Chemistry 1916

The Atom and the Molecule

Gilbert N. Lewis

A chemical bond is a single pair of electrons two atoms hold in common.

Choose your version

In depth · the introduction

What if two atoms could each get what they want — a full outer shell — not by stealing electrons from each other, but by sharing a pair?

The big idea

Atoms are pickier about their outermost electrons than about anything else. By 1916 chemists knew that many atoms are most content with eight electrons in their outer layer — the "octet" — and the old explanation was that one atom hands electrons to another, leaving both with tidy shells. That works for salt, where sodium gives an electron to chlorine. But it fails for most of chemistry: the gases, the oils, the carbon compounds of life, where no atom is willing to simply give an electron away.

Gilbert Lewis offered a better idea. Two atoms can each reach eight by sharing a pair of electrons that belongs to both of them at once. That shared pair is the chemical bond. Share one pair and you have a single bond; share two or three and you have a double or triple bond. To keep track, Lewis drew each outer electron as a dot around the atom's symbol — the dot pictures still taught in every chemistry class.

How it came about

Lewis traced the idea back to a teaching sketch he had drawn in 1902, of an atom as a little cube with an electron possible at each of its eight corners — a full shell being a full cube. He sat on it for years. When he finally published in 1916, he was not alone: Walther Kossel in Germany worked out the electron-transfer side of the story the same year, and Lewis's rule of eight built on an earlier one by Richard Abegg.

Then Irving Langmuir, a brilliant and tireless speaker, took up Lewis's scheme, named the "octet," and carried it across the chemical world so forcefully that people began calling it Langmuir's theory — which stung Lewis, who had thought of it first. And for all his influence, Lewis was nominated for the Nobel Prize dozens of times and never won it, one of the most famous oversights in the prize's history.

Why it mattered

The shared pair turned a pile of unexplained facts into a single clear rule, and it gave chemists a way to draw any molecule and predict how it would behave. Almost everything a chemistry student learns to do on paper — figure out a molecule's shape, follow a reaction by pushing electron pairs around with arrows, judge whether a structure is reasonable — descends directly from Lewis's dots. It is hard to find an idea more woven into the daily practice of chemistry.

A way to picture it

Think of two people who each need a full pair of gloves but each owns only one glove of a matching pair. Neither wants to give a glove away. Instead they agree to hold one pair of gloves jointly between them — a pair that counts as "complete" for both. That jointly-held pair is the bond. If they need more warmth they can hold a second or third shared pair. And just as no hand can wear more than a certain number of gloves, no small atom can hold more than eight outer electrons — try to force more and the arrangement simply won't form.

Where it sits

Lewis came just after the periodic table (mendeleev-1869) had sorted the elements and just as physicists were taking the atom apart — Thomson's electron, Rutherford's nucleus, Bohr's shells. He used their electrons but kept a chemist's eye, asking not what the atom is but how it bonds. The deeper why — what makes a shared pair stick — came next, when quantum mechanics reached chemistry in Pauling's hands (pauling-1931). Lewis's dots are the bridge between the periodic table and the quantum chemical bond.

The original document

Original source text

G. N. Lewis · Journal of the American Chemical Society 38 (1916) 762–785 · communicated from the University of California

One picture for every bond

Lewis opens by rejecting the assumption that atoms in a molecule are always held together by the outright transfer of electrons from one to another. Transfer accounts for the salts — the strongly polar compounds — but not for the vast class of non-polar substances, nor for the fact that most stable molecules contain an even number of electrons. He sets out to cover both kinds of compound with a single model of the atom.

He states his account as a short series of postulates, paraphrased here: every atom has an inner kernel (the nucleus plus the inner electrons) that survives ordinary chemical change; around it sits an outer shell that can hold from zero to eight electrons; the shell tends to hold an even number, and especially eight, normally arranged at the eight corners of a cube; and two atoms can complete their shells at once by sharing electrons that belong to both.

…the total difference between the maximum negative and positive valences or polar numbers of an element is frequently eight, and is in no case more than eight.

From this regularity — drawn from Abegg's earlier rule — Lewis builds the theory of the cubical atom: eight valence electrons sit at the corners of a cube, and two atoms that share an edge share the two electrons on it. A shared edge is a single bond; the shared pair counts toward the octet of both atoms at once.

The decisive proposal is that a chemical bond is a pair of electrons held jointly by two atoms. Sharing — not transfer — is the ordinary non-polar bond, and one, two or three shared pairs give the single, double and triple bonds of organic chemistry. To record it on paper, Lewis marks each valence electron as a dot around the symbol of its element: the dot diagrams still in use today.

[ … ]

Lewis notes that the pairing of electrons — what later authors called the rule of two — may run deeper than the rule of eight itself; he could not say why electrons pair, a question answered only by quantum mechanics a decade later. The paper closes by applying the scheme to a range of molecules, polar and non-polar alike.

G. N. Lewis · University of California, Berkeley · 1916