Consider the following bond lengths: C-O 143 pm C=O 123 pm C ≡O 109 pm In the CO3^2- i

Consider the following bond lengths: C-O 143 pm C=O ...

Consider the following bond lengths: $$\mathrm{C}-\mathrm{O} \quad 143 \mathrm{pm} \quad \mathrm{C}=\mathrm{O} \quad 123 \mathrm{pm} \quad \mathrm{C} \equiv \mathrm{O} \quad 109 \mathrm{pm}$$ In the $\mathrm{CO}_{3}^{2-}$ ion, all three $\mathrm{C}-\mathrm{O}$ bonds have identical bond lengths of 136 $\mathrm{pm} .$ Why?

Consider the following bond lengths:
$$\mathrm{C}-\mathrm{O} \quad 143 \mathrm{pm} \quad \mathrm{C}=\mathrm{O} \quad 123 \mathrm{pm} \quad \mathrm{C} \equiv \mathrm{O} \quad 109 \mathrm{pm}$$
In the $\mathrm{CO}_{3}^{2-}$ ion, all three $\mathrm{C}-\mathrm{O}$ bonds have identical bond lengths of 136 $\mathrm{pm} .$ Why?

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Key Concepts

Resonance

Resonance is the concept that a single molecular structure can be represented by two or more contributing structures which differ only in the placement of electrons. In the carbonate ion, resonance illustrates that the double bond character is delocalized across all three C-O bonds, leading to equivalence among them.

Bond Order and Delocalization

Bond order refers to the number of chemical bonds between a pair of atoms, and in molecules with delocalized electrons, the effective bond order is the average of the contributing structures. In the carbonate ion, the electron delocalization distributes a partial double bond character to each C-O bond, resulting in bond lengths that are intermediate between typical single and double bonds.

Transcript

00:01 So in this video we're going to talk about question 98 in chapter 8, which says, consider the following bond lengths.

00:06 A carbon oxygen single bond is 143 pica meters.

00:10 A carbon oxygen double bond is 123 pbometers, and a carbon oxygen triple bond is 109 pkometers.

00:16 So in the co3 minus 2 ion, all three carbon oxygen bonds have identical bond lengths of 136 pachometers.

00:23 Why? so 136 pica meters is neither a single bond nor double bond, but somewhere in between.

00:30 And all three of our co bonds have identical bond links.

00:33 So it's almost like they're a hybrid of a double bond and single bond.

00:37 But we'll definitely make more sense of this if we just draw the lewis structure for co3 minus 2.

00:43 So in co3 minus 2, our carbon atom brings four valence electrons.

00:47 Each of our oxygen atoms brings six valence electrons, and then our minus 2 charge gives us an additional two electrons.

00:53 So if we add that all together, we have 24 electrons total to work with.

00:58 So let's go ahead and make our carbon.

01:00 Atom or central atom and then bond to each of our oxygen atoms.

01:05 If we do that, we can fill in our octets.

01:11 And that actually uses up all 24 of our electrons.

01:15 So is this a good lewis structure? well, we can check by checking the formal charges...

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