20th century
In the early 20th century, the American chemist Gilbert N. Lewis began to use dots in lecture, while teaching undergraduates at Harvard, to represent the electrons around atoms. His students favored these drawings, which stimulated him in this direction. From these lectures, Lewis noted that elements with a certain number of electrons seemed to have a special stability.
Moreover, noting that a cube has eight corners Lewis envisioned an atom as having eight sides available for electrons, like the corner of a cube. Subsequently, in 1902 he devised a conception in which cubic atoms can bond on their sides to form cubic-structured molecules.
In other words, electron-pair bonds are formed when two atoms share an edge, as in structure
C below. This results in the sharing of two electrons. Similarly, charged ionic-bonds are formed by the transfer of an electron from one cube to another, without sharing an edge
A. An intermediate state
B where only one corner is shared was also postulated by Lewis.
Lewis cubic-atoms bonding to form
cubic-molecules Hence, double bonds are formed by sharing a face between two cubic atoms. This results in the sharing of four electrons.
In 1913, while working as the chair of the department of chemistry at the University of California, Berkeley, Lewis read a preliminary outline of paper by an English graduate student, Alfred Lauck Parson, who was visiting Berkeley for a year. In this paper, Parson suggested that the electron is not merely an electric charge but is also a small magnet (or "magneton" as he called it) and furthermore that a chemical bond results from two electrons being shared between two atoms.
[17] This, according to Lewis, meant that bonding occurred when two electrons formed a shared edge between two complete cubes.
On these views, in his famous 1916 article
The Atom and the Molecule, Lewis introduced the “Lewis structure” to represent atoms and molecules, where dots represent electrons and lines represent covalent bonds. In this article, he developed the concept of the electron-pair bond, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond.
In Lewis' own words:
| “ | An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively. | ” |
Moreover, he proposed that an atom tended to form an ion by gaining or losing the number of electrons needed to complete a cube. Thus, Lewis structures show each atom in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another; occasionally, pairs of dots are used instead of lines. Excess electrons that form lone pairs are represented as pair of dots, and are placed next to the atoms on which they reside:
Lewis dot structures of the Nitrite-ion
To summarize his views on his new bonding model, Lewis states:
[18]
The following year, in 1917, an unknown American undergraduate chemical engineer named Linus Pauling was learning the Dalton hook-and-eye bonding method at the Oregon Agricultural College, which was the vogue description of bonds between atoms at the time. Each atom had a certain number of hooks that allowed it to attach to other atoms, and a certain number of eyes that allowed other atoms to attach to it. A chemical bond resulted when a hook and eye connected. Pauling, however, wasn't satisfied with this archaic method and looked to the newly-emerging field of quantum physics for a new method.
In 1927, the physicits Fritz London and Walter Heitler applied the new quantum mechanics to the deal with the saturable, nondynamic forces of attraction and repulsion, i.e., exchange forces, of the hydrogen molecule. Their valence bond treatment of this problem, in their joint paper,
[19] was a landmark in that it brought chemistry under quantum mechanics. Their work was an influence on Pauling, who had just received his doctorate and visited Heitler and London in Zürich on a Guggenheim Fellowship .
Subsequently, in 1931, building on the work of Heitler and London and on theories found in Lewis' famous article, Pauling published his ground-breaking article "The Nature of the Chemical Bond"
[20] (see: manuscript) in which he used quantum mechanics to calculate properties and structures of molecules, such as angles between bonds and rotation about bonds. On these concepts, Pauling developed hybridization theory to account for bonds in molecules such as CH
4, in which four sp³ hybridised orbitals are overlapped by hydrogen's
1s orbital, yielding four sigma (σ) bonds. The four bonds are of the same length and strength, which yields a molecular structure as shown below:
A schematic presentation of hybrid orbitals overlapping hydrogens' s orbitals
Owing to these exceptional theories, Pauling won the 1954 Nobel Prize in Chemistry. Notably he has been the only person to ever win two unshared Nobel prizes, winning the Nobel Peace Prize in 1963.
In 1926, French physicist Jean Perrin received the Nobel Prize in physics for proving, conclusively, the existence of molecules. He did this by calculating Avogadro's number using three different methods, all involving liquid phase systems. First, he used a gamboge soap-like emulsion, second by doing experimental work on Brownian motion, and third by confirming Einstein’s theory of particle rotation in the liquid phase.
[21]
(Redirected from History of the molecule)
), earth (
), air (
), and water (
)) and "forces" of attraction and repulsion allowing the elements to interact. A modern conceptualization of molecules began to develop in the 19th century along with experimental evidence for pure chemical elements and how individual atoms of different chemical substances such as hydrogen and oxygen can combine to form chemically stable molecules such as water molecules.
