Nitrogen gas (\(\text{N}_2\)) makes up approximately 78% of the Earth’s atmosphere. This colorless and odorless gas is remarkably unreactive, a characteristic attributed to its highly stable molecular structure. This exceptional stability prompts questions about its bonding, specifically whether its structure can be represented by more than one arrangement of electrons, a concept known as resonance. Determining this requires a careful examination of the molecule’s accepted structure and the fundamental rules that govern electron distribution.
Understanding the Standard Structure of the Nitrogen Molecule (\(\text{N}_2\))
The fundamental structure of the nitrogen molecule is best described by its Lewis structure. Each nitrogen atom belongs to Group 15, meaning it possesses five valence electrons available for bonding. To achieve the stable electron configuration of a noble gas, each atom requires three additional electrons to complete its outer shell, known as the octet rule.
This requirement for three electrons from each atom results in the formation of a triple covalent bond between the two nitrogen atoms. A triple bond involves the sharing of three pairs of electrons, totaling six bonding electrons, which are positioned directly between the two nuclei. The remaining four valence electrons, two on each atom, are positioned as a non-bonding lone pair.
The accepted structure is therefore represented as \(\text{N}\equiv\text{N}\), with a single lone pair residing on each nitrogen atom. This configuration results in a perfectly symmetric, linear molecule. The total of eight electrons surrounding each nitrogen atom gives both atoms a formal charge of zero, which is the most preferred state for any molecule.
Defining Resonance Structures and Their Criteria
Resonance describes a phenomenon where the true electronic structure of a molecule or ion cannot be accurately depicted by a single Lewis structure. Instead, the actual structure is an average, or a hybrid, of several valid contributing structures, which are called resonance forms. This concept is fundamentally about electron delocalization, meaning the electrons are spread out over multiple atoms.
The basic condition for a molecule to exhibit resonance is that it must be possible to draw two or more distinct Lewis structures with the exact same arrangement of atoms. These structures must differ only in the placement of the non-bonding lone pair electrons or the electrons involved in pi bonds. The sigma bonds, which establish the atomic connectivity, must remain untouched.
Each contributing structure must also be a valid representation, meaning it must adhere to fundamental chemical rules, particularly the octet rule for second-row elements like nitrogen. An excellent example of resonance is the nitrite ion (\(\text{NO}_2^-\)), where the double bond could be drawn between the central nitrogen atom and either of the two oxygen atoms. Because both structures are chemically valid and equivalent, the actual structure is a hybrid where the electron density is equally distributed. The true structure is more stable than any single contributing form because the negative charge is spread out over the molecule.
Applying the Resonance Test to \(\text{N}_2\)
To determine if the nitrogen molecule has resonance structures, we must apply the criteria for resonance to its standard structure. The \(\text{N}_2\) molecule, with its \(\text{N}\equiv\text{N}\) triple bond, is perfectly symmetric and already satisfies the octet rule on both atoms with zero formal charge. The test requires proposing an alternative structure that maintains connectivity and the total number of valence electrons (ten) but moves the electrons.
One might attempt to draw a structure with a double bond (\(\text{N}=\text{N}\)) or a single bond (\(\text{N}-\text{N}\)). When electrons are rearranged to form a double bond, one nitrogen atom would have a positive formal charge, and the other would carry a negative formal charge. This charged arrangement is significantly less stable than the neutral structure with zero formal charges, rendering it an invalid resonance form.
Furthermore, any attempt to draw a valid Lewis structure for \(\text{N}_2\) that maintains the octet rule without the triple bond would invariably introduce substantial formal charges or violate the octet rule altogether. For instance, a hypothetical single bond structure would leave one or both nitrogen atoms with an incomplete octet, making it chemically unsound. Since no other valid Lewis structure can be drawn for \(\text{N}_2\) by simply moving non-bonding or pi electrons, the molecule fails the resonance test.
The stability of the nitrogen molecule, therefore, does not arise from electron delocalization across multiple structures. Instead, it comes from the sheer strength of the nitrogen-nitrogen triple bond. This bond is one of the strongest known in chemistry, requiring a large amount of energy to break, which directly accounts for the inert nature of \(\text{N}_2\) in the atmosphere. The standard \(\text{N}\equiv\text{N}\) structure is the one and only representation of the nitrogen molecule.