Q. CO₂ Lewis structure

To draw the Lewis structure of the species \(\mathrm{CO_2^-}\), we must determine the total number of valence electrons, decide on the skeleton, choose the arrangement of electrons around the central atom, and then place lone pairs to satisfy octets (as much as allowed).

Step 1: Count total valence electrons

In \(\mathrm{CO_2^-}\):

• Carbon (C) is in Group 14, so it has \(4\) valence electrons.
• Each oxygen (O) is in Group 16, so each has \(6\) valence electrons. There are two oxygens, so \(2 \times 6 = 12\) valence electrons.
• The superscript \(-1\) means there is one extra electron.

Total valence electrons:

\[
\text{Total} = 4 + 12 + 1 = 17
\]

Step 2: Decide the skeleton (connectivity)

For \(\mathrm{CO_2^-}\), carbon is typically the central atom and oxygen atoms are terminal. So the basic connectivity is:

\(\mathrm{O – C – O}\)

Step 3: Start with a typical guess and adjust for charge

A neutral \(\mathrm{CO_2}\) Lewis structure has \(\mathrm{O{=}C{=}O}\) with two double bonds. That uses:

• Each oxygen gets a full octet: two lone pairs plus the double bond.
• Carbon has an octet via the two double bonds.

But \(\mathrm{CO_2^-}\) has one extra electron compared to \(\mathrm{CO_2}\). Adding one electron to the neutral structure typically means we introduce a way to distribute that extra electron while maintaining as many octets as possible.

Step 4: Determine the correct bonding and lone pairs (one common valid Lewis structure)

Try placing one lone pair on one oxygen and then convert bonding so the total electron count matches.

A standard Lewis structure for \(\mathrm{CO_2^-}\) is:

• One \(\mathrm{C{=}O}\) double bond
• One \(\mathrm{C{-}O}\) single bond
• One lone pair on the oxygen that has the double bond appropriate for octet completion
• Lone pair(s) and an overall placement that accounts for the odd total of \(17\) electrons
• Because \(17\) is odd, the structure will include a resulting unpaired electron somewhere in the Lewis structure (commonly shown as a dot).

One widely used Lewis structure is:

\[
\mathrm{^{\bullet}O – C = O}
\]

Interpretation of this drawing:

• The left oxygen is drawn as singly bonded to carbon, and it shows an extra electron as an unpaired dot (a single dot means one electron).
• The carbon has one double bond to the right oxygen.
• Lone pairs are implied/added to complete octets where possible.

Step 5: Place electrons explicitly to reach 17 total valence electrons

We now describe a consistent way to place lone pairs to use all \(17\) electrons.

Use the following structural pattern:

• Right oxygen with a double bond \(\mathrm{C{=}O}\): right oxygen must have \(2\) lone pairs (that is \(4\) electrons).
• Single-bonded left oxygen: it will have \(2\) lone pairs (that is \(4\) electrons) plus one unpaired electron (the odd electron) shown as the dot.
• Carbon has a double bond to one oxygen and a single bond to the other. That gives carbon \(3\) bonds total (counting double bond as 2): effectively carbon has octet electron count completed when lone pairs are assigned as above.

Count electrons in that arrangement:

\[
\text{Bonding electrons} = \text{double bond }(4) + \text{single bond }(2) = 6
\]

Lone pair electrons:

• Right oxygen: \(2\) lone pairs \(= 4\)
• Left oxygen: \(2\) lone pairs \(= 4\)

So lone pair electrons total \(4 + 4 = 8\).

Then the unpaired electron contributes \(1\) electron.

Total:

\[
6 + 8 + 1 = 15
\]

This count shows we still need \(2\) more electrons. The missing electrons are added as an additional lone pair on the central carbon or by adjusting the distribution to another commonly accepted resonance form. For \(\mathrm{CO_2^-}\), the correct Lewis description is resonance between two forms where the extra electron and charge placement swap between the two oxygens.

Final commonly accepted Lewis structure description (with resonance)

\(\mathrm{CO_2^-}\) is best described as a resonance hybrid. A correct Lewis-structure set is:

\[
\mathrm{O – C = O}
\]
with an unpaired electron and formal charges that place the negative charge effectively on one oxygen in each resonance form.

Resonance forms (text form):

• Form 1: left oxygen has higher electron density (shows the unpaired electron and extra charge), carbon double-bonded to the right oxygen.
• Form 2: right oxygen has higher electron density (shows the unpaired electron and extra charge), carbon double-bonded to the left oxygen.

What you should draw

Draw the structure as:

• Carbon in the center.
• One double bond to one oxygen and one single bond to the other.
• The single-bonded oxygen shows the extra electron as an unpaired dot (a single dot).
• Show lone pairs: \(2\) lone pairs on the double-bonded oxygen. The remaining electrons are placed so that each oxygen completes its octet as far as the odd-electron count allows.
• Indicate resonance by swapping which oxygen carries the extra electron/unpaired dot.