by Thomas Hager
But when he was to start his junior year, Belle gave him a shock: She told him he couldn't have the money he had earned. She needed it, she told him, to make ends meet at home. He would have to quit college for a year and work a regular job in order to make enough to finish the program.
Despite his independence in other ways, Pauling was still a dutiful son. He didn't want to live with Belle, but he did almost everything she asked. The summer before, he had even joined the Masons at her request, a prerequisite for Belle to be able to join some of her friends as a member of the allied Eastern Star women's group. (As soon as she did, he quit going to meetings of the Masons.) Now he swallowed his disappointment, telling himself that it was his obligation to help his mother, not hers to send him through college, and made plans to keep working with the highway commission through the academic year. Then he received an unusual offer from OAC. Pauling had shown himself a prodigy in the chemistry program, never receiving less than an A in any chemistry course and in the winter term of his sophomore year managing a perfect 4.0 grade-point average. The chemistry department didn't want to lose its promising student; at the same time, there was a need for teaching help to handle the mushrooming number of students taking introductory chemistry courses. The solution was simple: Pauling, at age eighteen, was offered a job as an instructor teaching quantitative chemistry, a course he had finished just the year before.
Although the teaching job paid only a hundred dollars a month— twenty-five dollars a month less than testing pavement—Pauling didn't hesitate. He had already recognized that his interests were more academic than industrial, and he realized the importance of gaining some teaching experience. As a student, he had paid enough attention to the art of teaching to know what worked and what didn't in a classroom, and his grasp of chemistry through the sophomore level was at least as good as that of most of his professors. After an initial period of nervousness, he found that he enjoyed lecturing and that the students appreciated his enthusiasm. Following his first term the mining students petitioned the department to let Pauling teach them quantitative chemistry, and the department administrators began gratefully assigning a number of courses to him, including chemistry for home economics majors. "Lots of times the students would say, 'Well, hell, he knows more than the profs, anyway. He could conduct the class better than they did,'" remembered one of Pauling's OAC classmates. "He was regarded as quite a brain in those days."
Pauling's yearlong teaching stint gave him money, confidence as a lecturer, and time to catch up with the latest research in the field. He was given a desk next to a secretary in the small chemistry library, where he learned touch typing and pored over chemical journals between classes.
The journal reading was important. At OAC the chemistry professors not only didn't do much research themselves; they didn't teach their students much about any of the current investigations in their field. There was little attempt to place chemistry in a historical context, to outline recent trends in the field, or to transmit the excitement that comes from the pursuit of knowledge. Pauling had to find that excitement on his own, and he did—in the journals.
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One of the papers that caught his attention was written by Irving Langmuir, a General Electric Research Laboratory chemist who had made a substantial reputation for himself—and a fortune for GE—by discovering a way to greatly increase the life span of electric lights. Gifted with a roving curiosity and supported by all the power of one of the nation's leading industrial concerns, Langmuir would later become the first industrial chemist to win the Nobel Prize.
His work with electricity and its effects on metals led him to speculate on the way in which the basic unit of electricity, the electron, was involved in the structure of atoms and molecules. The paper Pauling read was a sixty-six-page tour de force called "The Arrangement of Electrons in Atoms and Molecules." It was an expansion, as Langmuir generously noted, of the ideas of another American chemist, the head of the University of California chemistry department, Gilbert N. Lewis, which had been published before the war. Pauling read Langmuir's work, went back and looked up Lewis's, and had his eyes opened to a new way of looking at chemistry.
The Lewis and Langmuir papers were an attempt to bring chemistry in line with some of the baffling things physicists were discovering about the structure of atoms. And atomic structure was the question of the day. For two thousand years atoms had been considered the ultimate units of nature, the smallest of the small. (The name atom itself means "not cut.") John Dalton in 1808 set nineteenth century chemistry firmly on the path of atomic theory in his treatise A New System of Chemical Philosophy, in which he persuasively argued that unbreakable atoms formed compounds by linking with other atoms in simple whole-number proportions: one carbon with two oxygens to make carbon dioxide; one carbon with one oxygen to form carbon monoxide; one oxygen with one hydrogen to form water. (He had that one slightly wrong.) It wasn't until the 1880s and 1890s that cracks began to appear in Dalton's solid little spheres, caused by the discovery of strange new phenomena that his atomic theory couldn't explain, among them x-rays and radioactivity. In 1897 the theory of the indivisible atom was finally exploded by J. J. Thomson, the British physicist and head of the venerable Cavendish Laboratory at Cambridge, who stunned the scientific world by reporting the existence of particles one thousand times smaller than the smallest atom. Thomson called them "corpuscles." When it was quickly found that they appeared to be the elemental unit of electricity, the name that stuck was "electron."
It was the first sighting of the subatomic world. And it turned science on its head. Thomson's discovery of the electron would force a crisis of understanding in science, compel the development of a new kind of physics and a vastly revised chemistry—would in fact, perhaps more than any other single event, usher in the twentieth century.
Electrons, it seemed, were normal parts of atoms. And the new problem became figuring out what atoms were, now that they appeared to be composed of smaller pieces. Electrons, Thomson found, carried a negative charge. But under normal circumstances the atoms of which they were a part had no overall charge; therefore, there had to be a positive electrical component somewhere to neutralize the electrons. Thomson thought that perhaps the electrons were stabilized in a field of positive electricity, like raisins in a pudding. He was proved wrong by one of his former students, a brash New Zealander with a genius for laboratory work named Ernest Rutherford, who announced in 1911 that he had found evidence of a surprisingly different structure: At the center of the atom, according to his elegant experiments, was a very small, very dense nucleus that carried a positive charge. The rest of the atom, with the exception of the electrons, was empty space. Expanded to the size of a football stadium, Rutherford's atom would have a nucleus the size of a grain of rice on the fifty-yard line, with barely visible electrons orbiting the outer bleachers.
This was as astounding as Thomson's discovery of the electron. Rather than solid balls, atoms had become webs of fairy gossamer. Solid matter was mostly empty space. Rutherford's findings set off another round of theorizing. If the nucleus was that small and positively charged and the electrons were that distant and negatively charged, what held the whole thing together? Opposites attract, so why didn't electrons simply dive into the nucleus?
Physicists understood a great deal about moving objects and forces; Newton's theories and the work of his successors had given them the power to predict the movement of the planets based on a few earthly experiments. Certainly the same well-understood and proven laws of nature—the body of knowledge that would be known as classical physics—could be adapted to explain the workings of the atom. Rutherford himself proposed that atoms might be something akin to little solar systems, with electrons whizzing around the nucleus like planets around the sun. The speed of their flight could, he theorized, counterbalance the electrical attraction of the nucleus. Like most physicists, he thought in terms of fast-moving electrons; Rutherford's was a dynamic model of the atom
. But his atom didn't work. One deadly shortcoming was the classical requirement that any moving charged particle lose energy. Applied to electrons, this meant that Rutherford's atom would run down like an unwound watch until the electrons spiraled into the nucleus.
If not a solar system, what was the atom like? In the early decades of the twentieth century, answering that question would become the Holy Grail for a new generation of physicists.
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They were not alone in the quest. "The problem of the structure of atoms has been attacked mainly by physicists," Langmuir wrote in the paper that first caught Pauling's attention, "who have given little consideration to the chemical properties which must ultimately be explained by a theory of atomic structure. The vast store of knowledge of chemical properties and relationships, such as is summarized by the Periodic Table, should serve as a better foundation for a theory of atomic structure than the relatively meager experimental data along purely physical lines."
There was bit of jostling for position here between chemistry, the king of sciences in the nineteenth century, and physics, which would prove the dominant field of the twentieth. Both Lewis and Langmuir knew and appreciated physics—both had studied in Germany with the pioneers of physical chemistry, and Lewis was one of the first Americans to champion Einstein's theory of relativity—but they were chemists at heart.
And at the heart of chemistry, as Langmuir pointed out, was the periodic table, a Rosetta stone for chemists eager to translate nature's elemental language. The table began to take shape in the 1860s, when several chemists noticed that when they arranged the elements according to increasing atomic weight, certain properties—melting points, boiling points, chemical reactivity—seemed to rise and fall and rise again in a roughly periodic way. At least early in the table, that period seemed to be eight elements long. Start with an inert gas like helium, one of nature's most unreactive substances. Move eight steps up the table and there was another inert gas, neon. Move eight steps more and there was another, argon. It was the same with the highly reactive alkali metals: Lithium was eight steps away from sodium, which behaved much like potassium, another eight steps away.
Why eight was the magic number was unclear. Then, around 1913, it became accepted that each new element in the table represented not only an increase in weight but the addition of one electron to the preceding element. Somehow this regular increase in electrons was intimately tied to the periodic nature of the elements.
Lewis explained it in a paper he published in 1916. The inert gases were inert, he wrote, because they possessed an unusually stable grouping of electrons. Eight steps between each inert gas meant the addition of eight electrons; whatever this stable organization of electrons was, it happened eight at a time. This electronic "rule of eight" had been proposed before, but Lewis employed it to explain more about chemistry than anyone had. And he used it to create a new model of the atom. Rather than the physicists' solar system, he placed eight electrons at equal distances from one another and from the nucleus in three-dimensional space, then connected the dots. The result was a cube enclosing the nucleus, with an electron at each of the eight corners. Going up the periodic table and adding electrons, new cubes would be formed one electron at a time around the inner cubes, like boxes around boxes.
Lewis's atomic cubes did more than explain the rule of eight. They also offered explanations for phenomena that the physicists' solar system model could not, such as how one atom could link with others to form a stable molecule. In this matter of the chemical bond, Lewis and Langmuir theorized, the explanation again came from the natural tendency of an element to want to form a perfect cube filled with eight electrons. An atom with four extra electrons beyond a perfect cube— carbon, for instance—would become most stable when it connected with some combination of other atoms that offered it four more electrons, filling its outer cube with a perfect eight. This filling of cubes could be achieved, Lewis wrote, by sharing electrons with other atoms. And his theory stated that this sharing should happen one pair at a time, using the two electrons along an edge of the cube. Four hydrogens, for instance, could each pair their single electrons with one of the four bachelors in the carbon shell, giving it the equivalent of eight in its outer cube and creating the stable molecule CH,, methane. Sharing pairs of electrons, Lewis and Langmuir said, was the glue that held molecules together.
Their cubical picture of the atom was simple but powerful, offering at least a preliminary explanation of the character of the inert gases, the periodic table, and chemical valency, the long-known but unexplained capacity of elements to combine most stably with certain numbers of other elements.
And the Lewis and Langmuir model did something else that the physicists' atom could not: It fit with what chemists knew about the shapes of molecules. Chemists knew that molecules were not random assortments of atoms joined willy-nilly; their shapes were specific. In methane, for instance, the four hydrogens were linked to the carbon to form a tetrahedron, a three-sided pyramid. There was no plausible explanation of how a solar-system atom could lead to that sort of three-dimensional specificity, but it was easy to see how, with a cubical atom, the sharing of electrons along various edges could lead to the patterns known to exist in nature.
It was a breakthrough paper, Lewis's 1916 effort, one that Pauling always thought should have won him the Nobel Prize. Like many milestone papers in science, it planted the seeds of several important ideas. It focused chemists' attention on electrons, reinforcing the growing belief that chemistry was rooted, in general, in the structure of groupings of electrons. This emphasis on structure—"We must first of all, from a study of chemical phenomena, learn the structure and the arrangement of the atoms," as Lewis put it—would also have a great effect on chemistry. It established the chemists' place in studying atomic structure, directly challenging the physicists' solar-system atom, which could not explain valency or molecular structure. And, most importantly, it proposed that the chemical bond was made of pairs of electrons.
Lewis's cubical atom had its own problems, however. It quickly became known as the "static" model of the atom, as opposed to the physicists' "dynamic" model, because Lewis demanded that electrons stay relatively still at the cube's corners. Static electrons were impossible, the physicists argued; a negatively charged particle could not sit still at a small distance from a positively charged particle—electrostatic forces would pull them together. Lewis's characteristically audacious response was to propose that his model might be right and the accepted Newtonian laws of nature wrong. "Indeed, if we find it necessary to alter the law of force acting between charged particles at small distances," he wrote in 1916, "it will not be the first time in the history of science that an increase in the range of observational material has required a modification of generalizations based upon a smaller field of observation."
The argument was set aside as the United States entered World War I. Lewis began researching ways to defend soldiers against gas warfare, Rutherford began focusing on radioactivity, and it would be left to a new generation of young physicists to take up the problem of atomic structure after the war. On the chemists' side, it wasn't until Langmuir popularized and expanded Lewis's ideas around 1919 that they got the attention they deserved.
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Until he read Lewis's and Langmuir's papers, Pauling had been teaching his students the same crude chemical-bond theory developed in ancient Greece, in which atoms had a certain number of hooks and eyes with which to latch on to other atoms. Sodium, for instance, had an eye; chlorine had a hook. Sodium chloride was easily formed, and two atoms of chlorine could latch their hooks together to make CL, but two atoms of sodium couldn't combine. While the theory gave a picture that satisfied Pauling at the time he was learning it, it didn't explain anything about what the hooks and eyes were, that is, the nature of the forces that held atoms into stable aggregates.
The Lewis and Langmuir model did. The idea of shared electron pairs tied chemistry to atomic physics
in a way that made it possible to discuss why things happened in chemistry rather than simply describing what happened.
Even as a student, Pauling had been thinking about molecular structure in his courses in materials and metallurgy, where questions about the hardening of iron into steel and the ductility of metals were discussed in the context of crude atomic models. One idea from his classes that he later remembered was that a slip occurs along planes of atoms as a metal is stretched, with atoms in one level sliding over others. He had found something pleasing in visualizing atoms bumping along each other as the basis for an observable property.
This was a rare pleasure in chemistry, where students could read hundreds of pages of text describing the properties of substances— chemical formulas, molecular weights, densities, colors, crystalline forms, melting points, boiling points, solubilities, reactivities, refractabilities—without finding a single paragraph explaining whv these properties existed: why water froze at one temperature and methane at another, why graphite was soft and diamonds hard, why neon didn't react with anything, while fluorine, one place away from it on the periodic table, reacted with almost everything. Now, through
Lewis and Langmuir, Pauling was introduced to a more sophisticated way of looking at molecules, a way that explained more. "It was then," he wrote, "that I developed a strong desire to understand the physical and chemical properties of substances in relation to the structure of atoms and molecules of which they are composed."