Gilbert Newton Lewis
Lewis, Gilbert Newton
Lewis, Gilbert Newton
(b. West Newton, Massachusetts, 25 October 1875; d. Berkeley, California, 23 March 1946),
Lewis received his primary education at home from his parents, Frank Wesley Lewis, a lawyer of independent character, and Mary Burr White Lewis. He read at age three and was intellectually precocious. In 1884 his family moved to Lincoln, Nebraska, and in 1889 he received his first formal education at the university preparatory school. In 1893, after two years at the University of Nebraska, Lewis transferred to Harvard, where he obtained his B.S. in 1896. After a year of teaching at Phillips Academy in Andover, Lewis returned to Harvard to study with the physical chemist T. W. Richards and obtained his Ph.D. in 1899 with a dissertation on electrochemical potentials. After a year of teaching at Harvard, Lewis made the pilgrimage to Germany, the center of physical chemistry, and studied with W. Nernst at Göttingen and with W. Ostwald at Leipzig. Upon his return to Harvard in 1901, he was appointed instructor in thermodynamics and electrochemistry. In 1904 Lewis was granted a leave of absence and became a chemist with the Bureau of Weights and Measures in Manila. After one year (a year he seldom spoke of he joined the group of progressive young physical chemists around A. A. Noyes at the Massachusetts Institute of Technology.
Most of Lewis’ lasting interests originated during his Harvard years. The most important was thermodynamics, a subject in which Richards was very active at that time. Although most of the important thermodynamic relations were known by 1895, they were seen as isolated equations, and had not yet been rationalized as a logical system, from which, given one relation, the rest could be derived. Moreover, these relations were inexact, applying only to ideal chemical systems. These were two outstanding problems of theoretical thermodynamics. In two long and ambitious theoretical papers in 1900 and 1901, Lewis tried to provide a solution. He proposed the new idea of “escaping tendency” or fugacity (a term he coined), a function with the dimensions of pressure which expressed the tendency of a substance to pass from one chemical phase to another. Lewis believed that fugacity was the fundamental principle from which a system of real thermodynamic relations could be derived. This hope was not realized, though fugacity did find a lasting place in the description of real gases.
Lewis’ early papers also reveal an unusually advanced awareness of J. W. Gibbs’s and P. Duhem’s ideas of free energy and thermodynamic potential. These ideas were well known to physicists and mathematicians, but not to most practical chemists, who regarded them as abstruse and inapplicable to chemical systems. Most chemists relied on the familiar thermodynamics of heat (enthalpy) of Berthelot, Ostwald, and Van’t Hoff, and the calorimetric school. Heat of reaction is not, of course, a measure of the tendency of chemical changes to occur, and Lewis realized that only free energy and entropy could provide an exact chemical thermodynamics. He derived free energy from fugacity; he tried, without success, to obtain an exact expression for the entropy function, which in 1901 had not been defined at low temperatures. Richards too tried and failed, and not until Nernst succeeded in 1907 was it possible to calculate entropies unambiguously. Although Lewis’ fugacity-based system did not last, his early interest in free energy and entropy proved most fruitful, and much of his career was devoted to making these useful concepts accessible to practical chemists.
At Harvard, Lewis also wrote a theoretical paper on the thermodynamics of blackbody radiation in which he postulated that light has a pressure. He later revealed that he had been discouraged from pursuing this idea by his older, more conservative colleagues, who were unaware that W. Wien and others were successfully pursuing the same line of thought. Lewis’ paper remained unpublished; but his interest in radiation and quantum theory, and (later) in relativity, sprang from this early, aborted effort. From the start of his career, Lewis regarded himself as both chemist and physicist.
A third major interest that originated during Lewis’ Harvard years was is valence theory. In 1902, while trying to explain the laws of valence to his students, Lewis conceived the idea that atoms were built up of a concentric series of cubes with electrons at each corner. This “cubic atom” explained the cycle of eight elements in the periodic table and was in accord with the widely accepted belief that chemical bonds were formed by transfer of electrons to give each atom a complete set of eight. This electrochemical theory of valence found its most elaborate expression in the work of Richard Abegg in 1904, but Lewis’ version of this theory was the only one to be embodied in a concrete atomic model. Again Lewis’ theory did not interest his Harvard mentors, who, like most American chemists of that time, had no taste for such speculation. Lewis did not publish his theory of the cubic atom, but in 1916 it became an important part of his theory of the shared electron pair bond.
Extremely bright and precocious as a young man, Lewis was also shy and lacking in self-confidence. His ideas were unorthodox and singular, perhaps owing in part to his unusual education. He was disappointed and resentful that his talents were not appreciated, especially by Richards. In 1928 he refused a call to Harvard and in 1929 refused an honorary degree.
Lewis remained at M.I.T. for seven years and there laid the foundations for his important work in thermodynamics. It was well-known that the equilibrium position of any chemical system could in theory be predicted from free-energy data—an invaluable aid to both pure and applied chemistry. But existing free-energy data were mostly unreliable, contradictory, and spotty. Existing methods for measuring free energies were imperfect. In practice, application of thermodynamics to chemistry was extremely difficult even for the specialist. Early in his career Lewis began to pursue this neglected opportunity. In Manila, he determined the oxygen-water electromotive potential, a key datum for many chemical reactions. At M.I.T. he systematically studied the free energy of formation of compounds of oxygen, nitrogen, the halogens, sulfur, and the alkali metals. Particularly important was his measurement of the free energy of formation of simple organic compounds, beginning in 1912 with ammonium cyanate and urea. Lewis realized that the most fruitful use of free-energy data would be in complex organic reactions, an insight that was borne out in the 1920’s and 1930’s.
In 1907 Lewis set forth a new system of thermodynamics based on the new concept of activity. A function with the dimensions of concentration, activity expresses the tendency of substances to cause change in chemical systems. Lewis derived activity by generalizing the idea of fugacity, but also emphasized that it could be derived directly from free energy, since change in free energy is proportional to change in activity. Lewis showed that in terms of activity, all the familiar thermodynamic equations for ideal systems became “perfectly exact and general” for real systems. He also defined the important new concept of partial molal properties and a more exact form of Nernst’s equation for the potential of a single concentration cell. Both ideas proved very useful in treating real chemical systems.
Like fugacity, the conception of activity never played the central theoretical role that Lewis thought it would; but it proved indispensable for treating deviations from ideal behavior in real solutions. In general, Lewis’ main contribution to thermodynamics was not in grand theory but, rather, in its practical applications to real systems. He made chemists aware of the importance of hydration of ions and clarified the theory of liquid boundary potentials and conductivity, all of which derived from the practical necessities of measuring free-energy data. His definition of ionic strength (1921) allowed the systematization of activity data. Lewis set new standards of experimental accuracy and reliability; to a fertile but unorganized field he began to bring new clarity and order.
Lewis’ other theoretical interests also flourished at M.I.T. The publication of Einstein’s theory of relativity (1905) and his mass-energy equation renewed Lewis’ interest in his early speculations on radiation. He derived the mass-energy equation from his early idea of the pressure of light without using the principle of relativity (1908). This striking concurrence of his view with Einstein’s convinced Lewis of the value of his youthful ideas and made him one of the very few early supporters of Einstein and relativity in America. A second paper with Richard Tolman (1912), deriving Einstein’s equation from conservation laws and the principle of relativity, illustrates Lewis’ delight in the bizarre paradoxes of relativity theory that most people found so profoundly disturbing: Lewis was an iconoclast and reveled in the overthrow of long-established ideas.
In 1912 Lewis accepted an offer to become dean and chairman of the College of Chemistry at the University of California at Berkeley. The chemistry department was badly run-down, and Lewis was given generous financial support and a free hand to recruit new faculty and to initiate reforms. With him he brought William Bray, Richard Tolman, and Joel Hildebrand, all of whom became distinguished teachers and authors; their textbooks trained the first generation of American chemists who rivaled the products of German universities. The most advanced ideas of German physical chemistry—above all, thermodynamics—still neglected at most American universities, were familiar at Berkeley.
Lewis himself taught no courses—he was always uneasy speaking before a large group—but his influence was felt everywhere. The curriculum was reformed, and introductory courses became models of clarity. Lewis kept the teaching load light and encouraged original research among both his colleagues and his students. No one was permitted to become a narrow specialist, and speculation and free discussion were encouraged. Lewis presided over weekly research conferences where the latest topics were discussed with the utmost freedom by professors and students alike. Lewis had missed this progressive and students alike. Lewis had missed this progressive, cooperative “spirit of research” in his student days at Harvard, and its results can be seen in the reputation that his department quickly attained and in the roster of fine chemists it turned out. Lewis’ reform and modernization of chemical education set the standard for American chemistry and is one of his most important and enduring achievements.
At Berkeley, Lewis’ work on thermodynamics grew more intense. In a long paper of 1913 he summarized and brought up to date the theory and methods for calculating free-energy data. The great utility of a complete table of free energies for predicting chemical behavior, he asserted, made the collection of such data “an imperative duty of chemistry.” in the next seven years Lewis published a series of lengthy papers, many in collaboration with Merle Randall, systematically collecting and reworking all the known free-energy data for each element. Especially important is a compendium of entropy data (1917) and an empirical verification of Nernst’s third law. All this material became the body of his book on thermodynamics (1923). Next to Nernst and Fritz Haber, Lewis was probably the most important figure in chemical thermodynamics at that time; and his clear, systematic organization of data was second to none.
Equally fruitful was Lewis’ theoretical achievement in valence theory. In 1913 Bray and Lewis proposed a dualistic theory of valence which distinguished two distinctly different kinds of bond: the familiar polar bond formed by electron transfer, as in Na+ c1-, and a nonpolar bond that did not involve electron transfer. The polar theory, exemplified by J. J. Thomson’s popular book The Corpuscular Theory of Matter (1907), was then at the peak of its popularity. Bray and Lewis were the first to challenge the view that all bonds, even those in the inert hydrocarbons, were polar; and their heresy was not well received. But other dissenters soon appeared. In 1914 Thomson himself postulated a nonpolar bond involving two electrons and two tubes of force. In 1915 Lewis saw in manuscript a paper by Alfred Parson, an English graduate student visiting Berkeley for a year, that postulated a two-electron nonpolar bond and also a cubic octet very similar to Lewis’ cubic atom. This striking coincidence apparently revived Lewis’ interest in his early speculations on atomic structure and valence. Lewis probably derived the shared electron pair bond from a combination of the novel and suggestive theories of Parson and Thomson with his own model of the cubic atom.
Early in 1916 Lewis published his germinal paper proposing that the chemical bond was a pair of electrons shared or held jointly by two atoms. The cubic atom was an integral part of Lewis’ theory of molecular structure. In terms of cubic atoms, the single bond was represented by two cubes with a shared edge, or more simply by double dots (Figure 1), a convention that has been universally adopted.
According to the octet rule, for a molecule to be stable, each atom must be surrounded by four pairs of electrons that are either shared or free pairs. From this simple idea Lewis derived structures for the halogen molecules, the ammonium ion, and the oxyacids, all of which had proved insoluble for previous theories of valence. Lewis conceived polar bonds simply as unequally shared electron pairs. Since complete transfer of electrons was only the extreme case of polarity, Lewis abandoned his earlier dualistic view. The polar theory became a special case of Lewis’ more general and unified theory.
Lewis’ theory of the shared-pair bond received no notice in 1916, and he followed it up with only one further paper on color in molecules with “odd” or unpaired electrons (1916). His cubic or “static” atom appeared to be inconsistent with the physicists’ view of the atom, based mainly on spectroscopy, which demanded moving electrons, as in Bohr’s planetary model (1913). The physicists’ “dynamic” atom, however, failed to explain the rigid stereochemistry of carbon compounds. This apparent paradox was much disputed in 1919-1923, and Lewis vigorously defended the static atom against Bohr’s atom in a lecture to the Physical Society in 1916 and in his book Valence (1923).
By 1916 World War I had halted scientific work in Europe, and in January 1918 Lewis went to France as a major in the Chemical Warfare Service. There he organized the Gas Defense School to train gas officers, and proved so excellent an organizer that he was decorated for his service upon his return in September 1918.
In the spring of 1919 the cubic atom and the shared-pair bond were taken up by Irving Langmuir, who was already famous for his invention of the gas-filled electric lamp and his theory of surface absorption. His dramatic lecture on the new theory of the atom, delivered to the American Chemical Society in April 1919 and often repeated by request elsewhere, suddenly kindled the interest of American chemists. In a series of long papers and lectures in 1919-1921 Langmuir elaborated Lewis’ theory so successfully that the Lewis-Langmuir theory, or the Langmuir theory as it was known to many, was talked of everywhere and soon was widely accepted. Lewis was resentful that Langmuir received so much of the credit for the ideas that he had originated. Langmuir always acknowledged his debt to Lewis but felt that he had added enough on his own to warrant the compound name. An exchange of polite but outspoken letters in 1919-1920 probably did not clear the air, but some years later Lewis and Langmuir were again on friendly terms.
Langmuir abruptly ceased publishing on valence in 1921, probably realizing the superiority of the increasingly sophisticated Bohr theory. Lewis, however, continued to support the static atom in a lecture to the Faraday Society in 1923 and in his Valence and the Structure of Atoms and Molecules (1923). The conflict between the static and dynamic atoms soon disappeared with the introduction of directed orbitals, and the cubic atom quickly became obsolete. But the shared-pair bond proved to be one of the most fruitful ideas in the history of chemistry. Valence became the textbook of the first generation of chemists for whom the chemical bond was more than a simple line. For the first time mechanisms of complex organic reactions could be explained in terms of shifting electron pairs; and in England a new school of physical-organic chemistry was formed by A. Lapworth, M. Lowry, R. Robinson, and others. In the late 1920’s the shared-pair bond was the starting point for the new quantum chemistry of E. Schrödinger, H. London, L. Pauling, and others, which transformed Lewis’ germinal idea into a quantum mechanical theory of molecular structure.
In 1923 Lewis also published, with Merle Randall, Thermodynamics and the Free Energy of Chemical Substances. This extremely influential textbook was for several generations the clearest and simplest presentation of chemical thermodynamics. Its summaries of reliable free-energy data made readily accessible to chemists, even to the novice, the powerful tools of thermodynamics. Once the luxury of specialists, after 1923 thermodynamics was increasingly regarded as an indispensable part of chemical education and research.
For Lewis 1923 marked an end to two of his most abiding interests. He had had enough of collecting free-energy data; and in organic and quantum chemistry, where the shared-pair theory proved most fruitful, Lewis was not at home. He thus found himself at loose ends; and for the next ten years he occupied himself almost exclusively with his third early interest, the theory of quantum and radiation by He tried to derive the laws of quantum radiation by thermodynamic reasoning from the law of microscopical reversibility. He proposed that light does not emit to all space but only to a receiver and that in the space-time manifold, emitter and receiver are in “virtual contact.” Such entrancing paradoxes of space and time formed the bulk of his Silliman lectures of 1925 and the resulting book, The Anatomy of Science (1926). The decade 1923-1933 was certainly the least successful period of Lewis’ career. He was a good enough mathematician to follow the contemporary developments in relativity and quantum theory, but he remained an amateur and an outsider. His novel ideas, stemming from youthful inspiration and his bent for thermodynamic reasoning, remained out of touch with the best professional thought. Lewis styled himself an enfant terrible and enjoyed shocking people with his unorthodox views. The profound revolution in physics in the 1920’s encouraged his taste for paradoxes; it was clear that some cherished beliefs would be brought down. Anything seemed possible, and it was difficult to distinguish brilliant ideas from absurd ones. A letter from Einstein to Lewis suggests that even he took Lewis’ ideas seriously; but in retrospect is clear that Lewis was out of his depth.
Lewis probably sensed this, for in 1933 he abruptly abandoned theorizing to exploit an unexpected opportunity in a field quite new to him, the separation of isotopes. Deuterium had been discovered in 1932 by Harold Urey, who noted that it might be isolated on a large scale by fractional electrolysis of water. Lewis had been trying to separate oxygen isotopes when he realized that deuterium, being twice the size of ordinary hydrogen, would be easier and more interesting to obtain. In 1933 he succeeded in obtaining nearly pure heavy water and in the next two years rushed out, against intense competition, twenty-eight reports on deuterium chemistry, including several in collaboration with E. O. Lawrence on the nuclear reactions of deuterium in the cyclotron. Since deuterium was markedly different from hydrogen, Lewis foresaw a whole new chemistry of deutero compounds with distinct and unusual properties. But by 1934 these high hopes had apparently paled, for Lewis abruptly ceased work on heavy water. Covalent carbondeuterium bonds are in fact not easy to make, and deutero compounds are not very different from ordinary compounds.
It is precisely these properties that make deuterium an ideal tracer for studying organic, and especially biochemical, reactions; but Lewis was not prepared to exploit this opportunity. He carried out several studies on the lethal effect of heavy water on germinating plant seeds and on living creatures, but he did not realize how deuterium could be used to study the microchemistry of living tissue. Lewis also tried to follow up the nuclear physics of deuterium in 1936-1937, but this attempt ended in failure when a report on the refraction of neutrons by wax had to be withdrawn as an experimental error. Again Lewis was not at home in either of the fields opened up by his own work.
In 1938 Lewis finally hit upon a fruitful combination of theory and experiment in photochemistry. He had long been interested in the theory of colored compounds. In 1914 he identified two different forms of the indicator methylene blue, and in 1916 his idea of the shared-pair bond led him to propose that color was due to the presence of “odd” electrons. (His work on color had won him the Nichols Medal in 1921.) The occasion for his return to this subject was an important lecture on the theory of acids and bases at the Franklin Institute in 1938. According to his generalized theory of acids and bases proposed in 1923, bases were molecules having free electron pairs to donate, whereas acids were electron-poor molecules that could accept an extra pair. The idea of a general base had been widely accepted; but because of the prevailing proton theory of acidity, Lewis’ conception of acids had not. His 1938 lecture did much to make the Lewis acid an important part of chemical theory.
Many of the Lewis acids, such as the triphenylcarbonium salts, exist in a variety of different-colored forms. Lewis proposed that these forms were of two distinct “electromeric” types which differed only in the distribution of electrons in the molecule. Neutralization of one kind required energy of activation, while the other was neutralized spontaneously. Ironically, the concept of electromerism had been an important part of the abandoned electrochemical theory of valence, which Lewis’ shared-pair bond and Pauling’s theory of resonance hybrids had rendered obsolete. Despite its outmoded terminology, Lewis’ theory of electromeric states proved extremely fruitful. Photochemistry was already becoming a popular field in the 1930’s, but like thermodynamics in 1910 it lacked a solid theoretical foundation. In a long review in 1939 Lewis and Melvin Calvin summarized the known facts in terms of their elaborate theory of color. There followed a series of fine experimental papers on fluorescence and phosphorescence spectra and on photochemical reactions in rigid media.
Lewis separated two emission bands in phosphorescence spectra, which he associated with “electromeric” excited states. From one, emission was delayed, that is, it required energy of activation; from the other, emission was instantaneous. At first Lewis had consciously avoided quantum mechanical interpretations, preferring a more classical chemical approach. But by 1943 his kinetic and optical studies of phosphorescence had led him to believe that the two electromeric states were in fact the singlet and triplet states of quantum theory. This striking conclusion, which was confirmed in several important papers in 1944-1945 with Calvin and Michael Kasha, was the starting point for the rapid development of photochemistry.
These were Lewis’ last papers. He died in 1946 while carrying out an experiment on fluorescence. Thus at the end of his career Lewis again found a field that combined theoretical interest and practical opportunities; to photochemistry he was able to bring the same rigor of experiment and vigor of imagination that characterized the best work of his early years.
I. Original Works. Lewis’ principal writings are “Outlines of a New System of Thermodynamic Chemistry,” in Proceedings of the American Academy of Arts and Sciences,43 (1907), 259-293; “The Atom and the Molecule,” in Journal of the American Chemical Society,38 (1916), 762-785; Thermodynamics and the Free Energy of Chemical Substances (New York, 1923), written with M. Randall; Valence and the Structure of Atoms and Molecules (New York, 1923; repr. New York, 1965); “The Isotope of Hydrogen,” in Journal of the American Chemical Society,55 (1933), 1297-1300; “Acids and Bases,” in Journal of the Franklin Institute,226 (1938), 293-318; and “Phosphorescence and the Triplet State,” in Journal of the American Chemical Society,66 (1944), 2100-2109, written with M. Kasha.
Two boxes of Lewis’ correspondence and some MS notes from his student years are in the Lewis Archive, University of California, Berkeley. The Harvard Archives contain Lewis’ Harvard records and a brief correspondence with T. W. Richards.
II. Secondary Literature. J. H. Hildebrand, “Gilbert N. Lewis,” Biographical Memoirs. National Academy of Sciences, 31 (1958), 209-235, includes a complete bibliography. See also W. F. Giauque, “Gilbert N. Lewis,” in Yearbook. American Philosophical Society for 1946, pp. 317-322; A. Lachmann, Borderland of the Unknown (New York, 1956); and R. E. Kohler, “The Origin of G. N. Lewis’s Theory of the Shared Pair Bond,” in Historical Studies in the Physical Sciences, 3 (1971), 343-376. The relation between Lewis and Langmuir is discussed in R. E. Kohler, “Irving Langmuir and the Octet Theory of Valence,” ibid., 4 (1972).
R. E. Kohler
Lewis, Gilbert N
Lewis, Gilbert N.
AMERICAN PHYSICAL CHEMIST
Gilbert Newton Lewis was born on October 25, 1875, in West Newton, Massachusetts. A precocious child, he received his early education at home and learned to read by the age of three. When Lewis was nine, his family moved to Lincoln, Nebraska. He attended the University of Nebraska for two years and in 1893 transferred to Harvard University, from which he received his B.S. in 1896.
After a brief stint as a teacher at Phillips Academy in Andover, Massachusetts, Lewis returned to Harvard, where he obtained his M.A. in 1898 and Ph.D. in 1899. He subsequently studied at the universities at Göttingen and Leipzig in Germany (1900–1901) and then returned to Harvard as an instructor (1901–1906). In 1907 Lewis became an assistant professor at the Massachusetts Institute of Technology, where he soon rose to the rank of full professor.
In 1912 Lewis accepted a position as dean and chairman of the College of Chemistry at the University of California, Berkeley. He remained at Berkeley for the rest of his life and transformed the chemistry department there into a world-class center for research and teaching. His reforms in the way chemistry was taught, a catalyst for the modernization of chemical education, were widely adopted throughout the United States. Lewis introduced thermodynamics to the curriculum, and his book on the same subject became a classic. He also brought to the study of physical chemistry such concepts as fugacity, activity and the activity coefficient, and ionic strength.
At the beginning of the twentieth century physicists tried to relate the electronic structure of atoms to two basic chemical phenomena: the chemical bond (the attraction between atoms in a molecule) and valence (the quality that determines the number of atoms and groups with which any single atom or group will unite chemically and also expresses this ability to combine relative to the hydrogen atom). German chemist Richard Abegg was the first to recognize in print the stability of the group of eight electrons, the arrangement of outer electrons that occurs in noble gases and is often attained when atoms lose or gain electrons to form ions. Lewis called this the "group of eight," and American chemist and physicist Irving Langmuir labeled it an "octet."
In 1902, while explaining the laws of valence to his students at Harvard, Lewis conceived a concrete model for this process, something Abegg had not done. He proposed that atoms were composed of a concentric series of cubes with electrons at each of the resulting eight corners. This "cubic atom" explained the cycle of eight elements in the Periodic Table and corresponded to the idea that chemical bonds were formed by the transfer of electrons so each atom had a complete set of eight electrons. Lewis did not publish his theory, but fourteen years later it became an important part of his theory on the shared electron-pair bond.
In 1913 Lewis and Berkeley colleague William C. Bray proposed a theory of valence that differentiated two different types of bond: a polar bond formed by the transfer of electrons and a nonpolar bond not involving electron transfer. In 1916 Lewis published his seminal article suggesting that the chemical bond is a pair of electrons shared or held jointly by two atoms. He depicted a single bond by two cubes sharing an edge, or more simply by double dots in what has become known as Lewis dot structure.
According to Lewis's octet rule, each atom should be surrounded by four pairs of electrons, either shared or free pairs. Lewis derived structures for halogen molecules, the ammonium ion, and oxy acids, inexplicable according to previous valence theories. He viewed polar bonds as unequally shared electron pairs. Because the complete transfer of electrons was only an extreme case of polarity, he abandoned his earlier dualistic view; the polar theory was just a special case of his more general theory.
Lewis's shared electron-pair theory languished until Langmuir revived and elaborated it beginning in 1919. It was soon accepted as the Lewis–Langmuir theory, one of the most fundamental concepts in the history of chemistry.
Lewis's acid-base concept is also well known to introductory-level chemistry students. A Lewis acid, for example, BF3, AlCl3, or SO3, is a substance that can accept a pair of electrons from a Lewis base, for example, NH3 or OH−, which is a substance that can donate a pair of electrons. It can be applied to various areas, for example, coordination chemistry : The metal ion is a Lewis acid, the ligand is a Lewis base, and the resulting formation of a coordinate covalent bond corresponds to a Lewis acid–base reaction.
Lewis made additional valuable contributions to the theory of colored substances, radiation, relativity, the separation of isotopes , heavy water, photochemistry, phosphorescence, and fluorescence. As a major in the U.S. Army Chemical Warfare Service during World War I, he worked on defense systems against poison gases. From 1922 to 1935 he was nominated numerous times for the Nobel Prize in chemistry. Lewis's death, while measuring the dielectric constant of hydrogen cyanide on March 23, 1946, precluded his receiving the prize, which is not awarded posthumously.
see also Acid-Base Chemistry; Lewis Structures.
George B. Kauffman
"Gilbert Newton Lewis: 1875–1946." Papers presented at the 183rd National Meeting of the American Chemical Society, Las Vegas, NV. Journal of Chemical Education 61: (January 1984) 3–21, (February 1984) 93–116, (March 1984) 185–215.
Hildebrand, Joel H. (1958). "Gilbert N. Lewis." Biographical Memoirs, National Academy of Sciences 31:209–235.
Leicester, Henry M., ed. (1968). Source Book in Chemistry 1900–1950, pp. 100–106. Cambridge, MA: Harvard University Press.
Lewis, Edward S. (1998). A Biography of Distinguished Scientist Gilbert Newton Lewis. Lewiston, NY: Edwin Mellen Press.
Lewis, Gilbert N. (1916). "The Atom and the Molecule." Journal of the American Chemical Society 38:762–785.
Lewis, Gilbert Newton, and Randall, Merle F. (1923). Thermodynamics and the Free Energy of Chemical Substances. New York: McGraw-Hill.
Gilbert Newton Lewis
Gilbert Newton Lewis
Gilbert Newton Lewis (1875-1946) was an American physical chemist whose concept of electron pairs led to modern theories of chemical bonding. His concept of acids and bases was another fundamental contribution.
Gilbert N. Lewis was born at Weymouth, Mass., on Oct. 23, 1875. He received his bachelor's degree in 1896 and his doctorate in 1899 from Harvard University and then served as instructor in chemistry at Harvard until 1900. After a year in Leipzig, Germany, he was in charge of the laboratories of the U.S. Bureau of Weights and Measures in the Philippine Islands in 1904-1905. He became assistant professor of physiochemical research at the Massachusetts Institute of Technology in 1907 and full professor in 1911. He married Mary H. Sheldon in 1912, and they had three children. Also in 1912 he accepted the chairmanship of a small chemistry department at the University of California at Berkeley, where he remained until his death.
In 1916 Lewis published his famous paper "The Atom and the Molecule," in which he proposed that nonionic molecular compounds were the result of the sharing of electrons among atoms. He suggested that a chemical bond was produced in the formation of a molecular compound. This involved the sharing of a pair of electrons by two atoms. He called this a covalent bond, and it became the basis of the electronic theory of the chemical bond.
Lewis made another important scientific observation in 1916, when he propounded the electron-pair concept of acids and bases, in which acids were classified more generally as electron-pair acceptors, and bases as electron-pair donors. This theory was useful in explaining many reactions otherwise difficult to classify. According to this theory, not only proton-donating compounds are classified as acids. Any compound or ion capable of accepting a pair of electrons to form a new compound is considered to be an acid. In 1923 he published Valence and the Structure of Atoms and Molecules. Three years later he wrote The Anatomy of Science.
At Berkeley, Lewis gradually built one of the most powerful and creative chemistry departments in the world. His lectures in thermodynamics drew students from all over the world, many of whom became famous. Among these were Linus Pauling, Harold Urey, Melvin Calvin, and William Giauque, each of whom received the Nobel Prize in chemistry. Many scientists believe that Lewis, who received a large number of science's most prestigious honors, should have become a Nobel laureate in chemistry, but this prize eluded him. He died on March 23, 1946.
A good account of Lewis and his work is in Great American Scientists, by the editors of Fortune (1961). His major contributions to chemistry are explained on a simple level in Gregory R. Choppin and Bernard Jaffe, Chemistry: Science of Matter, Energy and Change (1965).
Lachman, Arthur, Borderland of the unknown; the life story of Gilbert Newton Lewis, one of the world's great scientist, New York: Pageant Press, 1955. □