Lewis Structure & VSEPR Shape Calculator

Water, H₂O. Valence electrons: 6 from oxygen + 2 × 1 from hydrogen = 8. Oxygen is central, because hydrogen can only form one bond. Two O–H bonds use 4 electrons; hydrogen is full with 2, so the other 4 go on oxygen as 2 lone pairs, giving oxygen its octet. Every formal charge is 0 (oxygen: 6 − 4 − 2). Two bonds plus two lone pairs is a steric number of 4 (AX₂E₂): tetrahedral electron geometry, a bent molecule with an angle below the ideal 109.5° (measured at 104.5°), sp³ oxygen, and polar.

Nitrate, NO₃⁻. 5 + 3 × 6 + 1 = 24 electrons. Three N–O bonds and full octets on the three oxygens use all 24, leaving nitrogen with only 6, so one oxygen lone pair becomes a second bond: N=O. Nitrogen then has formal charge 5 − 0 − 4 = +1 and each N–O oxygen 6 − 6 − 1 = −1, adding up to the ion’s −1. The double bond can sit on any of the three oxygens, so there are 3 resonance structures; three domains and no lone pair make it trigonal planar, 120°, sp².

The common mistake: ignoring lone pairs. XeF₄ has four atoms around xenon, so it is tempting to call it tetrahedral at 109.5°. Count the electrons: 8 + 4 × 7 = 36; the four bonds and the fluorine lone pairs use 32, so 4 electrons, 2 lone pairs, stay on xenon. The steric number is 6, not 4: the electron geometry is octahedral, the lone pairs sit opposite each other, and the molecule is square planar with 90° angles and sp³d² xenon.

Gilbert N. Lewis proposed in his 1916 paper “The Atom and the Molecule” that a chemical bond is a pair of electrons shared between two atoms, and drew atoms with their outer electrons as dots. Irving Langmuir took up the idea in 1919, named the stable group of eight the octet, and made the dot structures widely known. Linus Pauling added hybrid orbitals in 1931 and his electronegativity scale in 1932; the electronegativities used here are Pauling values from the PubChem periodic table.

The shapes came later. Nevil Sidgwick and Herbert Powell noticed in 1940 that the arrangement of bonds follows the number of electron pairs around the central atom, and Ronald Gillespie and Ronald Nyholm turned that into valence-shell electron-pair repulsion (VSEPR) theory in 1957, including the rule that lone pairs push harder than bonding pairs. Labels such as sp³d² for SF₆ are the traditional textbook description; modern calculations show that d orbitals contribute little and that “expanded octets” are better described by charge separation, which is why sulfate is drawn both with two S=O bonds (smallest formal charges) and with four S–O bonds (octet rule).

This calculator draws the Lewis structure of any molecule or ion with a single central atom made of main-group elements. It counts the valence electrons including the charge, puts the least electronegative non-hydrogen atom in the centre, then tries every arrangement of single, double and triple bonds and keeps the one that gives the centre its octet and the smallest formal charges, with negative charge on the most electronegative atoms. Atoms from period 3 down may hold more than 8 electrons; boron, beryllium and aluminium may hold fewer. Equally good arrangements are reported as resonance structures, and when an expanded octet is chosen, the octet-rule alternative is shown in Show Work.

From the steric number it gives the electron geometry, molecular shape, ideal bond angle and hybridisation, and it judges polarity from the vector sum of the electronegativity differences along the ideal bond directions together with any lopsidedness caused by lone pairs. Molecules with several central atoms (ethane, methanol, oxyacids such as H₂SO₄), ionic compounds and transition-metal compounds get a clear message instead of a wrong picture. Everything runs in your browser; nothing is sent anywhere.

Related tools: Oxidation Number Calculator, Bond Enthalpy Calculator, and Interactive Periodic Table.