A carbonate is a salt of carbonic acid, H 2CO 3, characterized by the presence of the carbonate ion, a polyatomic ion with the formula CO2−3. The word “carbonate” may also refer to a carbonate ester, an organic compound containing the carbonate group O=C(−O−)2.
To draw the Lewis structure for the carbonate ion (CO3^2-), follow these steps:
Count the Total Valence Electrons:
- Carbon (C) has 4 valence electrons.
- Oxygen (O) has 6 valence electrons, and there are three of them: 6 * 3 = 18.
- Since it’s an ion with a -2 charge, add 2 electrons.
- Total = 4 + 18 + 2 = 24 valence electrons.
Determine the Central Atom:
- Carbon is the least electronegative, so it will be the central atom.
Place the Atoms:
- Arrange the three oxygen atoms around the carbon atom.
Form Bonds:
- Use 6 electrons to form single bonds (C-O) between carbon and each oxygen. This uses up 6 electrons, leaving 18.
Place Remaining Electrons:
- Place the remaining electrons around the oxygen atoms to complete their octets. Each oxygen now has 6 electrons in bonds, so give each 2 more pairs (4 electrons), using up 12 more electrons. Now, 6 electrons remain.
Check Octets and Form Double Bonds if Necessary:
- Carbon has only 6 electrons around it (from the bonds), and oxygen atoms have their octets but one can share more electrons to form a double bond with carbon to complete carbon’s octet.
- Convert one of the lone pairs from one oxygen to make a double bond with carbon. This gives carbon 8 electrons.
Resonance Structures:
- The carbonate ion exhibits resonance. This means you can move one of the double bond’s electron pairs to another oxygen to create three equivalent structures:
Formal Charges
- To ensure the best structure, calculate formal charges:
- For O with single bond: 6 (valence) – 6 (non-bonding) – 1 (bond) = -1
- For O with double bond: 6 – 4 – 2 = 0
- For C: 4 – 0 – 4 (bonds) = 0
The actual structure of CO3^2- would be a resonance hybrid of these structures where each C-O bond is somewhere between a single and double bond, all equivalent, with an overall -2 charge distributed over the molecule.
Geometry & Hybridization
Geometry:
The carbonate ion (CO3^2-) has a trigonal planar geometry. Here are the details:
- Molecular Geometry: Since the carbon atom is bonded to three oxygen atoms and there are no lone pairs on the carbon, the molecule adopts a shape where these bonds are as far apart as possible, leading to a planar structure with bond angles of 120°.
Hybridization:
- Carbon Atom Hybridization: The carbon atom in CO3^2- undergoes sp^2 hybridization.
Here’s why:
- Orbitals Involved:
- Carbon has a 2s orbital and three 2p orbitals in its valence shell.
- One of the 2s electrons is excited to a 2p orbital, leaving one 2s and two 2p orbitals to hybridize.
- Hybridization Process:
- These three orbitals (one s and two p) mix to form three sp^2 hybrid orbitals.
- These hybrid orbitals are oriented in a plane at 120° to each other, which explains the trigonal planar geometry.
- Remaining Orbitals:
- The remaining unhybridized p-orbital on carbon can overlap with a p-orbital from one of the oxygen atoms to form a π (pi) bond. Due to resonance, this π bonding is distributed over all three C-O bonds, contributing to the partial double bond character in each.
Additional Notes:
- Bonding: Each oxygen atom has one double bond character with carbon due to resonance, although in any single Lewis structure, one might show a double bond and two single bonds.
- Electron Pair Geometry: Since there are three regions of electron density (the three bonds) around the carbon, the electron pair geometry is also trigonal planar.
This combination of geometry and hybridization allows for the delocalization of the π electrons over the entire ion, contributing to its stability and explaining why all three C-O bonds are equivalent in length and strength.