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Every Lewis structure you draw in Class 11, with its dots, dashes and lone pairs, goes back to one paper published in April 1916. Gilbert Newton Lewis proposed that a chemical bond is a pair of electrons shared between two atoms. This post covers who he was, what the 1916 paper changed, how to draw and check a Lewis structure using formal charge, and what a student can take from the way he worked.
Who was Gilbert N. Lewis?
The Science History Institute biography gives his dates as 1875 to 1946. He was taught at home until he was 14, earned a doctorate at Harvard under Theodore W. Richards, and worked in Germany with the physical chemists Walther Nernst and Wilhelm Ostwald. He later held posts at MIT and the University of California, Berkeley, where he expanded chemistry and chemical engineering.
Bonding was only part of his work. He also turned the thermodynamics of Josiah Willard Gibbs into tools chemists could use to predict whether a reaction goes to completion or stops at equilibrium. The same biography notes that he was nominated 35 times for a Nobel Prize and never won.
What did the 1916 paper propose?
Lewis had been thinking about electrons in atoms since at least 1902. The Science History Institute shows a memorandum from that year in which he drew atoms as nested cubes with electrons at the corners. This "cubic atom" explained the eight groups of the periodic table and the idea that atoms tend towards a complete set of eight outer electrons, the octet.
In The Atom and the Molecule, published in the Journal of the American Chemical Society in April 1916, he made the key step: a bond can be a pair of electrons held jointly by two atoms. Irving Langmuir developed the idea further and introduced the term covalent bond. NCERT's Class 11 chapter on chemical bonding still opens with the Kössel-Lewis approach, Lewis symbols and the octet.
A review by Ronald Gillespie and Edward Robinson, Gilbert N. Lewis and the chemical bond, adds a point most textbooks skip. Lewis knew several exceptions to the octet rule and regarded it as less important than what he called the "rule of two", the electron pair. The authors argue that the pair has lasted, and is now understood as a consequence of the Pauli principle, while many more exceptions to the octet rule are known today than Lewis knew.
How do you draw and check a Lewis structure?
Worked example: carbon dioxide, CO₂.
Count valence electrons. Carbon has 4 and each oxygen has 6, so 4 + 6 + 6 = 16 electrons, or 8 pairs.
Put the least electronegative atom, carbon, in the middle and join each oxygen with a single bond. O–C–O uses 2 pairs, leaving 6.
Complete the outer atoms first. Three lone pairs on each oxygen use the remaining 6 pairs. Carbon now has only 4 electrons around it.
Move lone pairs into bonds until carbon has an octet. There are two ways: O=C=O, or O≡C–O.
To choose, use formal charge, as NCERT defines it: formal charge = valence electrons − lone-pair electrons − ½ × bonding electrons.
O=C=O: each oxygen is 6 − 4 − 2 = 0, and carbon is 4 − 0 − 4 = 0.
O≡C–O: the triple-bonded oxygen is 6 − 2 − 3 = +1, carbon is 4 − 0 − 4 = 0, and the single-bonded oxygen is 6 − 6 − 1 = −1.
The structure with the smallest formal charges, O=C=O, is preferred.
In 1923 Lewis carried the same idea into acids and bases. As NCERT's equilibrium unit puts it, he defined an acid as a species that accepts an electron pair and a base as one that donates a pair. Its example is BF₃ with NH₃. Boron in BF₃ has three bonds, so only six electrons. Nitrogen in NH₃ has a lone pair. That pair forms a B–N bond, giving F₃B–NH₃, and boron reaches eight.
What are the common mistakes with Lewis structures?
Forgetting the charge on ions. For NH₄⁺ subtract one electron (8 in total); for CO₃²⁻ add two (24 in total).
Forcing an octet on every atom. BF₃ has fewer than eight electrons on boron, and NCERT also covers expanded octets. Use formal charge rather than forcing the rule.
Treating formal charge as a real charge. It is bookkeeping for choosing between structures, not a measured charge on the atom.
Stopping at one structure. For ions such as NO₃⁻, no single Lewis structure describes the real ion; you need resonance.
What can a student take from how Lewis worked?
Keep notes on half-formed ideas. The 1902 cube sketch was wrong in its details, but it carried the seed of the octet for 14 years before the 1916 paper.
Know which part of a theory is the deep one. Lewis ranked the pair above the octet. Asking which rule matters most helps in any chapter.
Reuse a good idea. The same electron pair explained covalent bonds and, seven years later, acids and bases.
For the next chapter of the story, read our profile of Linus Pauling and the chemical bond, and for the shapes that electron pairs produce, our guide to VSEPR and molecular shapes. If ideas such as lone pairs and formal charge still feel shaky, the Fermi Concept Library covers every concept in chemistry, so you can go back to the one you are missing.
Try this now: draw the Lewis structure of the nitrate ion, NO₃⁻, and work out the formal charge on each atom.
