A Resonating-Valence-Bond Theory of Metals and Intermetallic Compounds, 1949

Linus Pauling, 1949

As we have written elsewhere, Linus Pauling developed and championed the theory of resonance early on in his career. But for almost 20 years, the ways that metals might conform to the theory remained elusive. Even though he had a hunch that they too adhered to the tenets of resonance, he was not able to prove it definitively at the outset. Not until 1949, with the publication of “A Resonating-Valence-Bond Theory of Metals and Intermetallic Compounds” was he able to demonstrate that metals do resonate. 

Pauling’s interest in metals dated at least as far back as his undergraduate years at Oregon Agriculture College and continued to flourish during his graduate training at the California Institute of Technology. Later in his career, when asked to reflect on his contributions to the field of chemistry, he often spoke of his early work with metals as being important. This was especially so with his work on metals and resonance.


“A Resonating-Valence-Bond Theory of Metals and Intermetallic Compounds,” which was published in the Proceedings of the Royal Society London, serves as an addendum of sorts to Pauling’s previous papers on the nature of the chemical bond. Even though Pauling’s theory of resonance had been well-received for several years, confusion still existed amongst chemists (including Pauling) about how the theory might apply to metals; particularly the iron-group transition metals. Pauling believed that, like other elements, resonance must be used to explain the way that these metals bond, but the specifics proved difficult to pin down.

Pauling had always aligned himself with the notion that the properties of an element were connected to the configuration of its valence electrons. For example, Pauling knew that the arrangement of valence electrons gave Carbon its stable tetrahedral properties and salts their ionic properties. Metals, however, were harder to define due to the fact that they “showed great ranges of values of their properties, such as melting and boiling point, hardness and strength, and magnetic properties.”

Pauling initially focused on magnetism, and from this work, it was determined that metals had a high ligancy of either 8 or 12. Ligancy, or the number of compounds that each metal could bond to, (often referred to as Coordination Number) is similar but still distinct from valency, and for the metals that Pauling was studying, a ligancy of 8 or 12 was higher than their valence. This discrepancy seemed to indicate that resonance could not be applied to explain how metals bonded. But Pauling still believed that resonance was the answer, and set about trying to confirm this belief.


In his quest to understand if or how metals resonated, Pauling first needed to clarify whether or not previous thinking about metals was correct. Prior to the publication of his 1949 paper, it was believed that the d orbitals did not participate in bonding with the iron-group transition metals, such as Manganese, Iron, Cobalt, Nickel, and Copper. Using these assumptions, Pauling modeled the predicted properties of such metals and found that they would have low melting points and weak bonds. The lab work indicated otherwise however, leading Pauling to conclude that the current understanding was incorrect, and that something else had to be going on with the bonds in order to account for their physical properties.

For several years, Pauling had been unable to devise a competing theory that would explain the unique properties of metals, but he felt certain that the answer would be revealed through the application of quantum mechanics. By specifically using the wave theory of quantum mechanics, Pauling calculated that metals used 8.28 orbitals instead of the expected 9. While he was sure that the math was correct, he struggled to understand what was accounting for the missing 0.72 orbitals, and for almost nine years he worked to reconcile the discrepancy. Ultimately though, he realized that 0.72 orbitals was, in fact, the answer he was looking for; they were what gave metals their unique properties.

In its essence, Pauling breakthrough was that the extra 0.72 orbitals was not a mathematical anomaly, but instead an extra orbital, one that he called a “metallic orbital.” The metallic orbital was, according to Pauling, “required” to give metals their “characteristic properties, especially that of electronic conductivity of electricity.” In Pauling’s view, substances lacking the metallic orbital that still maintained some metallic properties, such as electric conductivity, should be classified as metalloids.

Pauling’s conceptualization of the metallic orbital allowed him to subsequently reframe his understanding of resonance. When Pauling developed the theory, it was more specifically known as synchronized resonance; a state in which all bonds resonate in the same way and valence is undisturbed by the process. Pauling realized that, in order for metals to resonate, a different kind of resonance – unsynchronized resonance – would be needed. In unsynchronized resonance, instead of all bonds resonating simultaneously, a single bond could resonate on its own. In the case of metals this happened in the metallic orbital, and it was found that unsynchronized resonating metals conferred unusual stability and high ligancy.


Once Pauling made these connections, he was ready to publish his paper, which, first and foremost, sought to prove the existence of a metallic orbital. To do this, Pauling had to show that the existing understanding of valency was incorrect, and then to demonstrate that resonance – specifically unsynchronized resonance – could account for both the predicted and the observed properties of metals.

Using Lithium as his model element, Pauling explained that the current understanding posited a valence structure that consisted of 2s orbitals, and that this was the commonplace belief because, “the [proposed] molecular orbitals correspond to electron energies.” By using the wave function however, Pauling found that the assumed orbital structure did not correspond to observed electron energies. Specifically, he calculated that if Lithium had 2s orbitals, the predicted heat of formation would be much different than was the observed heat of formation; a calculated difference of 32.4 kcal/g. In short, the current understanding was incorrect.

Pauling’s next step was to use resonance to explain Lithium’s metallic properties, and to devise an arrangement of valence electrons that would match predicted energies with observed energies. Using the idea of unsynchronized resonance, Pauling suggested that if, instead of 2s orbitals, Lithium’s valence electrons were actually in resonance, this new arrangement could be “responsible for the difference in energy.” In other words, without resonance the electrons existed in a fixed energy state. But if the electrons were in resonance, their energy state would instead be a hybrid of all possible energy states. This circumstance, Pauling postulated, would confer lower energies to the molecule and bring predicted and observed energies into alignment. In the case of Lithium specifically, Pauling found that the valence electrons were in 2s orbitals, but also 2p orbitals. Pauling worked through different metals throughout the remainder of the paper to further support his thinking.

The theoretical work that Pauling put into this paper completed the framework for resonance. And interestingly, even though the ideas that he presented were accepted and widely used for decades, some of the math in the paper was not fully validated at the time. In fact, it wasn’t until 1984, when Pauling revisited the 0.72 orbitals, that the math was completed. By then Pauling had acknowledged that the calculations he used to get to 0.72 orbitals had been crude, and based on in part on educated guesses. But by the 1980s, many of the unknowns were known, which allowed Pauling to revisit the math with more precision and compare it to his work from the late 1940s. Using clean, contemporary data, Pauling confirmed that the new calculation was “in exact agreement with the observed value of 0.72.”

The Nature of the Intermolecular Forces Operative in Biological Processes, 1940

Linus Pauling, Max Delbrück and Max Perutz at the American Chemical Society centennial meeting, New York. April 6, 1976.

In 1940 Linus Pauling, along with colleague Max Delbrück, authored a three-page article that was published in the July issue of the journal Science. The length of the article was shorter than typical for Pauling, but what made it even more unusual was that it was not about Pauling’s findings. Instead, the piece served as a critique of a different article published earlier that year by a German scientist, Pascual Jordan.

In it, Jordan argued that when like molecules bonded, they were attracted more strongly than when dissimilar molecules bonded. Jordan believed that this stronger attraction of like molecules conferred special properties to these bonds, especially when they occurred in living cells. Pauling and Delbrück totally disagreed with this idea. Instead, the duo believed that it was a molecule’s complementary nature that conferred stability, an idea in opposition to Jordan’s concept of similarity.

In the two decades preceding these papers, chemists had come to look at their field in different ways, due mainly to advancements in quantum mechanics. This was certainly true for Pauling, who rapidly developed a reputation for using these new ideas to solve old problems. One line that he did not cross however, was the application of quantum mechanics to help “solve” topics that were already well understood and not in conflict.

For Pauling, one such instance was the basis of molecular attraction, and how that attraction created stability in a newly formed molecule. This idea, however, was something that other scientists found worth examining; Pascual Jordan in particular. Accordingly, and armed with a new set of quantum mechanical theories, Jordan set about attacking a question that others, including Pauling, believed not in need of answering.


Pascual Jordan

Pascual Jordan was born in Germany in 1902 of Spanish lineage. Though initially interested in the arts, Jordan studied math and physics in school, completing his physics Ph.D. in 1924. His ideas at this time were novel, with no less a figure than Albert Einstein taking note of his dissertation. But Einstein did not agree with certain of the hypotheses that Jordan was putting forth, many of which used quantum mechanics to consider the photon nature of light. While Einstein felt that there wasn’t necessarily anything wrong with Jordan’s ideas, he did not agree with the logic that informed them, and wrote missives in opposition.

But others supported Jordan’s work and, soon after graduation, he began working with a circle of colleagues that included Werner Heisenberg. During this time, Jordan became one of the biggest proponents of quantum mechanics and, along with Heisenberg, helped to unlock many of its secrets. Jordan was also a member of the Nazi party, joining when Germany entered World War II and remaining so until at least the end of the war. Nonetheless, Jordan helped to develop key theories in physics and math which are foundational to the fields today.


Though Jordan’s legacy today is marred by his political positions, when he wrote his 1940 paper about the attraction of molecules in biological cells, he did so from a position of authority. As noted, the foundation of the paper is the idea that identical molecules are attracted to one another in a special way that does not exist for dissimilar molecules and that, because of this, the bonds formed in molecules are more stable than is the case with other bonds. Jordan’s hypothesis, if true, would have been groundbreaking and consequential for all sorts of bonds, especially those in living cells.

Understandably, the paper created a lot of commotion when it was published. Pauling, who at that point was also an authority on quantum mechanics and resonance theory, was no doubt among those surprised by Jordan’s proposition. After reading it though, he immediately saw its flaws. In it, Jordan himself admitted some doubt that resonance could work in the manner that he was suggesting, and Pauling was sure that the ideas were wrong. Wishing to publish a rejoinder, Pauling began looking for a co-author whose expertise centered around bonds in living cells, and Max Delbrück was just such a figure.


Like Jordan, Delbrück was born in Germany in 1906. Interested in the stars, Delbrück began his studies in astrophysics, but changed directions upon meeting a physical chemist, Karl Bonhoeffer, who was eight years his elder. Fascinated by Bonhoeffer, Delbrück switched to physical chemistry in a ploy to become his friend, a tactic that ultimately worked well. The timing of the switch was also fortuitous as Delbrück entered the field at the beginning of the quantum revolution. After graduation, Delbrück studied all over Europe with scientists included Wolfgang Pauli and Niels Bohr. He eventually spent a few years at the California Institute of Technology on a Rockefeller Foundation fellowship, during which time he met Pauling and co-authored the 1940 paper. After leaving Caltech, Delbrück focused his research on bacteriophages and eventually won the 1969 Nobel Prize in Physiology or Medicine for this work. 

Even though Delbrück’s Nobel honor was nearly thirty years down the road, by 1940 he was already well-versed on the ways that living cells operated, making him a formidable writing companion. In their paper, Pauling and Delbrück argued that Jordan’s fundamental idea could not be correct because the stability of a molecule was conferred by the complementarity its components, not their similarity. By way of explanation, the duo first put forth the understanding that a stable molecule is one in which molecular distances are relatively short. This is a circumstance, they argued, that can best be achieved when complementary forces are working together, such as positive ions attracting negative ions. In other words, in a bonding pair “the two molecules must have complementary surfaces, like die and coin.” The like molecules that Jordan was advocating for were not complementary by definition; rather, they were identical, or close to it. Pauling and Delbrück acknowledged that “the case might occur in which the two complementary structures happen to be identical” but still their stability “would be due to their complementariness rather than their identity.”

Even though Pauling and Delbrück’s article was quite short, its message was clear: Jordan was plainly wrong. As they wrote, “We have reached the conclusion that the theory can not be applied in the ways indicated by him [Jordan], and his explanations of biological phenomena on this basis can not be accepted.” In short order, the scientific mainstream came to agree with their point of view, and Jordan’s ideas soon faded away.

The Nature of Interatomic Forces in Metals, 1938

Linus Pauling, ca. 1930s

“In recent years I have formed, on the basis of mainly empirical arguments, a conception of the nature of the interatomic forces in metals which has some novel features.”

­-Linus Pauling, 1938

Prior to the publication of this article, which appeared in the December 1938 issue of Physical Review, much about the interatomic forces operating in metals was either unknown, or theoretical predictions did not align properly with observed data. In publishing this paper, Linus Pauling first sought to align the incongruencies between theory and data for the transition metals, such as iron, cobalt, nickel, copper, palladium, and platinum. He was then able to correctly predict properties including “interatomic distance, characteristic temperature, hardness, compressibility, and coefficient of thermal expansion” by discarding previously held assumptions and inserting new – and correct – assumptions about transitions metals.

The most significant idea that Pauling introduced with this paper was the notion that the valence shell electrons – those in the outer shell – play a part in bonding. Previously, scientists believed that these electrons made “no significant contribution” to bond formation. Pauling was able to establish otherwise, and used this breakthrough to both align observable data with theoretical data, and make other predictions about transition metals.


The 1938 paper was written in the wake of a revolution within the world of chemistry. A raft of new theories brought about by a widening understanding of quantum mechanics was generating intense excitement for scientists world-wide, and the tools that quantum mechanics provided for helping to “correct” previous understandings of the chemical bond were of paramount interest to many. Pauling, of course, was a leader in this area, his body of work ultimately garnering the 1954 Nobel Chemistry Prize for “research into the nature of the chemical bond and its application to the elucidation of the structure of complex substances.”

Within this area of focus, many scientists were especially interested in exploring the ways that metals bonded because, as noted, the observed data did not match up with theory. Pauling sought to mend this gap by using quantum mechanics to look at interatomic forces in a novel way. Prior to the Physical Review paper, chemists believed that when metals bonded, their valance shell electrons played only a small role in the resulting structures. Pauling argued otherwise, and put forth an important new theory that the valence shell electrons contributed to the process through resonance, a theory that he had developed earlier in the decade and continued to champion.


Because the crux of Pauling’s scientific intervention was to prove that valence shell electrons are involved in bonding, most of the paper is devoted to supporting this claim. The primary tool that Pauling uses to craft his argument is an analysis of temperature predictions. According to the reigning theory regarding metals and valence shell bonds, when bonding occurred, the electrons would bond in a manner that would create a moment of ferromagnetism. Specifically, it was theorized that these ferromagnetic moments would be temperature dependent, meaning that as the temperature of the metal changed, its degree of magnetism would also change in a predictable way. Experiments had shown however, that when metals bonded, their ferromagnetism remained independent of temperature.

Pauling exploited this piece of information and used it to support his theory. According to Pauling, if metals bonded through resonance, they would create ferromagnetic moments that were temperature independent, a hypothesis that correctly aligned with the observed data.

To develop his argument, Pauling made specific use of the element Vanadium, which has an electron configuration of 3d34s2. Under the old model, Vanadium’s valence electrons could only interact weakly in bonding if, at most, two of the 4s2 electrons were involved in a bond. This, according to Pauling, would create ferromagnetism which would decrease with increasing temperature, meaning that it was temperature dependent. On the contrary, the experimental evidence showed that Vanadium’s magnetism was temperature independent. This meant, therefore, that weak valence interaction during bonding was not possible.

Pauling’s alternative suggestion was that all of Vanadium’s valence electrons were involved in bonding through resonance; not just the two 4s2 electrons, as previously believed. Further, if the valence electrons bonded through resonance, the ferromagnetism of their structure would be temperature independent, a prediction that aligned with the observed data.


Once Pauling was able to prove that the valence electrons in Vanadium bonded through resonance, he then began to apply the concept to all transition metals. As with the previous example, Pauling continued to support the concept by comparing predicted outcomes with empirical data. And once again, when viewing the bonding through the prism of resonance, predicted outcomes of magnetic moments began to align with the empirical data.

Pauling then took it another step by repeating the exercise with interatomic distances. As demonstrated in his paper, a resonant structure would correctly predict the interatomic distances that had been observed for many bonds. Pauling also claimed that other properties, such as the “compressibility, coefficient of thermal expansion, characteristic temperature, melting point, and hardness” would likewise correctly align with experimental evidence, once resonance was used to explain valence shell bonding.

Though clearly a significant breakthrough, the assertions that Pauling made in his paper were grounded in work done by others — notably the quantum mechanical theory of ferromagnetism developed by Heisenberg, Frenkel, Bloch, Slater, et al., and Wolfgang Pauli’s theory of the temperature-independent paramagnetism of the alkali metals. And while the 1938 article acknowledges these debts, it also attempts to improve upon them.

This was especially so with the quantum mechanical theory of ferromagnetism. As we have seen, Pauling successfully applied the idea of temperature independence and ferromagnetism to support his claims, but he also found one aspect of the theory to be needlessly bothersome. As Pauling noted, in order for much of the theory to work on a mathematical level, scientists were compelled to assign positive numbers to all unpaired valence electrons. Pauling recognized that this was only necessary if it was assumed that valence electrons did not play a large role in bonding for metals. Under a resonant scenario, Pauling was able to show that the math could still work if the valence electrons were negative and that, once again, “this conclusion agrees with the observation.”

The Theoretical Prediction of the Physical Properties of Many-Electron Atoms and Ions. Mole Refraction, Diamagnetic Susceptibility, and Extension in Space, 1927

Linus and Ava Helen Pauling in Copenhagen, May 1927

[Ed Note: Today and in the three posts that will follow, we will be taking a close look at four important scientific articles published by Linus Pauling between 1927 – 1949.]

In this ambitious and hugely influential paper, Linus Pauling applied his theory of screening constants to various problems, including electric polarizability, diamagnetic susceptibility, and the sizes of ions and atoms. Pauling was fundamentally interested in pursuing this topic because of his desire to merge the new quantum mechanics – which embraced wave functions – with older ideas in order to make predictions about molecular properties like mole refraction and diamagnetic susceptibility in space.

Especially during the early phases of his career, one of Pauling’s signature rhetorical tools was to put forth a bold assumption that would serve to simplify the predictions made later on in a given article. This paper is among the best examples of that approach. In it, Pauling developed mathematical relationships that, when applied, could help the reader make generalizations about molecules. But in moving through these calculations, Pauling had to make some assumptions, oftentimes without the aid of hindsight to determine whether or not they were correct. One enduring legacy of this paper is that many of Pauling’s assumptions were indeed correct, and its findings have thus remained relevant across the decades.


Another contributing factor to the paper’s success was that Pauling was in the right place at the right time. While working towards his PhD at Caltech, Pauling enthusiastically followed the rapid development of quantum mechanics in Europe and elsewhere. Pauling was particularly interested in the work of Arnold Somerfield in Munich and Niels Bohr in Copenhagen, and wrote to both to inquire about research opportunities. Bohr never responded but Sommerfeld did and, with his support, Pauling secured a Guggenheim Fellowship that allowed him to live and work in Europe for 19 months. During that period, he spent most of his time with Sommerfeld at the Institute of Theoretical Physics, though he did visit Bohr in Copenhagen as well as Erwin Schrödinger in Zurich.

Pauling’s residency in Europe proved auspicious, in part because Sommerfeld and his colleagues were working on uniting new ideas with old, a task not being readily pursued in the United States at the time. As they moved forward with their work, the European scientists began to solve more and more problems with quantum mechanics, cementing in Pauling’s mind the utility of the approach as a way forward.


Gregor Wentzel (Image credit: Emilio Segre Visual Archives)

One project important to Pauling’s paper was being led by University of Leipzig physicist Gregor Wentzel. A colleague of Sommerfeld, Wentzel was seeking to apply quantum mechanics to x-rays in order to calculate the screening constants of electrons in large and complex molecules. His project had hit a snag however, in that he was unable to find agreement between the observed data and those predicted by theory. After scrutinizing his work, the young Pauling found that Wentzel had made errors in his calculations. Once Pauling had corrected these miscalculations, he found that there was in fact agreement between the observed and predicted data, which meant that Wentzel’s work was actually correct. In so doing, Pauling had confirmed the value of quantum-mechanical calculations in predicting screening constants of electrons in complex molecules.

Armed with this information, Pauling recognized that he could use these same calculations to make predictions about electron arrangement in molecules and the relative size of ions, among other properties. This led to the publication of his paper, The Theoretical Prediction of the Physical Properties of Many-Electron Atoms and Ions. Mole Refraction, Diamagnetic Susceptibility, and Extension in Space, which appeared in 1927, published by the Proceedings of the Royal Society.

In its essence, the article used the wave mechanical feature of quantum mechanics to make predictions about molecules, an approach that emerged directly from Pauling’s exposure to European efforts to unify old ideas with new. And even though it was not the first time that Pauling had written a paper utilizing quantum mechanics, it was certainly his first publication in which he used these novel tools to make predictions about molecular properties. 


Fundamental to these predictions were three key assumptions that Pauling put forth at the beginning of his paper. The first was that,

each electron shell within the atom is idealized as a uniform surface charge of electricity of amount-zi e on a sphere whose radius is equal to the average value of the electron-nucleus distance of the electrons in the shell.

The second assumption stated that,

the motion of the electron under consideration is then determined by the use of the old quantum theory, the azimuthal quantum number being chosen so as to produce the closest approximation of the quantum mechanics.

And the third assumption was that,

since so does not depend on Z, it is evaluated for large values of Z, but expanding powers of zi/Z and neglecting powers higher than the first, and then comparing the expansion with that of the expression containing Z-so in powers of so/Z.

Armed with these assumptions, Pauling was able to issue a collection of predictions about molecules, particularly concerning mole refraction and diamagnetic susceptibility. Prior to his doing so, chemists lacked the necessary tools for making predictions of this sort, meaning that certain chemical properties remained hazy or unknown.

This issue was particularly salient for the hydrogen atom. In the months leading up to the paper’s publication, a huge debate had emerged concerning the polarizability of hydrogen. The prevailing formula had been proven incorrect in 1926, after which time a race ensued to find a new, more suitable equation. Eventually a successor formula was developed, but it was criticized as being “a conservative Newtonian” model. Agreeing that a more robust approach was needed, Pauling set about applying quantum mechanics, and based on his three assumptions, he derived the following:

Knowing full well that the equation was based on his three assumptions, and anticipating resistance, Pauling pre-emptively argued that “it might be thought that these values of ɣ are not correct because of the fact that the electron shells actually do not consist of hydrogen-like electrons, but rather themselves of ‘penetrating electrons.'” However, “as Z [a surface harmonic] increases, the ‘penetrating orbits’ become more hydrogen-like” and therefore should be ignored because any error found would be “negligible.” Having put forth this solution to the problem of hydrogen, Pauling was then able to more broadly demonstrate the utility of his ideas.

Indeed, even though much of the work in the paper made assumptions that were oftentimes crude – such as using data from the valence shell electrons only – Pauling was able to create complex (and, as it turned out, fairly accurate) tables of polarizability of ions, diamagnetism screening constants, and mole refraction, among predictions.

It is clear that Pauling believed strongly in his paper, which he felt would “make possible the accurate prediction of the properties of any atom or ion.” And though the approach would sometimes only yield “approximate values of the physical properties of ions” based on his three assumptions, the importance of the work was not diminished as, oftentimes, directly observed data “may not exist under conditions permitting experimental investigation.”

Ava Helen Miller in Corvallis, 1921-1923

Ava Helen Miller, ca. 1922. Note annotation by Linus Pauling: “Ava Helen about when I met her.”

[Part 2 of 2]

Ava Helen Miller entered Oregon Agriculture College, now Oregon State University, as a freshman in the fall of 1921. During that first year, she met Linus Pauling, and shortly thereafter the two fell in love and became engaged. This chance meeting derailed her plans to continue her formal education, and after five terms at OAC, she left school to start her life with Pauling. But even though she spent less than two academic years in Corvallis, Ava Helen made the most of her experience.

The tenth of twelve children, Ava Helen was born in 1903 in Beavercreek, Oregon, an unincorporated area southwest of Portland. When the time came to begin considering colleges, OAC proved a natural choice because, by 1921, three of her siblings were already attending the school. These siblings were her brother Milton, a senior majoring in agriculture; another brother Clay, a junior also studying agriculture; and sister Mary, a senior in home economics. Like Mary, Ava Helen entered OAC as a home economics major. And instead of living in a women’s dormitory like most first-year students, Ava Helen lived with her siblings in a Corvallis rental home that her mother Nora also resided in.


Though college was not an ordinary expectation for women of the era, a higher education seemed to be in the cards for Ava Helen. She had always excelled as a student, graduating from Salem High School in three years as the class president. Buoyed by her past successes, Ava Helen enrolled in a wide variety of courses during her OAC tenure, including Spanish and French, languages for which she showed a particular aptitude.

Generally speaking, her performance in college improved as she became more settled. During her first term, she received mostly Bs with a C in general chemistry and an A in English composition. Interestingly, during her second quarter, Ava Helen received her only F, which was also in English composition. That winter term was clearly challenging with respect to academics, as her grades were somewhat worse than was typical for the rest of her OAC time. (Coincidentally or not, this was also the term in which she met her future husband.) But later, her grades improved. In her last quarter, she took Child Care and received an A, perhaps a foreshadowing of her future responsibilities as a mother of four.

Chemistry was always one of her best subjects, and instructors besides the young Pauling took note of her chemistry aptitude. One of them, an organic chemistry instructor named Mr. Quigly, remarked that she was among his best students. Despite this aptitude, Ava Helen notably received a B in general chemistry during the winter term of her first year, a grade that she felt was lower than she deserved. This belief was accurate, but because the instructor for the course was her soon-to-be husband, the instructor – “boy professor” Linus Pauling – deliberately lowered her grade from an A to a B, in an effort to obfuscate any potential impropriety. This decision remained a source of good-natured ribbing between the couple for years to come.


Ava Helen pursued interests outside of academics during her stint in Corvallis. She joined the school’s drama club, known as the Mask and Dagger, during her first year. The club produced several plays including larger productions like Shakespeare’s “A Midsummer Night’s Dream” and “Clarence,” a modern “light comedy” by Booth Tarkington. Near the middle of the year, Ava Helen hit the stage in a one-act offering titled “Pierro by the Light of the Moon” by Virginia Church; Ms. Miller was cast as Columbine, the second lead.

During her second year at OAC, Ava Helen also served as secretary of the Lyceum Club, which consisted of “musicians, readers, and lecturers.” The club focused on coordinating cultural events for the local community and, during Ava Helen’s year, hosted several prominent artists and musicians, including the Flonzaley String Quartet.


A discussion of Ava Helen’s time at OAC would be remiss if it did not include mention of her romance with Linus Pauling. During the winter term of her first year, she was enrolled in a general chemistry course that was mandatory for all home economics majors. Pauling, himself a fifth-year senior at OAC, was assigned to teach the class, though he was only a few years older than the students that he was asked to teach. Shortly after their first encounter in January 1922, Ava Helen and Linus hit it off, and by the end of the school year they were engaged.

In the summer of 1922, Pauling left Oregon to begin his graduate studies at Caltech while Ava Helen continued her schooling, starting her second year at OAC. The initial plan was for Ava Helen to stay at OAC and complete her degree while Pauling worked on his Ph.D, and the pair wrote to each other nearly every day, and saw one another on breaks. But the waiting and separation proved too much, and by the end of the second quarter of her second year, in March 1923, Ava Helen dropped out of school in order to plan her wedding. In June 1923, the couple married and Ava Helen moved to Pasadena to begin a new life.

Ava Helen’s OAC

Ava Helen Miller, fall 1921

[Part 1 of 2]

Ava Helen Miller entered Oregon Agricultural College (OAC), now known as Oregon State University, in the fall of 1921. She spent a total of five quarters at the school; one quarter shy of two full academic years. And even though her experience in Corvallis was relatively short, it was a memorable period for Ava Helen, one culminating in her meeting her future husband, Linus Pauling.

During the 1921-22 school year, OAC was home to its largest student body to date – more than 7,000 enrolled for at least one quarter, and a graduating class of 3,147. (2,040 men and 1,107 women.) The freshman class was 300 students larger than had been the case the previous year and the student body included 47 international students from Bolivia, Palestine, Peru and South Africa, among other far away lands. An additional 995 out of state students pursued an OAC education, with Washington leading the way (384) followed by California (315) and Idaho (111). Within the college, Home Economics awarded the most degrees that year, tallying 889; the School of Agriculture came in a close second with 876 degrees.

Those enrolled from out of state did so despite the fact that, in the fall of 1921, OAC began charging non-residents tuition – $20 per term – for the first time. The change impacted new students only, meaning that existing out of state students could still attend for free. (Ava Helen was a resident of Oregon, so this did not affect her.) Understanding the financial burden that this decision could have on its students, the OAC Board decided that non-Oregon residents who had served in World War I could receive a 50% discount on tuition. According to the monthly OAC Alumnus, the fees were enacted to “prevent an undue influx of students from other states and at the same time [provide] additional income […] in the construction of buildings.”


OAC’s post-war population bump coincided with shifts in the types of activities available to the college’s women students. Several clubs for women were sponsored, including those dedicated to riflery and mandolins, but new sports were an area where women were really able to broaden their horizons. OAC co-eds competed in baseball, basketball, swimming, track, tennis, archery, volleyball, and the afore-mentioned riflery; dance was also organized as part of women’s athletics.

As was the case most everywhere however, traditional men’s athletics remained far more popular across campus and the community. The most prominent sport at OAC was football and the team, affectionately known as the “gridiron squad,” was lavished with rituals and processionals prior to each game. Among the most extravagant was a feeding party where the players dined on “big juicy steaks cooked rare, toast crisp through and through, big baked potatoes that rival the prizes of the Pullman diner [a popular upscale train car from the time] – all topped off with ice-cream.” And though the men ruled the sporting world at OAC, these epic meals took place in the women’s Home Economics Tea Room, as that was clearly the place for the best food on campus.

With more people at the school, more options to choose from, and requirements that first- and second-year women sign up for P.E. credits, its no surprise that, during Ava Helen’s second year, women’s physical education courses saw more enrollees than ever before. Elective options included seasonal sports, aesthetic dancing, swimming, or apparatus work, which is similar to gymnastics. The era’s women did not play softball, but were allowed to play baseball (with other women). To help manage growing numbers, and to harken back to times when the program was much smaller, the Physical Education Club arranged weekly dinners to assist students in getting to know one another.


Competitive speaking was another area opening up to OAC’s women. By the fall of 1921, women had three different debate teams available to them – varsity, junior varsity, and intramural. During Ava Helen’s first year, the women’s Dual Debate Team (junior varsity), competed against the University of California and was asked to address the possibility of Irish independence. Meanwhile, the varsity squad was charged with discussing: “That the principle of the closed shop should be applied to all American industries. Farming and all industries employing less than three men are excepted.”

Intramural debate tended to pit classes against one another, with a 1921 contest asking freshmen and juniors to consider federal government ownership and operation of coal mines. Other topics shed light on the culture of the state a century ago – notably, a men’s interclass competition considering whether or not “Oregon should enact a law prohibiting Orientals from acquiring land within the state.”


By Ava Helen’s time, women had clearly emerged as a priority for OAC. Importantly, in the fall of 1922, the college finished construction of its third women’s dormitory, Margaret Snell Hall. And a year prior, OAC had welcomed it newest Dean of Women, Mary A. Rolfe.

Rolfe had worked for various organizations before OAC, including the YWCA and the University of Iowa, but immediately prior to her tenure in Corvallis she had served in France as a “searcher.” In that role, she worked with other stenographers to write down wounded soldiers’ last words, which would then be conveyed to their families back home. As Dean, Rolfe believed that “My relation to young women should be that of friend,” an ideology that she upheld by creating opportunities for women and treating them as capable and inquisitive. One early example of this was her sponsoring the Oregon chapter of the Home Economics Conventions, a program “of interest to all women, whether professional home economics women, teachers, or homemakers.”

Ava Helen Miller took advantage of all that her college had to offer during a period of significant change. An exploration of her experiences at OAC will be the focus of our post next week.

Pauling’s OAC: Super Senior Year

[Ed Note: School starts today here at Oregon State University! As we have for the previous four years, we take this opportunity to look back at Linus Pauling’s undergraduate experience in Corvallis, this time documenting his “super senior” experience as a Beaver.]

The 1921-22 academic year was Linus Pauling’s fifth at Oregon Agricultural College (OAC), now known as Oregon State University. In the twelve months to come, Pauling would finish his coursework; graduate; become engaged; and move to Pasadena, California to begin his graduate studies. Needless to say, it was a year of great change for Pauling, but one that he embraced with excellence.

In the fall of 1921, seniors at OAC were welcomed back to the college through a series of “Get Acquainted” dances, which were aimed at helping them become more comfortable with their apical position in the social hierarchy. Though these dances were a running tradition at OAC, each senior class approached them in their own way. During the 1921-22 academic year the dances were themed, with one particularly memorable event, the Goof Dance, challenging participants to wear the craziest outfits they had. 

For Pauling, the start of the year marked a continuation of his effort to earn solid marks and gain entry into a good graduate program. Throughout his time at OAC, he had always applied himself, and by his senior year, those efforts were evident. As with other colleges, the OAC Beaver yearbook included basic information on all its seniors, as well as additional details documenting their participation in extracurricular organizations and clubs. These blurbs often consisted of a handful of words, but Pauling’s was, not surprisingly, several lines long.

As per OAC custom, Pauling’s entry lists his major (Chemical Engineering), his hometown (Portland), and his fraternal membership (Delta Upsilon). Other decorations included his membership in Sigma Tau, the engineering honor society into which he was inducted during his junior year and served as secretary during his final year at OAC. His participation in the Scabbard and Blade, a military honor society that he joined during his junior year, is also listed. During his senior year, Pauling served as a Captain in the Reserve Officer Training Corps, which was recognized by the yearbook as well. So too was he a member of the Chemical Engineering Association (junior year treasurer) and the Chi Epsilon civil engineering honor society (junior year president). His efforts in competitive speaking were noted as well.

Pauling’s accomplishments were likewise praised by outside organizations. One of them, the Oregon Alumni Society, heralded Pauling’s admittance into the Forum Honor Society, OAC’s most prestigious academic group. Pauling, along with sixteen other students, was admitted for his “excellence in scholarship, leadership in school activities and strength of character.” OAC President William Jasper Kerr welcomed the new members personally, a group that also included Pauling’s friend and future colleague, Paul Emmett. 


Linus Pauling on OAC graduation day, June 1922.

As graduation day neared, Pauling was asked to deliver the senior class speech. He was a likely choice to fill this role, given his strong academic standing and his success in a junior year debate contest. But unlike past years, where speakers tended to offer fairly generic observations, Pauling’s speech was notably more pointed.

Pauling viewed the speech as an opportunity for him to position himself as a scientist, and he focused his rhetoric on contemporary world events as observed through a scientific lens. A main thrust of the talk was his belief in scientists’ duty to use their tools for good. In exploring this, he referred to the scientific developments that had advanced weaponry options, including chemical weapons, during World War I. Pauling also expressed a feeling that science was being used to create income gaps and remove humanity from workspaces, before suggesting that “the country is crying for a solution to these difficulties, and is hopefully looking to the educated man for it.” This call to action was the real point of his talk, which ended with an exhortation to his classmates that they repay OAC in the years ahead through acts of service in their communities.


Newly arrived at Caltech, Pauling poses on the back of a student’s car.

Pauling was well-aware of the need to move beyond OAC to continue his learning, and throughout the year the decision of where to go for graduate studies weighed heavily on his mind. Always keen on a future in chemistry, Pauling stayed current on recent developments in the field and knew that there were a handful of institutions equipped to provide him with an advanced education that could keep up with the changing times.

He decided to apply to four graduate programs: Harvard University, the University of California at Berkeley, University of Illinois, and the California Institute of Technology. Of these schools, Pauling was perhaps most attracted to Berkeley because it was headed by G.N. Lewis, who had discovered that electron bonds are shared. Harvard was also enticing, in part because its program was led by Theodore Richards, who was America’s only Nobel Laureate in chemistry at the time. Richards had attended the University of Illinois for his graduate work, and this connection had helped to boost its program. Caltech, by comparison, was the smallest and newest of the possibilities in which Pauling was interested.

Pauling eventually opted for Caltech, a decision that was made, in part, because of a fortunate sequence of events. All of the universities that Pauling wanted to attend were interested in him, and Harvard offered an attractive fellowship that would cover his tuition. But shortly after receiving this offer, Caltech’s letter arrived. Like Harvard, the Pasadena school offered a full-ride fellowship, but Caltech’s package also included a $350 stipend to work as a teaching assistant in undergraduate chemistry courses. Importantly, Caltech had also accepted Pauling’s close friend, Paul Emmett, and the two would ultimately live together for their first year as graduate students. These two factors tilted the scale in Caltech’s favor, and Pauling would remain at the Institute for more than forty years.

Summer Break

Linus Pauling at the Oregon coast with his cousin Rowena, summer 1918

We’re taking a few weeks off! We’ll resume our usual posting schedule in September — please check back with us then for more from Pauling’s world.

Cancer and Vitamin C: Crossroads of New Research

[Part 9 of 9]

In 2018, a third edition of Ewan Cameron and Linus Pauling’s book, Cancer and Vitamin C, was published. Released nearly a quarter century after Pauling’s death, this edition marked the first time that someone other than Pauling or Cameron contributed to the volume. The second edition, which was published in 1993, included an updated preface and a few new indices, but the text itself remained entirely Cameron and Pauling’s. The 2018 edition also includes a new preface and a new index, but both were written by Stephen Lawson of the Linus Pauling Institute at Oregon State University. Lawson used the new additions to ably present recent scholarship connecting the efficacy of vitamin C to the treatment of cancer.


One bit of background that Lawson sought to clarify at the outset was the story of why vitamin C treatments for cancer had not been initially supported by outside research. As noted in our previous posts, Pauling and Cameron believed that external research corroborating their results would provide an effective avenue for convincing skeptics of the curative powers of vitamin C. The Mayo Clinic eventually agreed to conduct a study of this sort, but found that vitamin C did not improve cancer patients’ prognoses, and in some instances their outcomes were actually worse.

Even though Pauling and Cameron did not agree with the Mayo Clinic’s findings, they had a hard time making their case to the public about why the study was in error. But as described by Lawson in the 2018 preface, Pauling and Cameron recognized that the Mayo Clinic’s treatment protocols were different than their own in important ways. Specifically, Cameron and Pauling had always delivered vitamin C intravenously, whereas the Mayo Clinic researchers dosed their patients orally. At the time, Pauling and Cameron could not prove why this should make a difference, but they believed that it was a key reason why the Mayo Clinic did not get positive results. While science could not confidently address this scenario in 1979 – or even in 1993 at the time of the second edition – by 2018 researchers had developed a much clearer idea of the importance of intravenous application.

Beginning in 1999, a team of researchers began actively exploring the absorption of vitamin C and, in particular, whether or not there was a difference between oral and intravenous applications. As they conducted their work, the group discovered a previously unknown vitamin C transport molecule in the stomach, which helped to deliver ascorbic acid into the bloodstream. The team subsequently learned that there was an upper limit to how much vitamin C the transport molecules could carry. This effectively meant that no matter how much vitamin C a person ingested orally – vitamin C that would end up in the stomach – only a finite amount could actually be absorbed and utilized, due to the limited carrying capacity of the transport molecule. The exact amount of vitamin C that a transport molecule can carry is still being researched, with data to this point indicating that quantities may vary based on factors such as age.

In tandem with this discovery, researchers were also interested in understanding what happens when vitamin C enters the bloodstream, be it intravenously or through transport molecules in the stomach. As Lawson notes, studies found that one byproduct of high levels of vitamin C in the bloodstream (regardless of how it got there) is hydrogen peroxide. Hydrogen peroxide has the ability to fight cancer by altering its DNA, robbing its cells of ATP (the “muscle” of the cell), and fatally damaging its energy-producing mitochondria. When deployed at scale, this three-pronged attack might be presumed to fight cancer very effectively.

That said, hydrogen peroxide therapy has not been used with cancer patients, because there is no safe and effective way to deliver the substance into the bloodstream without damaging other healthy cells. While the connection between vitamin C dosing and internal hydrogen peroxide production is not well-understood, these preliminary findings suggest that high blood concentrations of vitamin C could create hydrogen peroxide in sufficient quantity as to be effective at neutralizing cancer cells.


Several other new areas of research were highlighted in the 2018 edition, including the connection between vitamin C and hypoxia inducing factor (HIF). Some cancer tumors grow so fast that blood vessels cannot be created quickly enough to deliver oxygen to the expanding mass. To compensate for this lack of blood vessels in these hypoxic (or low oxygen) environments, tumor cells induce HIF, which stimulate the growth of blood vessels within the tumor. This inducement of HIF allows fast-growing tumors to continue to propagate and wreak havoc, rather than succumbing to oxygen starvation. For reasons that remain unclear, when vitamin C is introduced into a fast-growing tumor it represses the activation of HIF, compelling the tumor to remain in its hypoxic state and eventually die.

Emerging research on the relationship between vitamin C and dehydroascorbic acid (DHA) is also included in the 2018 edition. DHA is an oxidation product of vitamin C, and when present has been shown to reduce the amount of colorectal cancer in the body.

A collection of sixteen relatively recent clinical trials are likewise surveyed in the book, all of which examined the outcomes of different kinds of cancer patients when given vitamin C in conjunction with other chemotherapeutics. All sixteen found that patients’ outcomes improved once they were given the vitamin C in concurrence with their chemotherapy treatment. Five other studies were conducted on cancer patients given vitamin C but not chemotherapy. Results of these trials were mixed but, as Lawson points out, the non-chemotherapy researchers were primarily interested in determining optimal doses of vitamin C and measuring how quickly patients depleted their vitamin C infusions. Many of the trials also found that patient outcomes improved.

As with the original text of the book, a collection of case studies were also discussed by Lawson. And though they serve mostly as anecdotal evidence, Lawson shows that significant improvements in patients’ health have been documented once they have been administered vitamin C treatment.


One of the biggest takeaways that a reader might glean from the 2018 edition is that, even though a significant amount of research has been conducted on the vitamin C and cancer connection, medical practitioners have been hesitant to deploy the treatment because the studies have seemed less than rigorous, or because other practitioners have found that it does not work with their patients. Lawson counters these sentiments by noting that most cancer research is novel and that only the most promising ideas go on to clinical trials. Further, because an optimum vitamin C dose has not yet been codified, doctors commonly administer vitamin C on an ad hoc basis. Worse still, doctors also sometimes administer the treatment as a “last ditch” effort after chemotherapy and all other treatments have failed. Pauling and Cameron demonstrated in their first edition that vitamin C needs to be administered continuously and before chemotherapy to be maximally effective. Clearly much remains to be done before the work that Pauling and Cameron started with their first edition can be called complete.

Cancer and Vitamin C: A Major Conference

Pauling in 1989. Photo by Paolo M. Sutter.

[Part 8 of 9]

When the first edition of their book, Cancer and Vitamin C, was published in 1979, popular support for the curative powers of vitamin C did not materialize in quite the way that the authors, Linus Pauling and Ewan Cameron, had hoped. In part due to negative press associated with a critical study conducted by researchers at the Mayo Clinic, mainstream readers found themselves disinclined to buy into the thesis that vitamin C could fight cancer in the ways that Pauling and Cameron had put forth.

But over time, the conventional wisdom began to shift a bit, and in 1989 scientists from around the world convened in Bethesda, Maryland to once again discuss the merits of vitamin C as a cancer fighting agent. Sponsored by the National Institutes of Health (NIH), the conference was attended by respected scientific leaders, many of whom left the three-day event with a sense that the world might finally start to agree with Pauling and Cameron.


Held in April 1989, “Ascorbic Acid: Biological Functions and Relation to Cancer” marked the first time that two agencies within the NIH – the National Cancer Institute and the National Institute of Diabetes and Digestive Kidney Diseases – had co-sponsored an event. In total, about 130 people attended the meeting, with some 40 talks and papers presented over the three days.

Because of the breadth of its participants, the research presented at the conference covered multiple angles of the vitamin C and cancer connection. One speaker was Dr. Balz Frei from the University of California, Berkeley (and later the director of the Linus Pauling Institute) who discussed how peroxidation and prooxidant reactions can lead to cancer. Frei pointed out that these reactions happen frequently, and that behaviors like smoking can make the reactions even more prolific, which helps to explain how smoking can lead to cancer. Frei also found that these reactions did not occur when cells were exposed to vitamin C, but as soon as the vitamin C was removed, the reactions began again. Similarly, a team from Japan led by Dr. Etsuo Diki found that free radicals, when present in the body, can lead to cancer, and that vitamin C was capable of destroying them more quickly than could any other substance under study.


Even though much of the research was presented by scholars new to the field, many of the old believers were in attendance as well, including Pauling and several colleagues from his institute. Pauling’s talk focused on a recent study involving mice who were afflicted with cancer via ultraviolet light, and then treated solely with vitamin C. Pauling’s team found that, once treated, cancerous growths did not develop further and incidences of cancer reduced in general. Another similar study conducted by Pauling and his team on mice and spontaneous cancer in the mammary glands found that when the mice were given vitamin C, the time to onset of cancer was notably delayed.

Others presented similar types of work. One paper demonstrated that the effectiveness of chemotherapy improved when vitamin C was added to the drinking water of mice with cancer. And a team from Pennsylvania found that mice that were given vitamins C and B12 exhibited a complete inhibition of cancer growth with no damage to healthy, non-cancerous cells. The team’s results had been so encouraging that they treated a collection of human patients with vitamins C and B12, with positive results.

A different group found that, when vitamin C was given in conjunction with chemotherapy, patients tended to retain a statistically significant amount of healthy tissue as compared to those who were not given vitamin C. Other presentations found that vitamin C in cancer patients is not excreted in urine, despite drops in blood concentration, leading to the conclusion that the vitamin was being “used up” in fighting cancer. Vitamin C also seemed to reduce the toxicity of certain chemotherapeutics, including Adriamycin, a well-known cancer drug.


The conference ended with a talk from the macro perspective by Dr. Gladys Block, an epidemiologist at the National Cancer Institute. Block determined that, to date, there had been 47 studies conducted which demonstrated that vitamin C provided a measure of protection against cancer, and that 34 of these put forth data that was statistically significant. Block argued that if chance was the only reason why vitamin C had been found useful, then only one or two of the studies would have been statistically significant. But the fact that 34 studies met the threshold of statistics was both decidedly meaningful and encouraging. Block’s remarks closed the event, which concluded with “thunderous applause and a standing ovation.”