Lewis Structure Generator (with VSEPR Shape)

About the Lewis Structure Generator (with VSEPR Shape)

Type a formula — CO2, NH4+, SO4^2-, XeF4, NO3- — or a name such as water or sulfate, and get its Lewis structure drawn with bonds, lone pairs and formal charges. The calculator counts the valence electrons, tries every arrangement of single, double and triple bonds around the central atom and keeps the ones that satisfy the octet rule (with expanded octets from period 3 and the incomplete octets of boron and beryllium) and minimise the formal charges. When several structures are equally good it shows all the resonance forms and the average bond order.

With the structure come its VSEPR description — AXₙEₘ class, electron-pair geometry, molecular shape, ideal bond angles and hybridisation — and whether the molecule is polar. For species such as sulfate, where textbooks disagree between the formal-charge structure (two S=O) and the octet structure (four S–O single bonds), both are shown. Drawings can be downloaded as SVG or PNG. Molecules that need more than one central atom (C₂H₆, H₂O₂, oxoacids) get a clear explanation instead of a wrong picture.

Common uses

• Chemistry homework on Lewis structures, formal charge, resonance and VSEPR, with every count shown.
• Checking a hand-drawn structure: are the electrons counted right, and is the formal charge distribution the best one?
• Seeing why expanded and incomplete octets happen, and how textbooks differ on species such as sulfate and phosphate.
• Making clean diagrams of molecules and ions for notes, worksheets and slides.

How the structure is found

The steps follow OpenStax Chemistry 2e §7.3:

• Count the valence electrons of all atoms; add one for each negative charge, subtract one for each positive charge (SO₄²⁻: 6 + 4 × 6 + 2 = 32).
• Choose the central atom — usually the least electronegative, never hydrogen — and bond every other atom to it.
• Complete the octets of the outer atoms (a duet for hydrogen), then put any electrons left over on the central atom.
• Make multiple bonds if the central atom lacks an octet, by sharing lone pairs of the outer atoms.

The calculator does this for every possible combination of bonds and every possible central atom, then ranks the results by formal charge.

Formal charge and the best structure

Formal charge = valence electrons − lone-pair electrons − ½ × bonding electrons. It is bookkeeping, not a real charge, and the formal charges always add up to the charge of the species. Following OpenStax §7.4, the preferred structure has formal charges of zero where possible, otherwise the smallest ones, no like charges on neighbouring atoms, and negative charges on the more electronegative atoms. That is how N₂O comes out as N≡N–O rather than N=N=O, and why carbon is the central atom in SCN⁻.

Expanded octets and the sulfate question

Atoms from period 3 on (P, S, Cl, Xe …) can have more than eight electrons around them, as in PCl₅ (10), SF₆ (12) and XeF₄ (12). Minimising formal charge then gives sulfate two S=O double bonds and a formal charge of 0 on sulfur — the structure many textbooks, including OpenStax, draw. The octet structure, with four S–O single bonds and +2 on sulfur, is also widely used: quantum-chemical studies find that sulfur’s d orbitals take little part in the bonding (for example E. Magnusson, J. Am. Chem. Soc. 112 (1990) 7940), so many chemists regard the S–O bonds as polar single bonds. Both describe the same tetrahedral ion; the calculator shows both so you can use your course’s convention.

VSEPR shapes

Electron domains around the central atom — each lone pair and each bond, single or multiple — repel each other and spread out as far as possible (Gillespie and Hargittai, The VSEPR Model of Molecular Geometry, OpenStax §7.6). Two to seven domains give the linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral and pentagonal bipyramidal electron-pair geometries; the molecular shape describes only the atoms. Lone pairs take more room than bonds, so they squeeze the bond angles (H₂O: 104.5° instead of 109.5°), sit in the equatorial positions of a trigonal bipyramid (SF₄ seesaw, ClF₃ T-shaped, XeF₂ linear) and opposite each other in an octahedron (XeF₄ square planar).

Polarity

A bond is polar when the two atoms differ in electronegativity; a molecule is polar when its bond dipoles (and lone pairs) do not cancel. Symmetric shapes with identical outer atoms — linear CO₂, trigonal planar BF₃, tetrahedral CCl₄, square planar XeF₄ — are nonpolar even though their bonds are polar; bent H₂O, pyramidal NH₃ and seesaw SF₄ are polar. The table of bond dipoles uses Pauling electronegativities.

Sources

• OpenStax, Chemistry 2e: §7.3 Lewis Symbols and Structures, §7.4 Formal Charges and Resonance and §7.6 Molecular Structure and Polarity.
• R. J. Gillespie and I. Hargittai, The VSEPR Model of Molecular Geometry (Allyn and Bacon, 1991).
• E. Magnusson, “Hypercoordinate molecules of second-row elements: d functions or d orbitals?”, J. Am. Chem. Soc. 112 (1990) 7940–7951.
• IUPAC Compendium of Chemical Terminology (Gold Book): “Lewis formula”, “formal charge”.
• Pauling electronegativities: PubChem periodic table (NCBI).

Limitations

• Only species built around one central atom are drawn: molecules with chains (C₂H₆, H₂O₂, N₂H₄) or with hydrogen on oxygen (oxoacids, alcohols) need more than one central atom.
• Transition-metal compounds and ionic compounds are outside the Lewis model used here; ions can be drawn one at a time (Cl-, NH4+).
• Lewis structures cannot show everything: O₂ comes out with all electrons paired although it is paramagnetic, and resonance is shown as separate forms of what is really one averaged structure.
• Bond angles are VSEPR’s ideal values with the direction of any distortion; measured angles are quoted only for a few molecules (H₂O, NH₃, CH₂O).

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