bond character delocalized over all three C–O bonds by resonance.

Total valence electrons in CO3^2-: $C(4) + 3\times O(6) + 2(\text{extra for charge}) = 4+18+2 = 24$ electrons (12 pairs).

One Lewis (contributing) structure: place carbon at the centre bonded to three oxygen atoms.

• One C–O bond is a double bond (C=O); that oxygen carries 2 lone pairs.
• The other two C–O bonds are single bonds; each of those oxygens carries 3 lone pairs and an overall −1 formal charge.
• This uses all 24 electrons: 4 bonding pairs (1 double + 2 single = 4 shared pairs = 8 electrons) + 8 lone pairs on the three oxygens (16 electrons) = 24 electrons. ✓

This structure has a net charge of −2 (each singly-bonded O is −1, the doubly bonded O and C are neutral), matching the ion's charge.

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A covalent bond is defined by the sharing of an electron pair between two atoms (option a), unlike an ionic bond which involves electron transfer.

According to the Lewis theory of bonding, a covalent bond forms when two atoms share one or more pairs of electrons so that both atoms attain a stable (often octet) electron configuration. For example, in H2, each hydrogen atom contributes one electron to the shared pair; in Cl2, each chlorine atom contributes one electron of the shar …