Lewis structures are chemical diagrams used for visualizing the arrangement of atoms and electrons within a molecule. They indicate which atoms are bonded together and show the location of all valence electrons, whether shared in a bond or existing as non-bonding lone pairs. The dinitrogen molecule (\(\text{N}_2\)), the primary component of Earth’s atmosphere, is a stable example characterized by a strong bond. Understanding its Lewis structure provides insight into this stability.
Foundational Concepts of Lewis Structures
Chemical bonding is governed by valence electrons, which are the outermost electrons of an atom. The number of these electrons is determined by an element’s group number on the periodic table. Atoms attempt to achieve the stable electron configuration of a noble gas, a tendency formalized by the Octet Rule. This rule states that atoms form bonds to be surrounded by eight valence electrons.
In a Lewis structure, a shared pair of electrons forms a covalent bond, typically drawn as a line. Electrons not involved in bonding are represented as dots, forming lone pairs. The goal is to arrange the total number of valence electrons so that the Octet Rule is satisfied for every atom, except hydrogen, which requires only two electrons.
General Procedure for Drawing Lewis Structures
The process begins by calculating the total number of valence electrons available from all atoms. For a neutral molecule, this is the sum of valence electrons contributed by each atom. If the structure is an ion, electrons must be added or subtracted to account for the overall charge.
Once the total count is established, the atoms are arranged into a basic skeletal structure. A single bond, representing two shared electrons, is drawn between connected atoms and subtracted from the total count. The remaining electrons are distributed as lone pairs, starting with the outer atoms to satisfy their octets. If any central atom still lacks an octet, lone pairs from the outer atoms must be shifted into the bonding region to form double or triple bonds until all atoms achieve a full octet.
Constructing the Lewis Structure for \(\text{N}_2\)
The procedure for \(\text{N}_2\) begins with the valence electron count. Nitrogen is in Group 15 of the periodic table, so each atom contributes five valence electrons, totaling ten electrons for the \(\text{N}_2\) molecule (\(5 \times 2 = 10\)).
The two nitrogen atoms are connected by a single bond, using two of the ten available electrons. This leaves eight electrons to distribute. Placing these eight electrons as lone pairs (four on each atom) results in each nitrogen atom being surrounded by only six electrons (four lone pair electrons + two bonding electrons). The Octet Rule is not satisfied.
To complete the octet for both atoms, two lone pairs—one from each nitrogen atom—must be moved into the bonding region between them. This converts the single bond into a triple bond (\(\text{N} \equiv \text{N}\)), sharing six electrons in total. The resulting structure has a triple bond and one lone pair remaining on each nitrogen atom. This final arrangement uses all ten valence electrons and ensures each nitrogen atom is surrounded by eight electrons, satisfying the Octet Rule.
Verifying the \(\text{N}_2\) Structure Using Formal Charge
While the triple-bond structure satisfies the Octet Rule, formal charge is used to confirm the most stable representation. Formal charge is a theoretical charge assigned to an atom, calculated by comparing its normal valence electrons to the electrons assigned in the Lewis structure. The calculation is: Formal Charge = (Valence Electrons) – (Non-bonding Electrons) – (1/2 Bonding Electrons).
For the nitrogen atom in the triple-bonded \(\text{N}_2\) structure, the calculation is direct. Each nitrogen atom has five valence electrons, two non-bonding electrons (one lone pair), and six bonding electrons (from the triple bond). Plugging these numbers into the formula yields: \(5 – 2 – (1/2 6) = 0\).
Since the formal charge on both nitrogen atoms is zero, and the overall molecule is neutral, this structure is confirmed as the most stable arrangement. Structures with minimized formal charges, especially zero on all atoms, represent the lowest-energy state for the molecule. This result validates the strong triple-bonded structure of dinitrogen.