Which of the following is the correct Lewis structure for ${\mathrm{CO}}_2$?

• A. $O=C=O$, with each oxygen atom having two lone pairs
• B. $O–C\equiv O$, with one oxygen atom having three lone pairs and the other having one lone pair
• C. $O–C–O$, with each oxygen atom having three lone pairs
• D. $O\equiv C–O$, with one oxygen atom having one lone pair and the other having three lone pairs

Answer: (A)

Verified step by step guidance

1. Step 1: Determine the total number of valence electrons available for CO\_2. Carbon has 4 valence electrons, and each oxygen has 6 valence electrons, so total valence electrons = 4 + 2 \times 6 = 16 electrons.
2. Step 2: Draw a skeletal structure with carbon as the central atom bonded to two oxygen atoms. Connect each oxygen to carbon with a single bond initially.
3. Step 3: Distribute the remaining valence electrons to satisfy the octet rule for the outer atoms (oxygen) first by adding lone pairs, then place any leftover electrons on the central atom (carbon).
4. Step 4: Check the octet rule for carbon. If carbon does not have a full octet, form double bonds by converting lone pairs from oxygen atoms into bonding pairs until carbon has 8 electrons around it.
5. Step 5: Confirm the formal charges on each atom to ensure the most stable Lewis structure. The best structure will have formal charges closest to zero, which typically results in carbon double bonded to each oxygen, with each oxygen having two lone pairs.