Draw all reasonable Lewis structures for the oxalate ion C2O4^2-? How many resonance structures exist for this ion? Assign formal charges to all atoms in the ion.

The oxalate ion, C₂O₄^2-, has a total of 18 valence electrons. To draw the reasonable Lewis structures for this ion, we start by arranging the carbon and oxygen atoms. Each carbon atom will be bonded to three oxygen atoms.

One way to draw the structure is by having one carbon double bonded to an oxygen and single bonded to another carbon, which is then double bonded to another oxygen and single bonded to another oxygen. This gives us the following structure:

    O     O
     ||   ||
H–C==C–O
     |   |
     O   O

In this structure, the double bonds are between the carbon and the oxygen atoms, while the carbon–carbon bond is single. Every oxygen has two pairs of electrons, and each carbon atom is associated with four bonds in total, satisfying the octet rule.

Another resonance structure can be created by shifting the position of the double bonds. For instance, if we switch which oxygen atoms have double bonds, we get a different arrangement that is still valid. This structure looks similar but changes the location of the double bonds:

    O     O
     ||   ||
H–C–C==O
     |   |
     O   O

These two structures are the primary resonance forms of the oxalate ion. Therefore, we can conclude that there are two resonance structures for the oxalate ion.

Now, let’s assign formal charges to all atoms in the ion. The formula for formal charge is:

Formal Charge = Valence Electrons – Non-bonding Electrons – 0.5 * Bonding Electrons

For each carbon atom:

• Valence Electrons = 4
• Non-bonding Electrons = 0
• Bonding Electrons = 4 (from the two double bonds)
• Formal Charge = 4 – 0 – 0.5 * 4 = 0

For each double-bonded oxygen:

• Valence Electrons = 6
• Non-bonding Electrons = 4
• Bonding Electrons = 4 (two bonds)
• Formal Charge = 6 – 4 – 0.5 * 4 = 0

For each single-bonded oxygen:

• Valence Electrons = 6
• Non-bonding Electrons = 6
• Bonding Electrons = 2 (one bond)
• Formal Charge = 6 – 6 – 0.5 * 2 = -1

So, in total, the formal charges are:

• Carbons: Formal Charge = 0
• Double-bonded Oxygens: Formal Charge = 0
• Single-bonded Oxygens: Formal Charge = -1

This results in a net charge of -2 for the ion, which matches the overall charge of C₂O₄^2-.