by Thomas Hager
It looked at first as though Slater had beaten Pauling. But after reading it several times, Pauling found some important differences in their work. Slater's paper, for one thing, was more descriptive than quantitative; it did not provide a way to get hard numbers for bond strengths and lengths. Pauling dashed off a note to the Physical Review calling readers' attention to his JACS paper's "very simple but powerful approximate quantitative treatment of bond strengths" (Pauling's italics), briefly sketching his six rules, and stressing that it was he who had first put forward the quantum-mechanical approach to tetrahedral binding in 1928. (Slater had not referred to Pauling's earlier paper in his own work.) Pauling then quickly reviewed Slater's work and pointed out ways in which his own ideas went further.
The jockeying for position was delicately done: Pauling and Slater liked each other and respected each other's work. A few weeks before either paper came out, Slater offered Pauling a full professorship at MIT in physics, chemistry, or any combination of the two; in Pauling's "thanks but no thanks" reply (he had, by that time, been promised his own full professorship at Caltech), he wrote Slater, "There is no theoretical physicist whose work interests me more than yours." After reading Pauling's JACS paper, Slater wrote, "I am glad things worked out as they did, we both deciding simultaneously to write up our ideas. I haven't had a chance to read yours in detail yet, but it looked at first sight as if we were in good agreement in general." Their agreement was so good, in fact, that before Slater came to speak at a chemical-bond symposium Pauling arranged in Pasadena for the summer of 1931, he cautioned Pauling, "... our general points of view seem so similar that we shall want to compare notes before the meeting, to avoid saying the same things."
The two young men had hit upon the same approach to the same problems and would end up sharing credit for what would for a time be awkwardly called the Heitler-London-Slater-Pauling theory of chemical bonding—later, more gracefully, the valence-bond theory—with the consensus that Slater and Pauling had independently reached nearly the same conclusions at almost exactly the same time.
Quantum mechanics, as the physicist Victor Weisskopf later noted, had finally united the two great fields of physics and chemistry. By using the rules of the new physics to explain the bonding of atoms into molecules, Slater and Pauling consummated the marriage. The physicists had been right: Electrons moved in their orbits, they did not sit still on dry goods boxes. But in very important ways, the chemists, especially G. N. Lewis, had also been right. Electron orbitals were concentrated in certain directions, and bonds resulted from sharing electrons. Slater and Pauling showed how the physicists' new quantum principles resulted, logically and with at least approximate mathematical validity, in the molecules seen by chemists.
The importance and novelty of Pauling's work was underscored when his first paper was published just six weeks after being received by the JACS, a surprisingly short time compared to the normal waiting period of several months. Pauling was later told that the journal's editor, Arthur Lamb, could not think of anyone qualified to referee it and ran it without review.
- - -
To provide support and encouragement for chemists early in their careers, Dr. A. C. Langmuir (the brother of the physical chemist Irving Langmuir) in 1931 began funding an annual prize of one thousand dollars, to be awarded to the best young chemist in the nation. He left the selection to the American Chemical Society (ACS), the field's leading professional association. The first year the Langmuir Prize was offered, former ACS president A. A. Noyes made sure that his favorite young chemist was nominated, and in August 1931, Pauling was thrilled to learn that he had won. The prize recognized Pauling's unusual productivity and promise—at age thirty he had already published more than fifty papers covering a wide range of theoretical and experimental topics—and was a particularly rich award for those days, equivalent to roughly one-quarter of Pauling's annual salary. It also provided the ACS and Caltech with the opportunity for some publicity. Pauling soon found himself a minor celebrity, interviewed by newspaper writers from Portland to New York, asked for photographs, and featured in magazines. Scientific American ran a large picture of him looking serious and scholarly and described him as a "prodigy of American science." Noyes told reporters that Pauling was "the most promising young man with whom I have ever come in contact in my many years of teaching." A. C. Langmuir gushed that he was "a rising star, who may yet win the Nobel Prize." At an evening plenary session of the national ACS meeting in Buffalo, New York, that September, Pauling loped across the stage to receive the award from the society's president to the sound of enthusiastic applause from two thousand of the nation's leading chemists. A cartoon of the occasion, drawn for a meeting newsletter, showed a tousle-haired young Pauling eagerly stretching his hands out toward a bag marked "$1,000." His only regret, he said later, was that Belle had not been there to see him win the award.
Four years after taking his first faculty position, Pauling had risen from promising youngster to national prominence. By the end of 1931, he was a full professor, associate editor of the JACS, master of crystallography, co-solver of the great problem of the chemical bond, and winner of the Langmuir Prize. He had always had a core of self-confidence; now it blossomed and filled him. He would rarely doubt himself after 1931. "I might well have become egotistical as a result" of it all, he remembered. "And I think that earlier I had developed a feeling of self-confidence in regard to science. But... I think that I just said I shouldn't let this go to my head. I shouldn't think I'm really better than other people even though I do this one thing better than other people."
CHAPTER 7
Resonance
The Avant-Garde
In 1931, Albert Einstein heard Pauling deliver a talk on the chemical bond. Einstein was in Pasadena for several months being wooed for a faculty position at Caltech. Knowing that he had the world's greatest living scientist in his audience, Pauling worked especially hard to explain at length his new ideas about the application of wave mechanics to chemistry. Afterward, Einstein was asked by a reporter what he thought of the young chemist's lecture. He shrugged his shoulders, smiled, and said, "It was too complicated for me.”
Einstein may simply have been brushing off another newshound, but Pauling's interpretation of the chemical bond was complicated when carried out with any mathematical rigor—certainly too complicated for most chemists. Chemists were not prepared, historically, mathematically, or philosophically, for what Pauling offered them. Chemistry at that time was still a polyglot of separate disciplines and specialties rooted in the last century—organic chemistry, inorganic chemistry, physical chemistry, colloid chemistry, agricultural chemistry—each with its own champions and sets of puzzles to solve. There were ionists and thermodynamicists and now quantum chemists, separate tribes, each with its own traditions, methods, and journals, gathering together only at a few general meetings each year. In the United States only the Noyes-influenced departments at Caltech and Berkeley were able to effectively blur the old disciplinary lines with an emphasis on open communication among professors and a devotion to the general ideas underlying all of chemistry.
Pauling, with his avant-garde ideas about the quantum-based chemical bond, was in 1931 a decade ahead of his time. The vast majority of chemists neither knew what quantum mechanics was nor cared what it meant to their field. Real chemistry, to most of its practitioners, was something done in the laboratory, not on a piece of paper; discoveries were made through the hands-on experience of manipulating compounds and observing their reactions, not by dreaming up mathematical equations. X-ray crystallography was a physicists' tool they might have heard about but never used—Caltech was still one of the few places where the technique was applied in any significant way to chemistry. As for Pauling's emphasis on the importance of molecular structure, well, that was something that organic chemists thought was important, but not many other chemists believed that it played a significant role in chemistry.
What was important was gett
ing into the lab and getting your hands dirty. The laboratory chemists' disdain of theoreticians like Pauling, who looked too much to physics for their inspiration, was expressed by the leading British chemical educator Henry Armstrong in the mid-1930s: "The fact is, the physical chemists never use their eyes and are most lamentably lacking in chemical culture. It is essential to cast out from our midst, root and branch, this physical element and return to our laboratories."
Their distrust of abstract thinking exceeded only by their ignorance of the new physics, most early-1930s chemists responded to the Heitler-London-Slater-Pauling chemical-bond ideas with a yawn. As chemical Nobelist Harold Urey recalled, they were "completely bland about the matter, didn't understand it, and largely, except for Pauling, nobody paid any attention to it."
But A. A. Noyes and G. N. Lewis made sure that Pauling's ideas were given a wide hearing at Caltech and Berkeley, where most chemistry students developed an understanding that this approach was important and a handful of the brightest began to follow Pauling into the field. From the physics side, researchers like Slater and London continued to refine the mathematics of blended wave functions, working out the structure of simple molecules from physical first principles. Their impact on chemistry was muted, however, because they had not mastered the huge masses of empirical facts important to chemists, did not share the same worldview, and did not know which questions were important. They were not, in short, chemists.
One of the few who knew both physics and chemistry was Robert Mulliken. Mulliken, the son of an MIT chemist who learned physics with Millikan at the University of Chicago, roomed next to Slater while on a fellowship at Harvard and, like Pauling, made a pilgrimage to Europe in the late 1920s to learn quantum mechanics. At Goettingen, Mulliken had come under the influence of one of Born's assistants, Friedrich Hund, who was thinking through an approach to the chemical bond different from Pauling's. Hund was interested in molecular spectroscopy, the study of the characteristic light absorbed and emitted by molecules, and he found that viewed this way, molecules behaved in important ways like individual atoms. Hund and Mulliken came up with a concept of the chemical bond that seemed radically different from Pauling's. Instead of electrons concentrating between two nuclei to bind atoms together, Hund and Mulliken theorized that binding electrons were spread around the molecule's surface, forming what Mulliken would call molecular orbitals. They conceived of the hydrogen molecule, H2, for instance, not as two hydrogen atoms approaching each other and forming a bond by pairing their electrons, as Heitler and London had proposed, but as a two-electron helium atom splitting into two nuclei, with its surrounding electron cloud reshaping into a new molecular orbital. "In general no attempt is made to treat the molecule as consisting of atoms or ions," Mulliken wrote in 1932. "Attempts to regard a molecule as consisting of specific atoms or ionic units held together by discrete numbers of bonding electrons or electron pairs are considered as more or less meaningless." This was radical thinking; the molecular-orbital concept seemed diametrically opposed to everything chemists had thought about the nature of the chemical bond for decades. It did, however, fit the spectroscopic data, and Mulliken stuck with his ideas after returning to the United States to teach at the University of Chicago.
For a few years it looked as if chemists would be forced to choose either Pauling's view or Mulliken's. But at root the two approaches were not as different as they seemed. Both were based on Schroedinger's wave equation, and Slater and others found in the mid-1930s that if the mathematics was carried through far enough, the two approaches ended up providing the same results. It was rather like the choice physicists had to make between Heisenberg's matrix approach to quantum mechanics and Schroedinger's wave equation: Although seemingly very different, both were paths to the same destination. The choice of paths depended on which was easier to use and which worked better in a given situation.
Pauling, of course, thought his was the better approach to understanding the chemical bond. He understood that the molecular-orbital approach was useful—he had employed it in some cases while searching for a breakthrough on the chemical bond—but he largely dropped it when he found how to make his own variations on the Heitler-London theme work in 1931. Once Slater showed the essential equivalence of his and Mulliken's methods, Pauling saw no need to refer to the molecular-orbital approach. His ideas worked out of what chemists already believed about the chemical bond; Mulliken's were by comparison anti-intuitive and, Pauling thought, confusing to students.
And Pauling's notion of the chemical bond took off, while Mulliken's languished in relative obscurity. There were several reasons, prominent among them the fact that Pauling was an eloquent lecturer and a persuasive writer who knew how to communicate in language chemists could understand. When Pauling spoke, the valence-bond approach seemed like revealed wisdom. When Mulliken talked, people went to sleep. He was a terrible teacher, ill at ease in front of crowds, his voice almost inaudible. He refused to pander to his chemistry students, and his lectures were notoriously digressive, heavy with mathematics, and hard to follow. He was not much better in print. As the years went by, Mulliken and a small band of followers would continue to improve their molecular-orbital approach, refining the equations and using it to successfully attack a number of problems. Twenty years later, a new generation of chemists would come to prefer it over Pauling's approach. But in the 1930s, Mulliken's ideas would be lost in a blizzard of razzle-dazzle coming from Pasadena.
- - -
Pauling's taming of the wave equation in his first "Nature of the Chemical Bond" paper released a flood of new ideas. In June 1931 he submitted a follow-up paper, the second in what was to become a series, this one examining how quantum mechanics could explain the existence of relatively rare one-electron and three-electron bonds. His calculations helped distinguish among alternative explanations for the unusual bonding properties of oxygen, boron, and nitroso compounds, the rare "odd-electron" molecules that had so interested G. N. Lewis. Lewis himself helped talk Pauling through some of the ideas in this paper, the two of them scribbling sketches and formulae on a blackboard in Lewis's office during Pauling's teaching stints at Berkeley, the older man puffing out clouds of cigar smoke and advice.
Then it was on to bigger mysteries. One long-standing puzzle in chemistry concerned the relationship between two seemingly different types of bonds between atoms, ionic and covalent. In Lewis's scheme, the bond was covalent when two atoms shared a pair of electrons equally and ionic when one electron-hungry atom pulled the entire electron pair to itself, resulting in a net negative charge on one atom and a positive charge on the other; the resulting bond was then due to the electrostatic attraction between the two. The question was whether ionic and covalent bonds were separate species with a sharp dividing point or, as Lewis thought, points along a continuum.
In Pauling's third "Nature of the Chemical Bond" paper, he showed that quantum mechanics again supported Lewis. At least in some cases, his equations showed that the existence of "partial ionic" bonds, links that had both ionic and covalent characteristics, was compatible with both quantum mechanics and observed properties. In other cases he found that the jump between bond types could be discontinuous; it depended on how strongly the elements involved attracted the electrons. He backed up his arguments with a number of real-world examples and a set of conditions necessary for such intermediate bonds to form.
Pauling was now using the term "resonance" in place of "electron exchange" when writing about the chemical bond, and he was expanding the concept into new areas. Heisenberg had used the electron-exchange idea to account for the interchangeability of electrons; Heitler and London had used it to explain the covalent chemical bond; Pauling and Slater employed it to account for the energy needed to form hybrid bonds like those in the tetrahedral carbon atom. Now Pauling proposed that when certain criteria were met, resonance could exist between the ionic and covalent forms of a molecule. Hydrogen chloride, for example, could be viewed either as a hydroge
n atom linked to a chlorine atom through a purely covalent bond or as a positively charged hydrogen ion and a negatively charged chloride ion linked with a purely ionic bond. The actual molecule, Pauling proposed, is a sort of hybrid, a structure that resonates between the two alternative extremes. And whenever that happened, "whenever there is resonance between two forms," Pauling said, "the structure is stabilized."
For Pauling, the entire chemical landscape now began to shift. Resonance, he realized excitedly, could be applied as well to the relationship between single and double bonds—they did not have to be one or the other but could resonate between the two forms, leading to a stabilized partial double bond with its own peculiar properties. Resonance explained all kinds of structures that didn't fit into the old classical cubbyholes.
Virtually all of chemistry could be reevaluated in the light of this new idea, and Pauling set about doing it through the early 1930s. By applying his resonance ideas to various types of chemical bonds, then cross-checking and amending his theoretical results to fit what was known empirically about bond lengths and strengths, Pauling was able to produce a string of papers that set chemistry on a new course.
- - -
Pauling's reputation grew with each paper. In the spring of 1932 he took up Slater's invitation to become a visiting professor for a term at MIT, leaving Ava Helen in Pasadena in the last trimester of her third pregnancy. His eastern visit introduced his ideas to a number of figures in the Harvard-MIT chemical establishment, and he continued working on them nonstop between lectures and dinner parties. He had been thinking of ways to estimate the relative contributions of ionic and covalent bonds to any molecule, developing with a fellow Caltech faculty member, Don Yost, a system for estimating the theoretical strength of pure covalent bonds. With his new numbers in hand, Pauling could now compare his theoretical numbers to the real behavior of different pairs of elements as they formed compounds. The real-world bonds were always stronger than predicted—an added strength, Pauling assumed, from the stabilizing effect of resonance with an ionic form of the bonds. The greater the deviation, the more ionic character the bond had and the more the two elements differed in their ability to attract electrons. Using this system, it was now possible to answer old questions such as whether hydrochloric acid, HCl, was ionic or covalent—it was both, Pauling discovered, in the ratio 20:80.
