Predict the ordering, from shortest to longest, of the bond lengths in CO, CO2, and CO3^2-.
The Correct Answer and Explanation is:
To predict the ordering of bond lengths in CO, CO2, and CO3^2-, we need to consider their structures, resonance, and bond orders.
- CO (Carbon Monoxide):
In CO, the molecule consists of a triple bond between carbon and oxygen, with one lone pair on carbon and one lone pair on oxygen. The bond order here is 3 (triple bond). Typically, triple bonds are short because the atoms are held tightly together by the strong attraction between them. Thus, CO has the shortest bond length of the three molecules. - CO2 (Carbon Dioxide):
CO2 has two carbon-oxygen double bonds, with each oxygen atom having two lone pairs. The molecule is linear, and the bond order between carbon and each oxygen is 2. Double bonds are shorter than single bonds but longer than triple bonds. Therefore, the bond length in CO2 is longer than in CO, but shorter than in CO3^2-. - CO3^2- (Carbonate Ion):
CO3^2- is a resonance-stabilized molecule. In the resonance structures, the negative charge is delocalized across the three oxygens, and each carbon-oxygen bond is equivalent. The bond order in CO3^2- is effectively 4/3 because the electron density is spread over three bonds rather than concentrated in one. As a result, the bond length in CO3^2- is longer than in CO2 but shorter than in a single bond, and typically, the bond length is longer than both CO and CO2 due to this delocalization.
Order of bond lengths from shortest to longest:
CO < CO2 < CO3^2-
Thus, the order of bond lengths is dictated by bond orders and resonance effects, with the shortest bond in CO, followed by CO2, and the longest in CO3^2-.
