= Lewis structure
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= Lewis formula
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= Lewis dot structure
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Before calculating electronic <energy>, ask a simpler question: where can the available <valence electrons> go? A <Lewis structure> makes that budget visible, and <nitrate> shows why a correct budget can still require several <resonance structures>.
A <Lewis structure> records which <chemical atoms> are connected in a <molecule> or <ion> and how its <valence electrons> are allocated. One line represents one <bond pair>, two lines two <bond pairs>, and three lines three <bond pairs>; two dots beside a <chemical atom> represent one <lone pair>. The marks are a counting language, rather than measured <Electron> positions. See the https://goldbook.iupac.org/terms/view/L03513[IUPAC Gold Book, Lewis formula].
For <water>, $\mathrm H_2\mathrm O$, the budget is $6+2\times1=8$ <valence electrons>. Two O–H lines use four <Electrons>; the other four form two <lone pairs> on <oxygen>. For <ammonia>, $\mathrm{NH}_3$, $5+3\times1=8$ <valence electrons> give three N–H lines and one <lone pair> on <nitrogen>. Both <molecules> contain four <electron pairs>, but their allocations differ. Each shared <bond pair> is counted once in the total budget, even though both endpoints count it when checking the <octet rule>.
\Image[/lewis-water-ammonia.png]
{title=Eight <valence electrons> in each <molecule>: <water> has two <bond pairs> and two <lone pairs>; <ammonia> has three <bond pairs> and one <lone pair>}
{description=Each blue line represents two <Electrons>; each orange dot represents one <Electron>. The <oxygen> and <nitrogen> each count eight <Electrons> locally, while each <hydrogen> counts two. These are schematic <Lewis structures>, not predictions of molecular shape or <Electron> trajectories.}
A practical construction method is:
* Add the neutral <chemical atoms>' <valence electron> counts. Add one <Electron> for each unit of negative <electric charge>, or subtract one for each unit of positive <electric charge>.
* Choose candidate connections and put one <bond pair> on each. <Hydrogen> normally has one connection in these simple <Lewis structures>.
* Place remaining <Electrons> as <lone pairs>, checking the <octet rule>. If an endpoint has eight <Electrons> but the central <chemical atom> lacks them, try converting an endpoint <lone pair> into an additional <bond pair>, keeping the total budget fixed.
* Compute every <formal charge> and check its sum against the net <electric charge>. Compare reasonable candidates, including <chemical resonance> when several allocations share the same connections.
For <carbon dioxide>, $4+2\times6=16$ <valence electrons>. Starting with two single C–O connections leaves twelve <Electrons>, enough for three <lone pairs> on each <oxygen>, but only four <Electrons> around <carbon>. Converting one <lone pair> from each <oxygen> into a second <bond pair> gives $\mathrm{O=C=O}$. Eight <Electrons> are in the four <bond pairs>, and eight are in four <lone pairs>, two on each <oxygen>. <Carbon> has no <lone pair>; all three <chemical atoms> satisfy the <octet rule>, and all three <formal charges> vanish.
For the <nitrate ion>, $\mathrm{NO}_3^-$, the <valence electron> budget is
$$
5+3\times6+1=24.
$$
<Nitrogen> contributes five, the three <oxygen> <chemical atoms> contribute eighteen, and the negative <electric charge> adds one <Electron>. Start with <nitrogen> connected to each <oxygen> by one <bond pair>: those three lines use six <Electrons>. Put the remaining eighteen <Electrons> into three <lone pairs> on each <oxygen>. Each <oxygen> now satisfies the <octet rule>, but <nitrogen> counts only six <Electrons>. Reassign one <oxygen> <lone pair> as a second <bond pair> on its N–O connection. This changes the allocation without adding an <Electron>; it applies the construction procedure in https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures[Flowers et al., Chemistry 2e, 7.3: drawing Lewis structures].
The resulting <nitrate> <Lewis structure> has one N=O connection and two N–O connections. Its four <bond pairs> use eight <Electrons>, and its eight <lone pairs> use sixteen: $8+16=24$. The doubly connected <oxygen> has two <lone pairs>; each singly connected <oxygen> has three; <nitrogen> has none. All four <chemical atoms> satisfy the <octet rule>. Any of the three <oxygen> <chemical atoms> can supply the additional <bond pair>, giving the three equivalent <resonance structures> below.
\Image[/nitrate-resonance-contributors.png]
{title=Three equivalent <nitrate> <resonance structures>, each with 24 <valence electrons>}
{description=The same labelled <chemical atoms> occupy the same positions in all three <Lewis structures>. Each line counts two <Electrons>; each dot counts one. Each drawing has four <bond pairs>, eight <lone pairs>, <formal charge> $+1$ on <nitrogen>, $-1$ on each singly connected <oxygen>, and zero on the doubly connected <oxygen>. The brackets give the net <electric charge>; the arrows indicate <chemical resonance>, not a <chemical reaction> or elapsed time.}
Why not draw three N=O connections to make the <ion> look symmetric? With two <lone pairs> on each <oxygen>, that drawing still spends $12+12=24$ <Electrons> and has the right total <formal charge>, but places twelve <Electrons> around <nitrogen>. It fails the <octet rule> for this second-period <chemical element>. Three equivalent <resonance structures> express the symmetry without that invalid local count.
The economy is beautiful: a handful of dots and lines rules out many impossible counts and identifies <lone pairs> that can participate in <acid-base chemistry>. A <Lewis structure> helps choose rational candidate connections before a detailed <quantum chemistry> calculation. It does not specify a complete many-<Electron> <wavefunction>, actual <electron density>, or molecular geometry. Adding the separate <VSEPR model>, described in https://openstax.org/books/chemistry-2e/pages/7-6-molecular-structure-and-polarity[Chemistry 2e, 7.6: molecular structure and polarity], predicts bent <water> and pyramidal <ammonia>; the layouts above alone make no such prediction. <Molecular orbital theory> supplies a richer electronic description, including the unpaired <Electrons> of <molecular oxygen> that an all-paired <Lewis structure> misses.
In 1916 Gilbert N. Lewis proposed two-dot notation for the shared <Electrons> of a <chemical bond> and fuller dot formulae for the outer <Electrons>, explicitly relating a shared pair to the existing single-line notation. See https://doi.org/10.1021/ja02261a002[Lewis, The Atom and the Molecule, JACS 38, 762–785 (1916), pp. 777–778], available as an https://zenodo.org/records/1429068[archived original paper]{ref}; the relevant passage is also readable in a https://www.chemteam.info/Chem-History/Lewis-1916/Lewis-1916.html[transcription of the original paper]{ref}. The modern <quantum mechanics> interpretation should be distinguished from his historical model of fixed outer <Electrons>.
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