A Lewis structure is a diagram representing the arrangement of valence electrons in a molecule or polyatomic ion. Chemical symbols identify the atoms, dots represent electrons, and lines commonly represent shared electron pairs forming covalent bonds. Nonbonding electrons and formal charges are included where appropriate. The notation describes connectivity and electron bookkeeping rather than the complete three-dimensional distribution of electrons. (goldbook.iupac.org)
Historical basis and notation
The method is named after Gilbert N. Lewis, whose paper The Atom and the Molecule, published in April 1916, developed the shared-electron-pair description of the chemical bond. This framework preceded modern quantum-mechanical accounts of bonding. (pubs.acs.org)
A single bond is represented by one line or one pair of dots between atoms. Double and triple bonds contain two and three shared pairs, respectively. A lone pair consists of two valence electrons assigned to one atom rather than shared in a bond. An unpaired electron appears as a single dot. Charges assigned to individual atoms are written beside their symbols; an ionic structure is ordinarily enclosed in brackets with its overall charge outside. (goldbook.iupac.org)
Lewis symbols for isolated atoms show their valence electrons around the elemental symbol. For many main-group elements, the number follows their group in the periodic table. Inner-shell electrons are normally omitted because the representation focuses on electrons involved in bonding. (openstax.org)
Constructing a structure
A typical construction procedure begins by adding the valence electrons of all constituent atoms. One electron is added for each unit of negative charge, or subtracted for each unit of positive charge. The atoms are then connected with single bonds. Hydrogen is terminal; among other atoms, lower electronegativity often favors a central position, although experimental connectivity may be needed. Remaining electrons complete terminal shells before being placed on central atoms. Lone pairs can become additional bonding pairs when necessary. The completed diagram must account for the original electron total. (openstax.org)
The octet rule guides many structures: atoms such as carbon, nitrogen, and oxygen commonly have eight electrons around them, counting both bonding and nonbonding electrons. Hydrogen instead has a two-electron shell. These are useful patterns, not universal requirements. (openstax.org)
For water, H₂O, eight valence electrons produce two O–H bonds and two lone pairs on oxygen. In ammonia, NH₃, three N–H bonds leave one lone pair on nitrogen. Carbon dioxide, CO₂, has sixteen valence electrons and is conventionally drawn O=C=O, with two lone pairs on each oxygen. Each example satisfies the relevant duet or octet pattern. (openstax.org)
Formal charge
Formal charge measures the electron allocation produced by assigning every nonbonding electron to its atom and dividing bonding electrons equally between the bonded atoms:
where is the neutral atom’s valence-electron count, its nonbonding-electron count, and its bonding-electron count. Formal charges sum to zero for a neutral molecule and to the overall electric charge for an ion. (openstax.org)
Among otherwise plausible structures, smaller formal-charge magnitudes and negative charges on more electronegative atoms are generally favored. Formal charge is a bookkeeping quantity, not a measurement of local charge. It differs from oxidation state, which assigns bonding electrons according to an ionic approximation rather than equal sharing. (openstax.org)
Resonance
Sometimes multiple valid diagrams share the same atomic connectivity but differ in electron placement. Such diagrams are contributors to resonance, conventionally joined by a double-headed arrow, ↔. They are not separate molecules rapidly interconverting; the actual electronic structure is represented by a resonance hybrid. (openstax.org)
For ozone, O₃, two equivalent contributors place the double bond on opposite sides of the central oxygen. The actual O–O bonds are equivalent and intermediate between ordinary single and double bonds. Thus, the integer bond orders in individual Lewis diagrams do not always describe the observed bonding directly. (openstax.org)
Exceptions to the octet rule
Odd-electron species, or radicals, cannot give every atom a closed-shell octet; nitric oxide, NO, is an example. Electron-deficient structures include BeH₂ and BF₃, whose central atoms conventionally have four and six surrounding electrons. Expanded-valence diagrams, such as PF₅ and SF₆, place more than eight electrons around a central atom from the third period or beyond. These categories demonstrate why octet counting alone cannot determine every structure. (moodle.cooper.edu)
Geometry and limitations
Lewis structures supply the bonding and lone-pair information used by valence-shell electron-pair repulsion theory to predict molecular shape. Electron-domain geometry includes lone pairs, whereas molecular geometry describes atomic positions. Water is bent, ammonia is trigonal pyramidal, and carbon dioxide is linear. Shape and bond polarity together determine whether bond dipoles cancel. (openstax.org)
The notation does not calculate electron energies, detailed electron distributions, or magnetic behavior. The conventional O=O diagram pairs all electrons, yet ground-state dioxygen has two unpaired electrons and exhibits paramagnetism. Molecular orbital theory, grounded in quantum mechanics, accounts for this behavior using orbitals extending across the molecule rather than only localized bonding pairs. (openstax.org)