Nitrogen (N), element number seven, is a nonmetal whose chemical behavior is focused on achieving maximum stability. This stability is dictated by the arrangement of electrons in its outermost shell, known as the valence shell. Atoms strive to attain a full valence shell, following the octet rule, which typically means acquiring eight electrons to mimic the stable configuration of a noble gas.
The Goal: Achieving Stability
The neutral nitrogen atom has five electrons in its valence shell. Since a full valence shell requires eight electrons, nitrogen needs to gain three additional electrons to complete its octet. This need for three electrons dictates the way nitrogen interacts with other atoms, driving all of its bonding and chemical reactions. Acquiring these three electrons allows the atom to reach the highly stable \(2s^22p^6\) configuration, resembling the noble gas neon.
Covalent Sharing to Complete the Shell
Nitrogen most commonly achieves its stable octet by sharing electrons with other nonmetal atoms through covalent bonding. This sharing forms neutral molecules where the shared electrons count toward the octet of both atoms.
Nitrogen frequently bonds with itself to form the diatomic molecule, \(\text{N}_2\), which makes up about 78% of the Earth’s atmosphere. In \(\text{N}_2\), the two nitrogen atoms share three pairs of electrons, forming an exceptionally strong triple covalent bond. This completes the octet for both atoms, resulting in an extremely stable, inert gas.
Another common covalent compound is ammonia, \(\text{NH}_3\), where one nitrogen atom shares electrons with three hydrogen atoms. Hydrogen atoms only need one electron to fill their shell, so each contributes one electron to share with nitrogen. Nitrogen uses three of its five valence electrons to form three single covalent bonds. The remaining two valence electrons form a non-bonding or “lone” pair, which gives ammonia its characteristic trigonal pyramidal shape and dictates its behavior as a weak base.
Gaining Electrons: The Nitride Ion
A less frequent path to a full valence shell is for nitrogen to completely gain three electrons from a highly electropositive element, typically a metal. This transfer of electrons leads to the formation of a charged particle, or ion. The resulting structure is called the nitride ion, represented by the formula \(\text{N}^{3-}\). The negative three charge indicates that the nitrogen atom has successfully acquired the three electrons needed to complete its valence shell.
This process forms ionic compounds known as nitrides, where the negatively charged nitride ion bonds electrostatically with a positively charged metal cation. For example, lithium nitride, \(\text{Li}_3\text{N}\), forms when three lithium atoms each donate one electron to a single nitrogen atom. All alkaline earth metals form nitrides, though most alkali metals do not readily form stable nitrides. The formation of the nitride ion contrasts sharply with covalent sharing, as it involves the full transfer of electrons and results in a crystalline salt structure.
Stability and Reactivity of Completed Structures
Once nitrogen has filled its valence shell, the resulting structures exhibit different chemical properties than the elemental nitrogen atom. The molecules or ions formed are less reactive because they have achieved a stable, low-energy state.
For instance, the triple bond in atmospheric \(\text{N}_2\) is one of the strongest chemical bonds known, requiring nearly 1,000 kilojoules of energy per mole to break. This bond strength is why nitrogen gas is inert at room temperature, only reacting under extreme conditions. Similarly, ammonia (\(\text{NH}_3\)) and the ammonium ion (\(\text{NH}_4^+\)) are stable forms of nitrogen readily absorbed by plants in the nitrogen cycle. The stability of these completed structures means that substantial energy is required to force them to react, which is why common forms of nitrogen are so prevalent in nature.