Final answer:
In decreasing nitrogen-oxygen bond length, the species are ordered as NO− (longest), followed by NO3−, and then NO+ (shortest). This order is based on the electronic structures, resonance, and bond orders of the ions, implying NO− has a bond length between a single and a double bond, NO3− bonds are slightly shorter, and NO+ has the shortest bond due to the highest bond order.
Step-by-step explanation:
To arrange the species NO+, NO3−, and NO− in decreasing nitrogen-oxygen bond length (longest to shortest), we must consider the electronic structure and resonance of each ion.
In NO− (nitrite ion), there is resonance between two structures, each with a different nitrogen-oxygen bond type (−N=O and −NO−), which results in equal bond lengths that are an average between a single and a double bond, implying a bond length shorter than a pure single bond but longer than a pure double bond.
The NO3− (nitrate ion) also displays resonance amongst its structures; however, it is well-known that bonds in NO3− are equivalent and slightly shorter than even the average N-O bond in NO− due to the delocalization over three oxygen atoms, suggesting that they are shorter than the equivalent bond in NO−.
Lastly, NO+ has a bond order greater than that of NO (neutral) due to the absence of an electron that would otherwise be in an antibonding orbital, leading to a bond length shorter than that of NO and, by extension, shorter than those in NO− and NO3−.
Therefore, the correct order from longest to shortest nitrogen-oxygen bond length is NO−, NO3−, NO+.