Q. \( \text{o}_2 \text{ 2- Lewis structure.} \)

Below, I show how to determine the Lewis structure of the doubly ion \( \mathrm{O_2^{2-}} \) step-by-step.

Step 1: Identify the species and its total charge.

\( \mathrm{O_2^{2-}} \) means two oxygen atoms with an overall charge of \( -2 \). So the molecule has two extra electrons compared to neutral \( \mathrm{O_2} \).

Step 2: Count total valence electrons.

Each oxygen is in Group 16, so each O has \( 6 \) valence electrons.

For two oxygens:

\[
6 + 6 = 12
\]

Add the extra electrons from the \( -2 \) charge:

\[
12 + 2 = 14
\]

Total valence electrons = \( 14 \).

Step 3: Draw the skeleton (connectivity).

With only two atoms, the skeleton is simply O–O.

Step 4: Place electrons to form a basic bond.

A common starting point is to place one bond between the oxygens (an O–O single bond). A single bond uses 2 electrons.

After placing an O–O single bond:

\[
14 – 2 = 12 \text{ electrons remaining}
\]

Now, give each oxygen an octet as best as possible using lone pairs.

Step 5: Distribute remaining electrons as lone pairs to satisfy octets.

There are 12 electrons left, which is 6 pairs.

Put lone pairs on the oxygen atoms. If we initially give each oxygen 3 lone pairs (which is typical for a doubly-bonding/charge adjustment process):

3 lone pairs on each oxygen uses:

\[
3 \text{ pairs} \times 2 \text{ oxygens} = 6 \text{ pairs} = 12 \text{ electrons}
\]

So all remaining electrons can be placed as lone pairs: each oxygen gets 3 lone pairs.

Step 6: Check octets and formal charges.

With an O–O single bond and 3 lone pairs on each oxygen, each oxygen has:

\( 2 \) electrons in the single bond plus \( 6 \) electrons in lone pairs \( = 8 \) electrons total around that atom.

So the octet rule is satisfied for both oxygens.

Step 7: Compute formal charges to confirm the structure.

Use the formal charge formula:

\[
\text{FC} = \text{valence e}^- – \left( \text{nonbonding e}^- + \frac{1}{2}\text{bonding e}^- \right)
\]

For each oxygen:

Valence electrons for O: \( 6 \)

Nonbonding electrons: \( 3 \) lone pairs \( = 6 \) nonbonding electrons.

Bonding electrons: one single bond means \( 2 \) bonding electrons, so half is \( 1 \).

So for each oxygen:

\[
\text{FC} = 6 – (6 + 1) = -1
\]

Each oxygen has formal charge \( -1 \). The total formal charge is:

\[
(-1) + (-1) = -2
\]

This matches the overall charge of \( \mathrm{O_2^{2-}} \).

Final Lewis structure (result).

The Lewis structure for \( \mathrm{O_2^{2-}} \) is:

• One single bond between the two oxygens.
• Three lone pairs on each oxygen.
• Each oxygen has a formal charge of \( -1 \).

A clear way to write it is:

\[
\mathrm{[:O:-]} \;-\; \mathrm{[:O:-]}
\]

Equivalently described in words: an O–O single bond with each oxygen carrying three lone pairs and formal charge \( -1 \) on each oxygen.