Q.All the C—O bonds in carbonate ion (CO3^2-) are equal in length. Explain.

The equality of C—O bond lengths in the carbonate ion ($CO_3^{2-}$) is explained by the concept of resonance.

1. A single Lewis structure for the carbonate ion shows one C=O double bond and two C—O single bonds.
2. However, the $\pi$ electrons forming the double bond are not localized between a specific carbon and oxygen atom. Instead, they are delocalized over all three C—O bonds.
3. This delocalization leads to three equivalent resonance structures, where the double bond character is distributed equally among all three C—O bonds. …

When we draw Lewis structures for molecules, we aim to represent the arrangement of electrons and bonds. However, for some molecules or ions, a single Lewis structure cannot accurately depict the true bonding. This is where the concept of resonance becomes crucial. Resonance describes the delocalization of electrons within molecules where the bonding cannot be expressed by a single Lewis formula. The carbonate ion is a classic example of this phenomenon.

Here's a step-by-step explanation:

1. Draw the Lewis Structure for the Carbonate Ion ($CO_3^{2-}$):

   First, we determine the total number of valence electrons. Carbon has 4, each oxygen has 6, and the $2-$ charge adds 2 electrons.

   Total valence electrons $= 4 + (3 \times 6) + 2 = 4 + 18 + 2 = 24$ electrons.

   Carbon is the central atom, as it is less electronegative than oxygen. We connect the central carbon to the three oxygen atoms with single bonds.

   $O-C-O$

   $|$

   $O$

   This uses $3 \times 2 = 6$ electrons. We have $24 - 6 = 18$ electrons remaining.

   Next, we distribute the remaining 18 electrons to the outer oxygen atoms to satisfy their octets. Each oxygen needs 6 more electrons (3 lone pairs).

   $O-C-O$

   $|$

   $O$

   This uses $3 \times 6 = 18$ electrons. All 24 electrons are now used.

   Now, let's check the octet for the central carbon atom. It currently has only 6 electrons (from the three single bonds). To satisfy carbon's octet, one of the lone pairs from an oxygen atom must be moved to form a double bond with carbon. This also helps minimize formal charges.

   Let's calculate formal charges for one possible structure:

   • For Carbon: $4 - (0 \text{ lone pair electrons} + 1/2 \times 8 \text{ bonding electrons}) = 4 - 4 = 0$.
   • For Oxygen with double bond: $6 - (4 \text{ lone pair electrons} + 1/2 \times 4 \text{ bonding electrons}) = 6 - 4 - 2 = 0$.
   • For Oxygen with single bond: $6 - (6 \text{ lone pair electrons} + 1/2 \times 2 \text{ bonding electrons}) = 6 - 6 - 1 = -1$.
   • The total charge is $0 + 0 + (-1) + (-1) = -2$, which matches the ion's charge.

   So, a valid Lewis structure for $CO_3^{2-}$ shows one C=O double bond and two C-O single bonds, with the two singly bonded oxygens carrying a negative charge.

2. Identify Possible Resonance Structures:

   The double bond can be formed with any of the three oxygen atoms. This leads to three equivalent Lewis structures:

These individual structures are called **contributing structures** or **resonance forms**.

> [!WARNING]
> It is a common misconception that the molecule rapidly switches between these structures. This is incorrect. The molecule does not oscillate between these forms; rather, its true structure is a single, unchanging hybrid of all of them.

3. Understand the Resonance Hybrid: …