Lewis structures are diagrams that represent the valence electrons and covalent bonding arrangement within a molecule or polyatomic ion. These drawings, sometimes called Lewis dot formulas, illustrate how atoms share electron pairs to form bonds and show non-bonding electrons, known as lone pairs. Constructing a Lewis structure helps predict a molecule’s geometry, electron distribution, and overall chemical stability, serving as a foundational tool for understanding its properties and reactivity.
The Foundation: Counting Valence Electrons
The process of drawing a Lewis structure begins with an accurate count of all available valence electrons from every atom in the chemical species. Valence electrons are the electrons in the outermost shell of an atom, and their number corresponds to the atom’s group number on the periodic table. To find the total number for the entire molecule, sum the valence electrons contributed by each individual atom.
This initial calculation requires specific adjustments if the species is an ion, meaning it carries a net electrical charge. For an anion (negatively charged), one electron must be added to the total valence count for every unit of negative charge. Conversely, for a cation (positively charged), one electron must be subtracted from the total for every unit of positive charge. This final, corrected total represents the exact number of electrons that must be placed in the completed Lewis structure.
Building the Skeletal Structure
Once the total electron count is established, determine the arrangement of the atoms to form the skeletal structure of the molecule. A single central atom is typically surrounded by the other atoms, referred to as terminal atoms. The choice for the central atom is usually the one with the lowest electronegativity.
Hydrogen atoms are consistently placed on the periphery because they can only form a single bond and are therefore never the central atom. Similarly, the most electronegative element, fluorine, is nearly always a terminal atom. Connect each terminal atom to the center using a single line, which represents a single covalent bond.
Each single bond accounts for two shared electrons. Subtract the total number of electrons used in the initial framework from the total valence count. The remaining pool of valence electrons will then be distributed in the subsequent steps.
Completing the Octets and Forming Multiple Bonds
The distribution of the remaining valence electrons begins with the terminal atoms to satisfy the octet rule. The octet rule states that most main group atoms strive to achieve a stable configuration of eight electrons in their outermost shell. Electrons are placed as lone pairs—non-bonding pairs—on the terminal atoms until each one is surrounded by eight electrons, counting both the lone pairs and the electrons in the single bond connecting it to the central atom.
An exception to the octet rule is hydrogen, which only requires two electrons (a duplet). Once all terminal atoms have satisfied their electron requirements, any remaining electrons must be placed on the central atom, typically as lone pairs. Placing these leftover electrons on the central atom is particularly relevant for elements in the third period and beyond, as they can sometimes accommodate more than eight valence electrons in an expanded octet.
If the central atom still has fewer than eight electrons after all valence electrons have been placed, lone pairs from terminal atoms must be converted into shared pairs to form multiple bonds. Converting a lone pair creates a double bond (four shared electrons) or a triple bond (six shared electrons). This conversion must be repeated until the central atom achieves a complete octet without exceeding the total valence electron count.
Final Verification: Calculating Formal Charge
After a Lewis structure is drawn, verify its plausibility using the concept of formal charge. Formal charge is a theoretical charge assigned to each atom in a molecule, assuming that the electrons in a bond are shared equally between the two atoms. This calculation is performed for every atom using the formula: Formal Charge = (Valence Electrons) – (Non-bonding Electrons) – 1/2 (Bonding Electrons).
Formal charge serves as a bookkeeping tool to help determine the most stable and accurate representation when multiple valid Lewis structures are possible. The most stable structure minimizes the formal charges on all atoms, ideally resulting in a formal charge of zero for every atom. If formal charges cannot be zero, the most plausible structure places any negative formal charge on the most electronegative atom.
The sum of the formal charges on all atoms in the structure must equal the overall charge of the molecule or ion, providing a mathematical check of the initial valence electron count.