The nitrate ion, a polyatomic ion with the chemical formula \(\text{NO}_3^-\), is a fundamental structure found in common substances like fertilizers and explosives. This seemingly simple combination of one nitrogen atom and three oxygen atoms carrying a negative charge poses a challenge to standard chemical drawing rules. Chemists rely on a unique concept called resonance to accurately describe the behavior of this ion because its true structure cannot be represented by a single, static drawing. The need for multiple representations arises from the way the electrons are shared among the atoms, which is more complex than a straightforward single or double bond.
The Shortcomings of a Single Lewis Structure
The conventional method for drawing molecular structures, known as the Lewis structure, attempts to depict all valence electrons and bonds as fixed, localized pairs. When this method is applied to the nitrate ion, the resulting diagram must include one nitrogen-oxygen double bond and two nitrogen-oxygen single bonds to satisfy the octet rule for all atoms. This single drawing is a flawed representation because it predicts that the nitrate ion should have two different types of bonds.
A single Lewis structure would therefore lead us to expect one short N-O bond and two longer N-O bonds within the ion. However, experimental evidence reveals a contradictory truth. Measurements consistently show that all three nitrogen-oxygen bonds in the nitrate ion are identical in both length and strength. This uniformity indicates that the ion is perfectly symmetrical, making any single Lewis structure, which implies asymmetry, scientifically inadequate.
Understanding Electron Delocalization
To reconcile the symmetrical experimental data with the asymmetrical single-drawing prediction, chemists developed the conceptual tool of resonance structures. Resonance is used when a single Lewis structure is insufficient to describe a molecule’s true bonding and electron arrangement. The ion is not rapidly switching back and forth between the three possible Lewis structures, as the name “resonance” might suggest. Instead, the true structure is a single, unchanging entity known as the resonance hybrid.
The resonance hybrid represents an average of all the valid contributing Lewis structures. In the nitrate ion, the electrons that would form the double bond in any one drawing are actually spread out, or delocalized, over the entire molecule. This electron delocalization means the bonding electrons are not confined to a single pair of atoms but are shared among the central nitrogen and all three surrounding oxygen atoms.
The delocalization of electrons across multiple bonds is a stabilizing phenomenon, lowering the overall energy of the molecule. The electron density is distributed more evenly, which makes the ion less reactive than a structure with electrons localized in one specific double bond. This concept of shared electron density is the core reason why the traditional rule of drawing two atoms sharing only two electrons fails for the nitrate ion.
The True Structure of the Nitrate Ion
The concept of resonance allows us to accurately describe the true, symmetrical structure of the nitrate ion. By considering the three equivalent resonance structures, we can calculate the true nature of the nitrogen-oxygen bonds, which are neither single nor double. The three resonance structures collectively show four shared electron pairs distributed across three N-O bonding regions. This distribution results in a fractional bond order of \(4/3\), or approximately 1.33, for every single bond in the ion.
This fractional bond order of 1.33 explains the experimental finding that all three N-O bond lengths are equal. Each bond is an intermediate between a single and a double bond, resulting in a bond length that is shorter than a single bond but longer than a typical double bond. Furthermore, resonance structures explain the distribution of the ion’s negative charge. The overall negative charge of \(-1\) is not fixed on one or two oxygen atoms as the Lewis structures suggest.
Instead, the negative charge is equally distributed across all three oxygen atoms in the resonance hybrid structure. This results in each oxygen carrying a partial negative charge of \(-1/3\). The symmetrical distribution of both electron density and negative charge contributes significantly to the ion’s stability and its symmetrical, trigonal planar geometry, with all bond angles measuring \(120^\circ\).