Lewis Structure and Diagram

FIELDS OF STUDY: Physical Chemistry; Inorganic Chemistry; Organic Chemistry

ABSTRACT

The visual representation of electron distribution known as Lewis structures and the Lewis dot diagram are discussed. The method represents the valence electrons in atoms and illustrates their combination in shared pairs to form bonds between atoms. Lewis structures are a simple means of relating molecular structure to valence.

Gilbert N. Lewis and Valence Theory

In 1916, American chemist Gilbert N. Lewis (1875–1946) published "The Atom and the Molecule," his first paper on the role of valence in chemical bonding. Electrons had been identified as charged subatomic particles just nineteen years earlier, and protons were about to be discovered the following year, while the existence of neutrons as the third subatomic particle would not be demonstrated until 1932. The modern theory of atomic structure was therefore in its very early stages and had not yet been defined by quantum mechanics, although the basic underlying mathematical principles had been developed. Any well-defined concept of atomic orbitals and the wave-particle duality of electrons (that is, their tendency to behave as both particles and waves) was still very controversial and very much open to discussion. These are all concepts now taken for granted by chemists and physicists the world over, but in Lewis’s time, they were very new ideas. The theoretical aspects did not yet have the support of experimental evidence. The electronic measurement devices to which he had access were crude by modern standards, predating the invention of the transistor by more than thirty years. Accordingly, explanations of chemical behavior based on electron shells and orbitals were subject to intense questioning by traditional chemists.

In this environment, Lewis proposed that atoms with an atomic mass greater than that of helium have inner shells of electrons with the same distribution as in the noble gas preceding them in the periodic table (as the six noble gases all have their outer electron shells filled and are therefore chemically inert). A sodium atom, for example, would have all of its electrons but one in the same configuration as the electrons in a neon atom; the extra electron would lie outside of these inner shells of electrons, in the outermost, or valence, shell. The natural corollary of this hypothesis was that the valence electrons of an atom could be easily given up to form a positively charged ion, or cation, with all of its electron shells filled in the same way as those of the preceding noble-gas element. Conversely, electrons could be added from other atoms to form a negatively charged ion, or anion, with all of its electron shells filled in the same way as those of the next noble-gas element in the periodic table.

Lewis Dot Structures

According to the modern theory of atomic structure, electrons are arranged in shells about the nucleus of an atom. The first shell can hold up to two electrons, the second shell can hold a total of eight, the third shell a total of eighteen, and the fourth shell a total of thirty-two; while subsequent shells could theoretically hold more, in practice, the fifth, sixth, and seventh electron shells also hold a maximum of thirty-two electrons. The outermost shell of an atom, however, only ever has a maximum of eight electrons (at least in theory—transition metals can have more), beyond which the next shell starts filling. For example, a calcium atom has twenty electrons, but it has two electrons in the first shell, eight in the second, eight in the third, and two in the fourth, rather than ten in the third and none in the fourth. The tendency of atoms to combine or form ions in such a way that they have eight electrons in their valence shells is known as the octet rule.

In a Lewis dot diagram of an atom, the chemical symbol of the element is shown, representing the nucleus and the inner electron shells, and is surrounded by up to eight dots, generally in pairs, representing the electrons in the atom’s outermost shell. For example, using calcium again, the Lewis diagram of this element consists of the element symbol, "Ca," with two dots next to it, representing the two electrons in the fourth shell.

In a Lewis structure of a molecule, the chemical symbols of the constituent elements are shown connected by lines in place of dots, each one representing the shared pair of electrons that forms a covalent bond between the atoms; a single line represents a single bond, while double and triple lines represent multiple bonds. (Ionic bonds are represented differently; an ionic compound is depicted as adjacent but separate ions.) Any electrons not involved in chemical bonds, be they unpaired electrons or lone pairs, are still shown as dots next to their respective elements. There are variations on this system, some of which retain the circles representing the atomic orbitals, others of which use the dots in place of lines to show the bonds between atoms in a compound; however, the system described here is the most widely used, as it permits a comprehensible two-dimensional representation of the molecule and the bond system it contains.

The Cubical Atom

Lewis’s 1916 paper introduced several important ideas: the sharing of electrons to form a covalent bond, the transfer of electrons from one atom to another to form an ionic bond, the octet rule, and of course his dot diagram. He also proposed that, counter to the planetary model of the atom introduced by Niels Bohr (1885–1962) in 1913, atoms were in fact cubical in shape, with valence electrons positioned at some or all of the cube’s eight corners. Two cubical atoms could form a single covalent bond by sharing a single edge so that they had two corners in common, while a double covalent bond was formed by the atoms sharing a full face, giving them four corners in common. Because this model could not account for triple bonds—two cubes cannot share more than four corners at one time—Lewis suggested that in some cases, the electrons of an atom would rearrange themselves from a cubical to a tetrahedral shape (a three-sided pyramid) with two electrons at each corner, allowing two atoms to share six electrons by sharing a single face.

While the cubical model of the atom was consistent with Lewis’s valence theory, it never found widespread acceptance, although Irving Langmuir (1881–1957) built on Lewis’s ideas to further refine valence theory and propose his own model of atomic structure. Lewis’s atomic model, like all other contemporary models, was eventually disproved and replaced by the quantum mechanical model, which was superficially similar to Bohr’s model but incorporated the idea of wave-particle duality and defined electron orbitals as the areas of the atom with the greatest probability of containing a given electron, rather than defined paths for electrons to follow around the nucleus.

PRINCIPAL TERMS

• covalent bond: a type of chemical bond in which electrons are shared between two adjacent atoms.
• electron: a fundamental subatomic particle with a single negative electrical charge, found in a large, diffuse cloud around the nucleus.
• lone pair: two valence electrons that share an orbital and are not involved in the formation of a chemical bond; also called a nonbonding pair.
• multiple bond: a bond formed by two atoms sharing two or more electron pairs; includes double bonds and triple bonds.
• shared pair: the two electrons shared between two atoms in a normal covalent bond.

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