Lewis structures are fundamental diagrams used in chemistry to visually represent the distribution of valence electrons around atoms in a molecule. These diagrams, sometimes called electron dot structures, illustrate the bonding between atoms and any unshared electron pairs (lone pairs) that may be present. By mapping the arrangement of these outermost electrons, Lewis structures help predict a molecule’s geometry and understand its chemical reactivity. Although they are two-dimensional, they serve as the starting point for theories that describe the three-dimensional shape of a molecule.
Essential Building Blocks: Valence Electrons and the Octet Rule
The process of drawing Lewis structures relies on accurately accounting for valence electrons, which are the electrons in an atom’s outermost shell. For main group elements, the number of valence electrons is determined by the element’s group number on the periodic table. For example, elements in Group 15, such as nitrogen and phosphorus, possess five valence electrons.
The concept of stability is governed by the Octet Rule, which states that main group atoms tend to form bonds until they are surrounded by eight valence electrons. This configuration mimics the electron arrangement of noble gases. The Duet Rule applies to hydrogen, which only requires two electrons to achieve a stable outer shell. Electrons are categorized as either bonding pairs (shared between atoms and represented by a line) or lone pairs (unshared and localized on a single atom, represented by two dots).
Step-by-Step Guide to Drawing Simple Lewis Structures
Lewis structure construction begins by calculating the total number of valence electrons in the molecule or ion. For an ion, this total must be adjusted by adding electrons for each negative charge or subtracting electrons for each positive charge. Once the total count is established, the atoms are arranged into a skeletal structure by placing the least electronegative atom in the center. Note that hydrogen can never be a central atom.
Next, a single covalent bond is drawn between the central atom and each surrounding terminal atom, with each line representing two shared electrons. For a molecule like phosphorus trichloride (\(\text{PCl}_3\)), phosphorus is the central atom bonded to three chlorine atoms. After accounting for the electrons used in these initial bonds, the remaining valence electrons are distributed as lone pairs to satisfy the octet of the terminal atoms first. In \(\text{PCl}_3\), each chlorine atom requires three lone pairs to complete its octet.
Any remaining electrons after the terminal atoms have achieved their octets are then placed onto the central atom as lone pairs. In the \(\text{PCl}_3\) example, 18 electrons are placed on the chlorine atoms, leaving a final lone pair for the central phosphorus atom. The final step involves checking that the central atom also possesses a complete octet. If it does not, lone pairs from the terminal atoms must be converted into double or triple bonds to satisfy the central atom’s eight-electron requirement.
Evaluating Stability: Using Formal Charge and Resonance
After drawing a Lewis structure, its plausibility and stability are evaluated using the concept of formal charge, a theoretical charge assigned to an atom in a molecule. Formal charge is calculated by subtracting the number of non-bonding electrons and half the number of bonding electrons from the atom’s original number of valence electrons. This accounting system assumes bonding electrons are shared equally between the bonded atoms. The most stable Lewis structure is generally the one that results in formal charges of zero for all atoms, or where the sum of the absolute formal charges is minimized.
If non-zero formal charges are unavoidable, the most stable structure is the one where any negative charge resides on the most electronegative atom. Formal charge helps differentiate between multiple possible structures that satisfy the octet rule, guiding the selection toward the most favorable arrangement. If a single Lewis structure cannot accurately describe the electron distribution, a phenomenon known as resonance occurs. This happens when the location of multiple bonds or lone pairs can be drawn in two or more equally valid ways, such as in the nitrate ion (\(\text{NO}_3^-\)).
These multiple valid representations, called resonance structures, are connected by a double-headed arrow. This indicates that the true electronic structure is not a single form but a hybrid or average of all contributing structures. This electron delocalization, where electrons are spread out over several atoms, increases the stability of the molecule. The resonance hybrid accurately reflects the equivalent bond lengths and charges observed experimentally.
Handling Structures That Deviate from the Octet Rule
While the octet rule is a reliable guideline, some molecules are stable even though their central atom does not follow the eight-electron rule. One category of exception is the incomplete octet, which occurs with elements from Groups 2 and 13, such as beryllium and boron. For instance, in boron trifluoride (\(\text{BF}_3\)), the central boron atom is shown with only six valence electrons, which avoids placing a positive formal charge on the highly electronegative fluorine atoms.
A second exception involves expanded octets, where the central atom is surrounded by more than eight valence electrons. This is only possible for elements in the third period and beyond, such as phosphorus and sulfur. These elements possess accessible d-orbitals that can accommodate extra electrons for bonding. For example, phosphorus pentachloride (\(\text{PCl}_5\)) has ten electrons around the central phosphorus atom.
Finally, odd-electron molecules, also known as free radicals, contain an odd total number of valence electrons. This makes it mathematically impossible for every atom to achieve a complete octet. These species, like nitrogen monoxide (\(\text{NO}\)), are reactive due to the presence of an unpaired electron.