How to draw CO32- lewis structure resonance?

If you find interesting the resonance structures of carbonate (CO₃^2-), a polyatomic molecular ion, then you may like reading this article till the end.

In this article, we will walk you through the possible resonance forms of CO₃^2-, how to draw them, and how the CO₃^2- resonance hybrid is formed from these contributing structures.

 So let’s start!

How many resonance structures are possible for CO₃^2-?

The following three resonance structures represent the carbonate (CO₃^2-) ion.

You can draw these resonance structures one by one, following the simple steps given below.

Steps to draw CO₃^2- resonance structures

1. Draw the initial Lewis structure of CO₃^2-

The initial Lewis structure of CO₃^2- is shown below.

It has a total of 24 valence electrons.

A carbon (C) atom at the center is surrounded by three oxygen (O) atoms at the sides.

Two O-atoms are single-covalently bonded to a C-atom at the center while the third O-atom is single bonded to the central C-atom.

The single-bonded O-atoms have 3 lone pairs each. In contrast, there are 2 lone pairs of electrons on the double-bonded O-atom. 

2. Identify the resonance contributing atoms

The terminal oxygen atoms containing lone pairs act as resonance-contributing atoms in CO₃^2-.

3. Determine electron flow and draw the resonance structures

The electron flow is shown using curved arrows.

The C=O double bond keeps shifting its position from one point to another, yielding three different CO₃^2- resonance structures.

Starting from the Lewis structure drawn in step 1, a lone pair from the O-atom (marked red) shifts to form a C=O double bond on the left side of CO₃^2- while the C=O bond at the top of the molecular ion breaks, placing an extra lone pair on the O-atom (marked blue).

Alternately, from Lewis structure 2, a lone pair from the O-atom (marked green) shifts to form a C=O double bond on the right side of CO₃^2-.

This increase in electron density is accompanied by a simultaneous decrease in electron density as the already present C=O double bond breaks into a single bond, placing an extra lone pair on the O-atom (marked red).

The above two structures along with the Lewis structure drawn in Step 1 make a total of three resonance structures for CO₃^2-.

4. Assign formal charges

The formal charges are assigned using the formula given below:

⇒ Formal charge = (Valence electrons – non-bonding electrons -1/2 bonding electrons)

Applying the formal charge formula on resonance structure I:

For C-atom:

• Valence electrons of carbon = 4
• Bonding electrons = 1 double bond + 2 single bonds = 4 + 2(2) = 8 electrons
• Non-bonding electrons = no lone pair = 0 electrons

Formal charge = 4-0-8/2 = 4-0-4 = 4-4 = 0

 For O-atom (marked red):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 single bond = 2 electrons
• Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons

Formal charge = 6-6-2/2 = 6-6-1= 6 – 7 = -1

 For O-atom (marked blue):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 double bond = 4 electrons
• Non-bonding electrons =2 lone pairs = 2(2) = 4 electrons

Formal charge = 6-4-4/2 = 6-4-2= 6 –6 = 0

For O-atom (marked green):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 single bond = 2 electrons
• Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons

Formal charge = 6-6-2/2 = 6-6-1= 6 – 7 = -1

Applying the formal charge formula on resonance structure II:

The formal charge present on the C-atom stays the same i.e., 0.

For O-atom (marked red):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 double bond = 4 electrons
• Non-bonding electrons =2 lone pairs = 2(2) = 4 electrons

Formal charge = 6-4-4/2 = 6-4-2= 6 –6 = 0

For O-atom (marked blue):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 single bond = 2 electrons
• Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons

Formal charge = 6-6-2/2 = 6-6-1= 6 – 7 = -1

For O-atom (marked green):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 single bond = 2 electrons
• Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons

Formal charge = 6-6-2/2 = 6-6-1= 6 – 7 = -1

Applying the formal charge formula on resonance structure III:

The formal charge present on the C-atom stays the same i.e., 0.

 For O-atom (marked red):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 single bond = 2 electrons
• Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons

Formal charge = 6-6-2/2 = 6-6-1= 6 –7 = -1

For O-atom (marked blue):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 single bond = 2 electrons
• Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons

Formal charge = 6-6-2/2 = 6-6-1= 6 –7 = -1

For O-atom (marked green):

• Valence electrons of oxygen = 6
• Bonding electrons = 1 double bond = 4 electrons
• Non-bonding electrons = 2 lone pairs = 2(2) = 4 electrons

Formal charge = 6-4-4/2 = 6-4-2= 6 –6 = 0

5. Combine the resonance structures

As a final step, the three resonance structures drawn above are combined using double-headed arrows and enclosed in square brackets to represent a single entity i.e., CO₃^2- as shown below. 

The actual CO₃^2- structure is a weighted average of the above resonance forms. It is known as the resonance hybrid. Each resonance form contributes equally to the resonance hybrid.

The pi-bonds are shown using dotted lines on the resonance hybrid, indicating that it keeps changing its position from one place to another on the ion.  

The net charge present on the resonance hybrid = -1 + (-1) = -2.

Therefore, the CO₃^2- resonance hybrid is enclosed in square brackets and a -2 formal charge is placed on the top right corner as shown below.

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