Pauling Develops His Theory of the Chemical Bond

Through his ingenious use of quantum mechanics, Linus Pauling developed a theory of the chemical bond that strongly influenced the fields of chemistry, physics, biology, and mineralogy.

Summary of Event

When interviewers asked Linus Pauling to name what he considered to be his most important discovery, he often responded that he was extremely pleased by his work on the nature of the chemical bond. His series of papers on this topic in the early 1930’s and his book The Nature of the Chemical Bond and the Structure of Molecules and Crystals (1939) Nature of the Chemical Bond and the Structure of Molecules and Crystals, The (Pauling) clarified large areas of chemistry and contributed to important advances in such fields as biochemistry, mineralogy, and medicine. Pauling’s work grew out of the contributions of his predecessors, and in several ways he was able to recapitulate the thinking of scientists who had lived many years before him. The idea of an atom’s definite combining power had been developed in the nineteenth century, and it formed the theoretical backbone of a dynamic structural theory that was able, under the insightful tutelage of the Dutch physical chemist Jacobus Henricus van’t Hoff Van’t Hoff, Jacobus Henricus (1852-1911), to reveal how the four valence bonds of certain carbon compounds are directed toward the corners of a regular tetrahedron, thus accounting for the right- and left-handed forms of particular organic molecules. [kw]Pauling Develops His Theory of the Chemical Bond (1930-1931)
[kw]Chemical Bond, Pauling Develops His Theory of the (1930-1931)
[kw]Bond, Pauling Develops His Theory of the Chemical (1930-1931)
Chemical bond theory
Chemistry;chemical bond theory
[g]United States;1930-1931: Pauling Develops His Theory of the Chemical Bond[07500]
[c]Chemistry;1930-1931: Pauling Develops His Theory of the Chemical Bond[07500]
[c]Physics;1930-1931: Pauling Develops His Theory of the Chemical Bond[07500]
Pauling, Linus
Lewis, Gilbert N.
Slater, John C.
Heitler, Walter Heinrich
London, Fritz Wolfgang

Linus Pauling.

(The Nobel Foundation)

With the discovery of the electron by Sir Joseph John Thomson in 1897 and the discovery of the nucleus by Ernest Rutherford in 1911, theories about how atoms are linked together in compounds became more dependent on a physical model of the atom, especially how electrons are arranged around the nucleus. Pauling became fascinated by the ideas of Gilbert N. Lewis, who had proposed a theory of the inner structure of the atom. Atoms;structure In Lewis’s model, electrons occupy the space around the nucleus in certain patterns that are responsible for an element’s specific properties, including its bonding properties. Lewis pictured the atom as a series of concentric cubes whose corners could be occupied by electrons. He used this “cubical atom” to explain the formation of chemical bonds by the transfer of electrons from one atom to another. Lewis proposed that electrons could be shared between atoms in forming a covalent bond. The idea of this covalent (or shared electron-pair) bond had a great impact on Pauling.

As a graduate student at the California Institute of Technology, Pauling made use of Niels Bohr’s Bohr, Niels
Bohr model of the atom quantum theory of the atom in his own theoretical speculations. Bohr had been successful in applying the quantum idea to the hydrogen atom, showing how its lone electron could exist solely in specific orbits (or quantized states) around its nuclear proton. Bohr’s theory, however, was not very successful in accounting for the electronic structures of complex atoms or for the combinations of atoms in molecules. Pauling tried to use Arnold Sommerfeld’s Sommerfeld, Arnold version (with elliptical rather than circular orbits) of Bohr’s atomic model to explain the structure and properties of benzene, a ring compound of six carbon atoms, but the old quantum theory was simply too weak to enable Pauling or any other scientist to develop a satisfactory theory of chemical structures.

With the development of quantum mechanics Quantum mechanics in the mid-1920’s, a tool was now available to resolve many of the deficiencies of the old theory. Pauling was influenced principally by Erwin Schrödinger’s Schrödinger, Erwin formulation of wave mechanics Wave mechanics and was able to use the new quantum mechanics to predict the properties of various ions and to develop a perspicacious picture of the ionic bond (an interatomic link in which one of the atoms has a disproportionate share of the electronic charge). One of the ideas on which wave mechanics was built is the wave nature of the electron. Using his wave equation, Schrödinger was able to reveal many advantages in treating the electron as a wave rather than as a particle in a strictly defined orbit. In one interpretation of Schrödinger’s equation, electrons occupy orbitals, often represented by such shaded geometric figures as spheres or somewhat distorted dumbbells, the intensity of the shading corresponding to the relative probability of finding an electron in this particular region. Although Schrödinger’s treatment made it difficult to speak of atoms as rigid mechanical structures, it permitted a new view of chemical bonding, particularly in the hands of Walter Heinrich Heitler and Fritz Wolfgang London.

While in Munich, Pauling met Heitler, who was working toward his doctoral degree. In the summer of 1927, Pauling discussed quantum mechanics with Heitler and met London, who had a Rockefeller Foundation grant to work with Schrödinger. After Pauling returned to the United States, Heitler and London published a paper on the quantum mechanics of the chemical bond that Pauling once called “the greatest single contribution to the clarification of the chemist’s concept of valence.” In the paper, Heitler and London made use of the idea of resonance, which had been introduced by Werner Heisenberg in 1926, to explain how two hydrogen atoms could bond when brought near to each other. In the resonance phenomenon, an interchange in position of the two electrons reduces the system’s energy and causes the formation of a bond. Heitler and London supplied a quantum mechanical justification for Lewis’s electron-pair idea. Their quantum mechanical method allowed them to calculate approximate values for various properties of the hydrogen molecule, for example, how much energy it would take to split the molecule into its component atoms.

In the fall of 1927, Pauling began a period of intense scientific creativity. He developed a theory for predicting the structures of crystals. He refined his wave mechanical treatment of the hydrogen molecule-ion (a hydrogen molecule with one rather than two electrons). Most important, he enunciated for the first time what later came to be called the hybridization of atomic orbitals. The idea of hybrid orbitals grew out of his chemical intuition. To physicists, it seemed strange that carbon, with two different types of orbitals (the spherical 2s and the dumbbell-shaped 2p), should generate, in its chemical compounds, four identical bonds that were directed to the corners of a tetrahedron. To Pauling, it seemed strange that such tetrahedral compounds of carbon as methane (CH
) could not be explained by atomic orbitals. A physicist would regard it as a very serious error to mix s and p orbitals, but Pauling recognized that the energy separation between the two orbital states was small, compared with the energy of the bond formed. For him, the ability to make the best possible bond was the most important consideration. In 1928, he published a short paper in which he reported that he had used quantum mechanical resonance to derive the four equivalent orbitals used in bonding by the carbon atom. Furthermore, these orbitals, which he also referred to as hybrid orbitals, are directed toward the corners of a regular tetrahedron.

Pauling had difficulties in translating his insight into a mathematical treatment that would be convincing to most scientists. When he was unable to solve the complex mathematical expressions, he set the matter aside. Meanwhile, in 1929, John C. Slater became interested in the quantum mechanics of the chemical bond. He had developed a method for interpreting the complex spectra of certain atoms, and when he applied his methods to molecules, he derived information about the valence and directional properties of these molecules. He first described his results in an informal talk at the Washington meeting of the American Physical Society in April, 1930. He published a paper in 1931 in the Physical Review, suggesting that the dumbbell-shaped charge clouds of certain orbitals were responsible for the directional properties of many chemical compounds. He also introduced the criterion of maximum overlapping of orbitals for bond strength, and he made extensive use of resonance. Further, he discussed the directional bonds formed in molecules with many atoms, and he even attempted to explain why the four valences of carbon have tetrahedral symmetry.

Slater’s work stimulated Pauling to return to the problem he had set aside in 1928. In December, 1930, while he was doing some calculations, he had an idea about making an assumption to simplify the quantum mechanical equations describing the bonding orbitals of carbon so that they could be solved in an approximate way. The quantum mechanical equations describing the orbitals with which he was concerned have radial and angular parts, and given that the radial part of the 2s wave function of the carbon atom (from which the orbital can be readily derived) is not very different from the radial part of the three 2p functions, Pauling concluded that little error would be introduced if he ignored the radial factor in the p function. This facilitated his calculations of various hybrid orbitals, for example, the tetrahedral hybrids, called sp3 because they involve a combination of the 2s and the three 2p orbitals. His semiquantitative approach to the chemical bond proved so successful that he went on to develop the implications of his ideas. He was able to give explanations from the kinked structure of the water molecule to the transition from covalent to ionic bonding.

In January and February, 1931, Pauling continued to make refinements and extend their application. His most famous paper, titled “The Nature of the Chemical Bond: Applications of Results Obtained from the Quantum Mechanics and from a Theory of Paramagnetic Susceptibility to the Structure of Molecules,” appeared in the Journal of the American Chemical Society on April 6, 1931. With this work Pauling established a framework for understanding the electronic and geometric structures of molecules and ions in terms of hybrid bond orbitals. Also, the framework permitted him to use the magnetic properties of molecules to distinguish between ionic and covalent bonding. Pauling explained how the energy barrier between s and p orbitals can be broken when strong bonds are formed. Because a large orbital overlap is associated with the formation of a strong bond, he was able to relate bond strength to the nature of the orbitals from which the bond is formed. He used the ideas of resonance and hybridization to explain the tetrahedral, square, and octahedral configurations of certain molecules.

From 1931 to 1933, Pauling published a series of papers on the nature of the chemical bond, in which he used the concepts of hybridization of bond orbitals and the resonance of molecules among two or more alternative structures to elucidate many basic chemical phenomena. For example, he was able to use hybrid orbitals and resonance to explain the properties of the benzene molecule, the carbon-to-carbon bond lengths, planar structure, hexagonal symmetry, and great stability of which had puzzled chemists for generations. He was able to formulate an electronegativity scale of the elements, which proved to be extremely useful to chemists. Electronegativity is the tendency of an element in a compound to attract electrons to itself, and Pauling found a way to assign numbers to the elements to represent this power of attracting electrons.

Pauling’s papers were so successful that, by 1935, he believed that he had a complete understanding of the nature of the chemical bond. He was able to give form to this understanding in the late 1930’s, when he served as a professor of chemistry at Cornell University. Out of his lectures came a book, The Nature of the Chemical Bond and the Structure of Molecules and Crystals, the first edition of which was published in 1939. The book was received enthusiastically, with one reviewer calling it “epoch-making.” It was an excellent summation of Pauling’s decade of studies on the nature of the chemical bond.


Pauling’s theory of the chemical bond was a milestone in the development of modern chemistry. Chemists were deeply impressed by the power of Pauling’s methods to unify a vast amount of previously unintegrated data. They were also impressed by the large amount of experimental data Pauling had collected, tabulated, and rationalized. Pauling’s papers and book on the chemical bond accumulated thousands of citations in the 1930’s and 1940’s. Physicists, biochemists, molecular biologists, and mineralogists used his data and ideas to solve many problems in their disciplines. It is interesting to note that James D. Watson Watson, James D. and Francis Crick Crick, Francis made extensive use of Pauling’s book while they were trying to figure out the structure of deoxyribonucleic acid (DNA). DNA

Despite the great success of Pauling’s work on the chemical bond, it shared the contingencies of all scientific achievement. Because of the many approximations involved in his approach, it had a number of limitations, which critics soon pointed out. Charles A. Coulson, Coulson, Charles A. for example, asserted that although Pauling’s comparison of bond length and orbital overlap is very useful in a qualitative sense, it is not adequate for the quantitative determination of bond energies for certain hybrid orbitals. Nevil Sidgwick Sidgwick, Nevil also argued that Pauling’s theory of the chemical bond should be “received with a certain reserve” because of its lack of mathematical rigor. Other scientists developed a theory—often called the molecular orbital (MO) method or theory Molecular orbital theory —that rivaled Pauling’s. Pauling and Slater’s method was a natural outgrowth of the work of Heitler and London; therefore, this approach, which viewed the chemical bond from the perspective of two separate atoms coming together, was sometimes called the Heitler-London-Slater-Pauling (HLSP) method, but it is known most often as the valence bond (VB) method or theory. Valence bond theory Other names given to the theory, the “method of localized pairs” and the “method of the directed valence bond,” are not used often.

In the molecular orbital method, the problem of the chemical bond is attacked from the viewpoint of the single molecule produced by the coalescence of the atoms. During the late 1940’s and through the 1950’s, the MO method began to find favor with a number of chemists, and Pauling’s VB theory came increasingly under attack. Some European and American scientists pointed out that the VB theory was encountering difficulties in explaining the excited states of molecules, whereas the MO method was successful in its quantitative discussions of these excited states. During this same period, Pauling’s theory came under attack on ideological grounds. Certain Soviet critics found Pauling’s notion of resonance irrational, because it went against some of the basic tenets of dialectical materialism. Essentially, certain Soviet chemists criticized Pauling, whom they called a “decadent bourgeois scientist,” for his attributing reality to his formal equations and models (for these Soviet critics, resonance structures were human-made, not real). In place of Pauling’s theory, the Soviet chemists preferred to turn to the Soviet chemist Aleksandr Mikhailovich Butlerov, Butlerov, Aleksandr Mikhailovich whose theory of the mutual influence of atoms in molecules they found both scientifically and ideologically acceptable. It is ironic that Pauling, who shared with many Soviet philosophers and scientists a deterministic and realistic interpretation of quantum mechanics, should have encountered such animosity toward his theory of resonance, especially given that he admitted that his resonance descriptions bore a close resemblance to the actual molecular structures. A further irony is that at the same time Soviet chemists were attacking him for his resonance theory, American politicians were criticizing him for his attitude toward the Soviet Union.

Despite these criticisms from the East and the West, Pauling’s theory of the chemical bond proved to be useful for scientists for three decades. When the VB theory began to be replaced by the MO theory in the 1960’s, Pauling continued to retain a deep loyalty to his views of the chemical bond. This was characteristic of his approach to science. When he was convinced of the value of a scientific idea, he clung to it tenaciously and used it boldly. Resonance was such an idea. Early in his career, Pauling became convinced of its power and efficacy. Continued success in applying this idea to a great variety of chemical problems confirmed this attitude. Because of the strength of the resonance concept, he was less attracted to the new ideas about chemical bonding that came on the scene. Confronted with new problems, he tended to solve them in terms that he understood. This tenacity, which had a negative side, was also a part of his genius. Without his championing of the VB theory, modern structural chemistry’s great successes, such as the determinations of the structures of DNA and proteins, would have been significantly delayed. It is thus impossible to understand the evolution of modern structural chemistry without taking account of Pauling’s theories. Chemical bond theory
Chemistry;chemical bond theory

Further Reading

  • Coulson, Charles A. Valence. 3d ed. London: Oxford University Press, 1985. Volume by an Oxford professor of mathematics who became a distinguished theoretical chemist discusses how wave mechanics transformed valence theory. Includes some use of mathematical formulas, but the treatment is mostly qualitative. Intended for “the novice chemist with few mathematical attainments.” Features author, substance, and subject indexes.
  • Hager, Thomas. Force of Nature: The Life of Linus Pauling. New York: Simon & Schuster, 1995. In-depth biography places Pauling’s scientific work within the context of his long life and many wide-ranging interests.
  • Lagowski, J. J. The Chemical Bond. Boston: Houghton Mifflin, 1966. Uses the words of scientists to re-create for students the exciting process of discovery. Approach is historical and accessible to both students of science and readers taking chemistry for the first time. Contains several helpful tables and diagrams and includes lists of suggested readings.
  • Lewis, Gilbert Newton. Valence and the Structure of Atoms and Molecules. 1923. Reprint. New York: Dover, 1966. Although written before the development of quantum mechanics, this interesting account is still valuable for the light it sheds on the ideas of Lewis and other scientists. Includes an introduction by Kenneth Pitzer, illustrations, references, and index.
  • Mead, Clifford, and Thomas Hager, eds. Linus Pauling: Scientist and Peacemaker. Corvallis: Oregon State University Press, 2001. Collection of essays both by Pauling and about him is divided into sections on the man, the science, and Pauling’s work as a peace activist. Includes photographs, reproductions of original Pauling manuscripts, and a bibliography of Pauling’s works.
  • Palmer, William G. A History of the Concept of Valency to 1930. Cambridge, England: Cambridge University Press, 1965. Uses the theme of the increasingly fertile cooperation between physics and chemistry to explain how the idea of chemical valence was established and how the modern theory of the chemical bond evolved from it. Valuable source of information on both the early history of valence theories and the later development of electronic theories.
  • Pauling, Linus. The Architecture of Molecules. San Francisco: W. H. Freeman, 1964. Demonstrates Pauling’s excellent ability to communicate clearly his understanding of molecules. Illustrator Roger Hayward captures Pauling’s structural imagination in beautifully drawn and colored plates. Intended for young people who are beginning to develop an interest in science, but valuable for anyone with a curiosity about the natural world. Includes a periodic table of the elements and some tables of atomic radii.
  • _______. The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry. 3d ed. Ithaca, N.Y.: Cornell University Press, 1960. Pauling’s magnum opus. Because his approach is structural rather than mathematical, the book is accessible to readers with modest backgrounds in chemistry, physics, and mathematics. More difficult material appears in twelve appendixes. Includes author and subject indexes.
  • Russell, C. A. The History of Valency. Leicester, England: Leicester University Press, 1971. Detailed, coherently organized study of the history of physical and chemical bonding ideas. Treats controversial material in a balanced and insightful way. Includes indexes.
  • Servos, John W. Physical Chemistry from Ostwald to Pauling: The Making of a Science in America. Princeton, N.J.: Princeton University Press, 1990. Explores the evolution of physical chemistry in the United States through an analysis of the key institutions and scientists who made the discipline into a fertile source of innovative ideas. Includes endnotes and index.

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