The Slow March Toward a Cure for Sickle Cell Disease

Pastel drawing of sickled hemoglobin cells by Roger Hayward, 1964

By Dr. Marcus Calkins

[Ed Note: This is the first of three posts examining the history of sickle cell treatment up to present day. It is authored by Marcus J. Calkins, Ph.D., “a proud OSU alumnus” (Chemical Engineering, B.S., 1999), who now works as a scientific communications service provider and educator in Taipei, Taiwan. In submitting this piece, Calkins emphasized that he has “taken inspiration from Linus Pauling’s research activities, teaching methods and moral character for many years.”]

In 2020, Jennifer Doudna and Emmanuelle Charpentier shared a Nobel Prize for their discovery and development of the CRISPR gene editor. One of the first clinical applications for CRISPR promises to be an ex vivo gene therapy for Sickle cell anemia. If it works, this medical technology will be a major breakthrough in biomedicine, representing the culmination of more than a century of research on Sickle cell disease that encompasses a wide range of topics.

Despite the lifetimes of work that have led to our current exciting position on the precipice of a cure for Sickle cell disease, the basic molecular features of the disease were defined seven decades ago by another Nobel Prize winner, Linus Pauling. The intervening 70 years of work have been required for scientists to learn how we might apply the foundational knowledge to actual patients in a real-life clinical setting. While the pace of progress may seem agonizingly slow to those outside biomedical research, the ground that has been covered is immense, and entire fields of biomedicine needed to be built and optimized before a truly feasible treatment technology could be invented.

Sickle Cell Disease (1910)

Sickle cell disease was first described over the period from 1910 to about 1924. During this time, a series of case reports detailed approximately 80 people of African descent, who had an odd morphology of red blood cells resembling a crescent or a sickle. In many cases, this sickle-like morphology was associated with a devastating condition involving severe anemia and early death. Furthermore, scientists learned that the red blood cell sickling could be exacerbated by depriving the blood of oxygen, either by adding carbon dioxide to cells in a dish or restricting blood flow in the patient. These clinical observations laid the foundation for basic scientists to postulate that the condition was related to hemoglobin, the protein that carries oxygen in red blood cells.

The first person to make this suggestion was Pauling. At some time in 1945, he was chatting with a colleague on the train from Denver to Chicago, when he learned about the difference in sickling between oxygenated and deoxygenated blood. According to the account of his colleague, Pauling was also informed that the sickled red blood cells show birefringence when viewed under a polarizing microscope, which would suggest an alignment of molecules within the cells.

However, by Pauling’s account of the conversation, his immediate guess that Sickle cell disease is caused by a defect in the hemoglobin protein complex was based entirely on the difference in the sickling properties of oxygenated and deoxygenated blood. Notably, Pauling later stated that the idea of Sickle cell disease being singularly caused by the hemoglobin molecule came to him in “two seconds,” but gathering evidence and refinement of the idea took at least three years.

In his public talks, Pauling often emphasized the fact that in the first years of the study, his students performed many experiments but could not identify any obvious biochemical differences between the hemoglobin molecules of patients and control individuals. From his repeated emphasis of this fact, one might speculate that the translation of a two-second idea to a three- or four-year demonstration would have been frustrating for such a quick-minded individual, though Pauling never said as much. Alternatively, he may have simply been emphasizing the challenges and slow, steady nature of rigorous scientific pursuit.

The Molecular Defect and a Potential Cure (1949)

In 1949, prior to the double helix model of DNA and before stem cells were described, the Pauling lab published a paper titled “Sickle Cell Anemia, a Molecular Disease.” In this work, Pauling and his students definitively showed that a slightly abnormal form of hemoglobin is found exclusively in patients with the cell sickling phenotype. Using a 30-foot-long Tiselius apparatus that they had constructed for electrophoresis, a small two-electron difference could be detected in the overall charge of hemoglobin molecules from Sickle cell patients and unaffected individuals. Meanwhile, carriers of the disease had a mixture of the two hemoglobin isoforms.

Importantly, Pauling’s group found that the defect in hemoglobin is not related with its ability to bind oxygen. Instead, it was later shown that the slight change in molecular charge affects the way hemoglobin proteins interact with each other, as would be predicted from the birefringence observation. This aberrant interaction causes the formation of long molecular scaffolds that change the shape of the red blood cell and lead to its dysfunction.

With this publication, Sickle cell disease became the first disorder to be associated with a single molecule. It was also the first with a known genetic basis. In his publication of the same year, J.V. Neel showed that Sickle cell disease follows an autosomal recessive inheritance pattern, meaning that each parent must contribute one copy of the mutated gene for a child to develop the disease. The cell sickling phenotype can occur to some degree in people who only carry one mutant allele, but only those with two copies experience the pernicious effects of the disease. This information, combined with Pauling’s study, established the essential basis for our understanding of Sickle cell disease and serves as a model for many other genetic diseases.

Surprisingly, James Watson (prior to his famed work on the structure of DNA) contributed a prescient idea to Sickle cell disease treatment, when he speculated that cells could be protected by expression of another form of hemoglobin, fetal hemoglobin. Watson made this prediction in 1948, just one year before Pauling’s powerhouse publication. His suspicion was an extension of reports that red blood cell sickling did not happen in the blood of infants who would later develop the condition as children and adults.

The stage was thus set for a Sickle cell disease cure. After the theoretical basis was determined, onlookers might have expected a cure for the disease to be found within a few years. However, extension of the ideas of Pauling and Watson has required incredible efforts by myriad scientists over the course of the next seven decades to create a potential new clinical reality.

Pauling’s Seventh Paper on the Nature of the Chemical Bond

[Part 7 of 7]

“The Nature of the Chemical Bond. VII. The Calculation of Resonance Energy in
Conjugated Systems.” The Journal of Chemical Physics, October 1933

The final paper in Linus Pauling’s earthshaking series on the nature of the chemical bond was the shortest of the seven and made less of a splash than had most of its predecessors. This lesser impact was anticipated and was due primarily to the guiding purpose of the paper: to apply previously developed postulates to compounds that had not been addressed by Pauling in his prior writings. As with the sixth paper in the series, the final publication was co-authored by Caltech colleague Jack Sherman.

In paper seven, Pauling demonstrated how to calculate resonance energy in conjugated systems. A conjugated system is one in which there exists a plane – or alignment – of three or more connecting electrons located in the p orbital. While it was commonly understood by the era’s organic chemists that conjugated systems supplied a compound with more stability than would ordinarily be expected, Pauling’s paper offered the calculations needed to codify this knowledge.

The paper also put forth a collection of rules to help researchers better understand the properties of conjugated systems. For example, Pauling found that “a phenyl group is 20 or 30 percent less effective in conjugation than a double bond, and a naphthyl group is less effective than a phenyl group.” To arrive at these conclusions, Pauling used the equations that he had developed in his previous two papers, applying them this time around to conjugated systems.


Jack Sherman and Linus Pauling, 1935.

Pauling’s seven papers on the nature of the chemical bond came to print over the course of thirty months, from article one in April 1931 to article seven in October 1933. The first three papers laid the groundwork for what was to come by defining chemical bonds in quantum mechanical terms. The fourth paper, published in September 1932, appeared at the midpoint of Pauling’s publishing chronology and also served as a kind of transition paper, connecting the concepts introduced in the first three publications to those in the three more that were forthcoming. (Paper four also contained Pauling’s vital electronegativity scale.) The last three articles were devoted to the concept of resonance and its application to a fuller understanding of the chemical bond.

Taken as a whole, this body of work proved hugely important to the future direction of chemistry. By reconciling and applying the principles of quantum mechanics to the world of chemistry, the articles showed that what had once been mostly a tool for physicists could indeed have great applicability to chemical research. In the process, Pauling and his collaborators also rendered quantum mechanics far more accessible to their colleagues across the field of chemistry. The end result was, to quote Pauling himself, “a way of thinking that might not have been introduced by anyone else, at least not for quite a while.”


This is our forty-eighth and final post for 2020. We’ll look forward to seeing you again in early January!

Pauling’s Sixth Paper on the Nature of the Chemical Bond

Table of resonance energy calculations for condensed ring systems

[Part 6 of 7]

“The Nature of the Chemical Bond. VI. The Calculation from Thermochemical Data of the Energy of Resonance of Molecules Among Several Electronic Structures.” The Journal of Chemical Physics, July 1933.

In paper number five in his Nature of the Chemical Bond series, Linus Pauling argued that the theory of resonance could be used to accurately discern the structure of many compounds, and he used Valence Bond theory to substantiate that claim. However, much of the argumentation put forth in the paper relied upon fairly generalized calculations, some of which were subsequently shown to be in error.

In his sixth paper, published one month later, Pauling put forth more definitive calculations that used thermochemical data that were more empirically based, and therefore less prone to errors. As with the previous publication, this paper was co-authored. However, instead of G.W. Wheland, Pauling’s collaborator this time around was Jack Sherman, a theoretical chemist who had received his PhD from Caltech the year before.


The data used in the paper weren’t anything new; in fact, they had been used by chemists for years to calculate energy values and to determine bond energies. However, in many cases these calculations failed because chemists, who were rooted in classic organic or physical models, always assumed that the molecules under study were consistently similar.

Relying instead on a quantum mechanical approach, Pauling and Sherman argued that compounds could (and should) be organized into two broad categories. In one group, there resided those molecules that were well-approximated by their Lewis structures (classical representations of molecules using lines and dots to represent bonds and electrons). The other group consisted of compounds whose structures could only be accurately explained through resonance.

By organizing compounds into these two discrete bins, Pauling and Sherman were then able to make more accurate calculations of bond energies. More specifically, the duo was able to calculate energies of formation for various molecules by using extant experimental data on heats of combustion. Pretty quickly they realized that energies of formation could be accurately calculated for Lewis structure (non-resonating) compounds.

For resonating compounds however, the tandem found that calculated energies of formation were much higher than what would have been predicted by theory. Higher energies of formation yield more stable molecules, and the co-authors concluded that the “difference in energy is interpreted as the resonance energy of the molecule among several electronic structures” and that “in this way, the existence of resonance is shown for many molecules.”

Pauling’s Fifth Paper on the Nature of the Chemical Bond

[What follows is Part 5 of 7 in this series. It is also the 800th blog post published by the Pauling Blog.]

The Nature of the Chemical Bond. V. The Quantum-Mechanical Calculation of the Resonance Energy of Benzene and Naphthalene and the Hydrocarbon Free Radicals.” The Journal of Chemical Physics, June 1933.

With his fifth paper in the nature of the chemical bond series, Linus Pauling communicated a new understanding of the structures of benzene and naphthalene. While it had been long accepted that benzene (C6H6) was arranged as a six-carbon ring and naphthalene (C10H8) as two six-carbon rings, the specific organization of electrons and bonds within these structures were not known. Before the publication of Pauling’s fifth paper, several ideas on these matters had been proposed, but all were viewed as flawed in some way or another. But where others had been stymied, Pauling found success, and he did so by fully embracing and utilizing the theory of resonance.


At the time that Pauling began this work, there were five competing structures for benzene, each burdened by its own problems. The one that was the most accepted, despite its inability to connect theory to experimental data, was the Kekulé model. Put forth several decades earlier by the German chemist August Kekulé, this model centered around a six-carbon ring that possessed alternating double bonds. Because the arrangement of these double bonds could differ, Kekulé’s model was actually proposing two potential isomers for benzene. The standard understanding at the time was that these two isomers constantly oscillated between one another.

One major problem with the Kekulé approach was that scientists of his generation had never found evidence of the oscillating structures. Furthermore, the Kekulé structures should have been quite unstable, which was contrary to what researchers were able to observe in the laboratory. As such, even though it was compelling in the abstract, the Kekulé model was known to be imperfect.

In his paper, Pauling pointed out the flaws in Kekulé’s work as well as four other concepts published by other researchers. In doing so, he suggested that a common hindrance to all of the approaches was a reliance upon the laws of classical organic chemistry, and a concomitant lack of application of the new quantum mechanics. It was Pauling’s belief that the structure of benzene could be explained using quantum mechanics, as could the structures of all aromatic compounds.


In a handful of previous papers, Pauling had used the theory of resonance to explain a variety of chemical phenomena, but in thinking about benzene and naphthalene he committed more fully to its principles. According to Pauling, all observable data that had been collected for benzene, particularly its bond energies, suggested that benzene was much stronger than any models had yet to predict. But none of the previous models had entertained the possibility of a resonate structure, by which he meant an aggregate structure that was essentially a blend of all possible structures. A structure of this sort, Pauling argued, would conform to a lower, more stable energy state, and would accurately map with the observed data.

For Pauling, therefore, the structure of benzene was not the result of rapid isomerization as put forth by Kekulé, but rather a blend of states. “In a sense,” he wrote, “it may be said that all structures based on a plane hexagonal arrangement of the atoms – Kekulé, Dewar, Claus, etc. – play a part” but “it is the resonance among these structures which imparts to the molecules its peculiar aromatic properties.”

To support his theory, Pauling considered all five possible structures of benzene – which he called “canonical forms” – calculating the energy of each structure as well as the combined resonance energy. Having done so, Pauling then noted that it was the resonance energy that most closely matched the observed data.


In addition to its utility, the elegance of Pauling’s approach compared favorably with similar work being published by a contemporary, the German chemist Erich Hückel. Situating this thinking within Molecular Orbital theory, Hückel was able to arrive at a similar conclusion for benzene, but his calculations were quite cumbersome and could not be applied to larger aromatic compounds. By contrast, Pauling was now firmly rooted in Valence Bond theory and his formulae could be applied to all aromatics, not just benzene. In particular, by simplifying some of the calculations that Hückel had made, Pauling was able to overcome some of the mathematical hurdles posed by the free radicals in benzene and other aromatics.

To demonstrate the broad applicability of his ideas, Pauling applied his theoretical framework to naphthalene, which consists of two six-carbon rings and had forty-two canonical structures — a great many more than benzene’s five. Despite this significant difference, Pauling was successful in applying the same basic math to determine that the structure was also in resonance.

Indeed, Pauling was certain that his calculations were relevant to all aromatic compounds, noting specifically that “this treatment could be applied to anthracene [a three-ringed carbon molecule] and phenanthrene [a four-ringed carbon molecule], with 429 linearly independent structures, and to still larger condensed systems, though not without considerable labor.” Were one willing to expend this labor, the calculations would show that the “resonance energy and the number of benzene rings in the molecule would be substantiated” and the structure correctly predicted.


G.W. Wheland

The fifth paper was unique in part because it was the first in the series to be co-authored. The article also marked a switch in publishing forum: whereas the first four had appeared in The Journal of the American Chemical Society, this paper (and the two more still to come) was published in volume 1 of The Journal of Chemical Physics.

Pauling’s co-author for the paper was George W. Wheland, a recent doctoral graduate from Harvard who worked with Pauling from 1932-1936 with the support of a National Research Fellowship. This collaboration proved noteworthy both for the quality of the work that was produced and also because Wheland later became a vocal supporter, advocate and contributor to resonance theory.

Pauling’s Fourth Paper on the Nature of the Chemical Bond

[Part 4 of 7]

“The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms.” Journal of the American Chemical Society, September 1932.

The first three papers published by Linus Pauling in his nature of the chemical bond series were all novel, and the first paper in particular made a significant impact. But it is the fourth paper that has proven to be perhaps the most influential of all. In it, Pauling introduced his idea of the electronegativity scale, a cohesive and logical tool that proved to be of major import to the discipline.


The concept of electronegativity can be understood in terms of the likelihood that an atom will attract a pair of valence (or bonding) electrons. The more electronegative an atom is, the more likely that it will attract electrons. The most electronegative element is fluorine and the least electronegative element is francium.

Pauling was able to develop a scale for electronegativity using insights into valence bond energies. The ideas that he had put forth in his three preceding papers did not always firmly commit to either Molecular Orbital theory or Valence Bond theory, and this vacillation led to significant flaws in paper number two. Beginning with the fourth paper however, Pauling chose to base his work on the tenants of Valence Bond theory.

The electronegativity scale that Pauling developed was also quite intuitive in that it could not be calculated directly, but instead had to be inferred using the atomic and molecular properties of a given element. And even though its creation relied upon a series of assumptions and simplifications, the tool was nonetheless quite sophisticated. Prior to Pauling’s publication of his electronegativity scale, chemists either had to rely upon their own best guesses to determine bond affinities or, if they wanted more precision, compute bond energies for every interaction. Pauling’s scale both standardized and simplified these processes while creating a context where chemists could make predictions of “the energies of bonds for which no experimental data are available.”

Clearly one key piece of utility that the electronegativity scale provided was the ability for chemists to draw conclusions without the need for a lot of computation. For example, based on a molecule’s electronegativity, chemists could roughly deduce the ionic nature of a bond. However, in order to predict the ionic character of a bond, Pauling had first needed to make some assumptions, such as creating an “arbitrarily chosen starting point” to which everything is relative. Even though this approach was not precise, Pauling argued that its usefulness justified the simplification of the approach. And naturally, over time, Pauling’s electronegativity scale became more honed and more specific, an evolution that Pauling also predicted in the fourth paper.


Most of Pauling’s article was devoted to a discussion of how he developed the scale and what kinds of measurements were used in the calculations. To begin, he focused on the qualities of covalent attraction or repulsion in compounds formed by identical elements, such as H:H or Cl:Cl. From there, Pauling used quantum mechanical wave function properties to establish that “energies of normal covalent bonds are additive.” It was from this central theorem that Pauling built out the rest of his conceptual work and also his scale, predicting, for example, bond energies for light atoms and halogens.

Pauling also relied upon others’ helpful calculations in creating his scale, and especially those related to heats of “formation and combustion of gaseous materials” put forth in the International Critical Tables (National Research Council) and elsewhere. The heats of formation were especially useful because they helped to correct for any unknown bond energies. In fact, by using these experimental data for twenty-one different single bond energies, Pauling was able to derive a formula that predicted where a given element should reside on the electronegativity scale. For bonds where data were not available, Pauling used his predictive models to extrapolate approximate energies.


The formula that Pauling derived and published was Δa:b=(χAB)2 where χA and χB represent the coordinates of each atom A and B on an electronegativity map (see Fig. 5 above), and Δ represents the degree of electronegativity.

Here, for example, is how Pauling calculated the electronegativity value for oxygen. He began by establishing the heat of formation of water:

2H + O=H2O(g) +9.493 v.e.

Then, based on the fact that the H:O bond is found to have a bond energy of 4.747 v.e., Pauling further calculated that:

H2O2(l)=H2O(l) + ½ O2 + 1.02 v.e.

Pauling next combined this value with the heat of vaporization of H2O2, which is .50 v.e., and found that:

2H + 2O = H2O2(g) +10.99 v.e.

From there, using his original postulate about consistent bond energies, Pauling subtracted 4.75 v.e. for each H:O bond to yield 1.49 v.e. for the O:O bond. However, since Pauling had previously concluded in his previous papers that O:O was a double bond, the actual electronegativity for the oxygen molecule was found to be 3.47. (The present-day electronegativity value for oxygen has been revised to 3.44.)


Even though it relied in part on a collection of assumptions and simplifications, Pauling’s electronegativity scale has been widely used and stands as a lasting component of his legacy. In addition to its ability to approximate values for a wide variety of compounds, the scale was also important for establishing the idea that electronegativity is not a fixed number that never changes. Instead, Pauling understood that an element’s electronegativity value emerges from its bonding relationships. Because of this, calculating absolute values has been difficult, but Pauling’s scale continues to be useful and predictive.

Pauling’s Third Paper on the Nature of the Chemical Bond

[Part 3 of 7]

“The Nature of the Chemical Bond. III. The Transition from One Extreme Bond Type to Another.” Journal of the American Chemical Society, March 1932.

In his third paper exploring the nature of the chemical bond, Linus Pauling dug into the unsolved question of how molecules transition from one kind of bond type to another. While it had been determined that molecules do switch from one kind of bond to another – from an ionic bond to an electron-pair bond, for example – the specifics of how that transition happens remained elusive.

Prior to the third paper, two prevailing ideas were being debated by chemists. One concept, as Pauling wrote, was that “all intermediate bond types between the pure iconic bond and the pure electron-pair bond” exist in some kind of infinite transitionary state. A contrary viewpoint put forth instead that molecules “transition from one extreme bond type to another” in an abrupt manner. Pauling suggested that the answer lie somewhere in between.


In order to determine how molecules transition, Pauling first needed to establish the bond structures of given molecules in their initial states. He did so by defining the bonding characteristics of molecules, a task that takes up the majority of the paper. But amidst this discussion, Pauling arrived at several key conclusions.

To begin, Pauling described many cases where a relationship existed between atomic arrangement – as determined by x-ray crystallographic analysis – and bond energies. When, for instance, a strongly electropositive and strongly electronegative molecule bonded, it was reasonable to assume that the bond was ionic. This presumed, Pauling then used electron energy curves to show that an example group, the alkali halide molecules, were strongly ionic, and that they might generally be thought to form ionic bonds.

As Pauling pointed out however, these presumptions were faulty. In fact, studies of the bonding in hydrochloric acid (HCl) and hydrobromic acid (HBr) indicated that both molecules were essentially covalent in make-up, whereas hydrofluoric acid (HF) was ionic. So even though it might reasonably have been assumed that the initial states for HCl, HBr and HF would be similar in their bonding, the experimental data indicated otherwise. These findings led Pauling toward the conclusion that there is no single universal answer to the question of how molecules transition, because there is no steadfast rule determining the types of bonds that hold molecules together before they transition.  


Having arrived at the conclusion that one could not lean upon a guaranteed universal bond type, Pauling then turned to his burgeoning theory of resonance to develop more precise thinking about transition mechanisms. Pauling specifically argued that when bonds transition from one type to another, rather than shifting either abruptly or in a continuous state – as the two competing models then prevailing put forth – they instead shift to an intermediate resonant state before switching to a new bond type.

Pauling was in essence suggesting that, in between classically-defined “completed” bond states, there also existed an intermediate bonding state that could best be understood through the theory of resonance. Moreover, Pauling argued that idealized bonds, such as pure covalent bonds or pure ionic bonds, did not technically exist. Rather, bonds might more accurately be described as constantly transitioning through resonant states, some of which more closely approximated a classic bond type.

Pauling understood that the concept he was putting forth was quite theoretical and that, in practical terms, it was hard to work with molecules if they existed in a constant state of transition. As such, Pauling allowed that, for purposes of discussion, it was acceptable to think of molecules as residing in discrete bonding states. He likewise acknowledged the convenience of using more traditional names (ionic, covalent, etc.) when referring to bonds, even if they never fully existed.

Pauling then concluded that, even though bonds were constantly transitioning, for certain bond types – such as “when the normal states for the two extremes have the same number of unpaired electrons” – it could be assumed that they had transitioned in a continuous state. But, of course, continuous state transition was definitely not always the case and could not be universally applied.


In conducting the work that led to his third paper, Pauling had sought to define a universal rule that would govern the transition between bond types. By the time that he delivered his manuscript though, he had recognized that not only was a universal law unattainable, but that what he did find had its limitations. In particular, Pauling suggested of his approach that “It is not possible at the present time to carry out similar calculations for more complicated molecules,” though “certain less specific conclusions can, however, be drawn.”

Regardless, Pauling’s third paper broke new ground on a topic of keen importance to structural chemists. By applying the theory of resonance, Pauling helped chemists to understand that there was a spectrum of polarity, and that bonds were not always strictly of one kind or the other. Importantly, in this same paper Pauling did not fall prey to dogmatism, and allowed that bonds residing near the ends of one spectrum or another might fairly be said to represent so-called “classic” bond types.

Pauling’s Second Paper on the Nature of the Chemical Bond

Linus Pauling, 1931.

[Part 2 of 7]

“The nature of the chemical bond. II. The one-electron bond and the three-electron bond.” Journal of the American Chemical Society, September 1931.

Linus Pauling’s first paper on the nature of the chemical bond made huge waves throughout the field, catching the attention of many. The second paper in the series however, was not quite so memorable. In it, Pauling tried to use quantum mechanics to explain molecules formed by either one- or three-electron bonds. This work stood in contrast to his first paper, which focused solely on electron pair bonds. And though they were similarly novel to Pauling’s first article, the ideas put forth in the second paper did not stand the test of time.


The single-electron and three-electron bonding environments had long proven difficult for chemists to understand. In tackling the challenge, Pauling admitted at the outset that they, “have not the importance of the electron-pair bond, for they occur in only a few compounds.” Nonetheless, they were “of special interest on account of their unusual and previously puzzling properties.”

Pauling’s first step in approaching the problem was to focus first on single-electron bonding, after which he would then apply the rules and mechanisms that he had formulated to the three-electron bond. Like all of the papers in his nature of the chemical bond series, the foundational groundwork for his ideas was based in quantum mechanics.

Accordingly, he began his explanation of the one-electron bond by delving into quantum mechanical principles. In doing so, Pauling argued that resonance – which, simply stated, is the sharing of energies, and a concept that Pauling would detail in greater depth in his fifth paper – could be predicted to arise if two different theoretical states of a bond possessed equal energy.

For example, the two possible bonding states for H2+ were H· H+ and H+ H·. Given that the two configurations possessed equal energies, Pauling put forth that resonating bonds between the two would be found with “essentially the same energy.” In other words, the one-electron bonds holding H2+ together were actually resonating between two different configurations. The mechanisms underlying three-electron bonds were detailed in much the same way.


After offering this explanation for how these bonds might form, Pauling used the rest of his paper to apply the mechanism to known molecules. In doing so, Pauling noted that, to date, it had been difficult to determine which molecules even possessed one- or three-electron bonds. Both types of bonds were known to exist, but this was mostly because observed energy levels for certain compounds were much more stable than would be predicted by models using other types of bonds.

For compounds suspected to be of this type, Pauling used resonance to demonstrate equivalence between observed bond energies and theoretical bond energies. Success in showing that the two sets of bond energies were equal or “differ[ing] by only one or two volt-electrons” would mean that the “criterion for the formation of a one-electron [or three-electron] bond is satisfied,” thus validating his ideas.

In his paper, Pauling applied his theory specifically to the boron hydrides. Once again based upon compatibilities between observed and predicted energy values for these compounds, Pauling found his criteria to be met and concluded that “the one-electron bond is to be expected.” In later years, as Pauling’s thinking matured, his views on these compounds became more sophisticated. But at the time of the second paper, he argued that the hydrogen ion (H+3), and the triatomic hydrogen ion (H+3) in the boron hydrides must be single-electron bonds based on their observed energy levels.

Pauling likewise suggested that the oxygen molecule, the helium ion (He+2), and nitric acid (NO), were all formed by three-electron bonds. For oxygen in particular, Pauling believed that a three-electron bond was required to explain the molecule’s slight magnetic moment. This connection between magnetism and electron bonding was something that Pauling had addressed in his first paper, and remained a topic to which he would turn with some frequency throughout the seven paper series.


As time moved forward, many of the notions put forth in Pauling’s second paper were largely debunked, with Pauling himself eventually speaking out about their irregularities. The primary issue with the work was its reliance upon Molecular Orbital theory to explain how double and triple bonds work. At this early stage in Pauling’s career, Molecular Orbital theory and Valence Bond theory did not stand apart as distinct models and, as such, Pauling’s thinking sometimes blended the two. Later, of course, Pauling’s writings would prove foundational to the advancement of Valence Bond theory.

Despite their limitations, the concepts introduced by the second paper were still important in helping Pauling to frame the larger thrust of his series. Though he sometimes went astray, Pauling more commonly found success in building off of concepts from his preceding papers in his work to create a fundamental understanding of the nature of the chemical bond.

Pauling’s First Paper on the Nature of the Chemical Bond

[Part 1 of 7]

Over the span of just two short years beginning in 1931, Linus Pauling published seven decidedly influential papers on the nature of the chemical bond. In the series, which formed the foundation for his 1939 book, The Nature of the Chemical Bond, Pauling introduced many chemists to the burgeoning field of quantum mechanics and demonstrated its applicability to structural chemistry. As a result of this work, Pauling was awarded the 1954 Nobel Prize in chemistry, “for his research into the nature of the chemical bond and its application to the elucidation of the structure of complex substances.” With today’s post, we begin a series that attempts to explore the scientific advancements and ideas put forth in each of the seven papers.


The nature of the chemical bond. Application of results obtained from the quantum mechanics and from a theory of paramagnetic susceptibility to the structure of molecules.” Journal of the American Chemical Society, April 1931.

The first paper in the series was perhaps the most important, in that it set the theoretical foundation for the six papers – some of them more practical than theoretical – yet to come. In it, Pauling sought to lay the groundwork for ideas that he would expand upon in future articles, notably pointing out that “many more results of chemical significance can be obtained from the quantum mechanical equations.” A statement of this sort was necessary because, up until that point, quantum mechanics had mostly been seen as a tool for physicists to use within their discipline. In order to convince chemists of the usefulness of quantum mechanics to their own work, Pauling put forth a conceptual framework consisting of a series of rules as well as guidance on how to apply them to molecular orbitals.

Three of the six rules that Pauling stressed in his paper were already well-known to chemists, but were needed in order to generate “buy-in” for the three additional rules that were new to the field. The most important of these was the fifth rule, which stipulated that when two electrons are in the process of forming a bond, the electron with the larger eigenfunction value will dictate the direction and shape of the bonds that are subsequently created. Pauling argued that this type of electron pairing would result in the most stable bonds possible.

Importantly, Pauling’s rules also combined ideas about wave functions with quantum mechanical thinking in a way that pushed the field forward. Earlier iterations of quantum ideas were not holding up well to certain experimental data. For example, it was known that electrons occupied quanta states, but the observations of the energy associated with these states did not always match what might be predicted by quantum theory. Through his study of both wave theory and quantum mechanics, Pauling recognized that shared bonding energies could explain certain observed bond energies and angles. As such, with his rules, Pauling helped chemists to merge quantum mechanics and wave functions, in the process creating a model that was predictive for all molecules.


In addition to conceptual rules for electron-pair bonding, the first article also helped to establish rules regarding the splitting or breaking of molecular orbitals, a concept that was completely new. The traditionally held belief was that orbitals were firmly set, but Pauling believed that if this were not true, then a whole new set of tantalizing possibilities was on the table. (Indeed, Pauling’s idea that orbitals could be split or broken formed the germ of his later work on the theory of resonance, which proved hugely influential.)

In this section of his paper, Pauling’s main focus was the s and p orbitals. Prior to his article, these orbitals were thought to occupy discrete quantum states that could not be broken. However, based on his earlier assertions that combined wave functions and quantum mechanics, Pauling argued that quantized states could, in fact, sometimes be broken. In subsequent writings, Pauling would conclude that the hybridization of orbitals allowed for the breaking of quantized states, but in this earlier phase of his thinking, he instead put forth that the orbitals were simply broken.

This was, of course, a big theoretical leap, but throughout his series of seven papers, Pauling often chose to base theory on informed guesses, even if he did not always have a precise understanding of how the rules worked. In this particular instance, Pauling used his fifth rule to rationalize the idea of breaking orbitals, positing that, when bonding, the stronger of two electrons would force the weaker to overlap and ultimately create new bond angles.


Though he did not yet fully understand all the minutiae of how it could be possible for orbitals to break, Pauling did know that the rules governing chemistry at the time were not sufficient, since they were unable to explain why molecules were sometimes observed to be more stable than predicted. By leaning on the notion that orbitals could break, Pauling was able to devise a set of rules that overlapped almost perfectly with the experimental data. In his paper, Pauling spoke particularly of the tetrahedral carbon atom,

in which only the s and p eigenfunctions contribute to bond formation and in which the quantization in the polar coordinates is broken can form one, two, three, or four equivalent bonds, which are directed towards the corners of a regular tetrahedron.

The idea of a tetrahedral angle is well-known within structural chemistry today, but it was a novel concept in 1931. As such, with his first paper, Pauling was not only proposing that orbitals could do things previously unheard of (i.e. break), but that they also formed angles that were completely new. Pauling knew that these ideas were revolutionary and devoted a significant component of his article to describing the tetrahedral angle in detail.


The s and p orbitals that Pauling addressed were important to many chemists because they formed the building blocks of carbon atoms. Analysis of larger orbitals however, such as the d orbitals, was often kept separate from discussions of the s and p orbitals, because the chemistry of the time lacked a unified law that could apply to both smaller and larger orbitals alike.

Pauling combated this by explaining that his six rules were constant, and that they could be applied to all scenarios, not just the s and p orbitals. From there he reasoned that, while more complicated molecules might open up additional possibilities for bond angles, the proclivity for bonds to form tetrahedral angles still applied. Pauling further argued that this idea was supported by the experimental data. In the case of cobalt for example, Pauling noted that both the predicted and observed angles are six equivalent sp3d2 bonds, and that because of his fifth rule, the bonds are pulled towards the corners of the octahedron that is formed, rather than the center.


As if that weren’t enough, Pauling also addressed magnetization in his paper, a new concept that had long fascinated him. Even during his years as an undergraduate student at Oregon Agriculture College, Pauling was deeply interested in understanding how it was possible that certain compounds were magnetic, while others were not. What, Pauling asked himself, was causing the difference?

In the first paper, Pauling does not quite offer an answer, but he did lay out a series of observations that would lead to new insights on electronegativity, the subject of his fourth paper in the series. In paper one, Pauling specifically contended that, since unpaired electrons were fundamental to magnetic compounds, a model of the bond types that comprise a given molecule could be built using data on the magnetic properties of the molecule. Offering the transition group elements as an example, Pauling pointed out that, without exception, they pair with CN  to form electron-pair bonds; with F to form ionic bonds; and with H2O to form ion-dipole bonds.


Later in life, Pauling reflected that his first paper on the nature of the chemical bond was “the best work I’ve ever done,” and indeed it is difficult to overstate the importance of the publication. In a single article, Pauling was able to put forth crucial new ideas on bonding in both simple and complex molecular structures using a standardized set of rules. The paper also began the process of applying the new quantum mechanics to help explain the structure of molecules in ways that better supported experimental observations. And while Pauling was later criticized by some for the assumptions that he had made, the value of the paper increasingly shone through as chemists came to understand the practical utility of the work to their own research.

Pauling’s OAC: A New Decade

[Looking back on Linus Pauling’s junior year at Oregon Agricultural College. This is part 3 of 3.]

The 1920-21 academic year at Oregon Agricultural College was, in many respects, a period of growth and change. With World War I now concluded, the school grew in size, many of its programs became better known, and the campus was buoyed by a sense of optimism.

One signal that the Roaring Twenties had reached Corvallis was the aptly named “Inter-Fraternity Smoker” contest. The cheeky competition – in which Linus Pauling participated – saw fraternity members “v[ying] with one another to produce the best characterization of womankind.”

Sports at OAC also enjoyed growth and success during the school year. Football, basketball, baseball and track all competed without the encumbrances mandated by the flu years, a new tennis program came into being, and OAC’s wrestling team won an intercollegiate championship. The college likewise sponsored several sports-adjacent groups, including a varsity “yell team” that was comprised of three students who rooted especially hard for the Beavers at all athletic events. “Women’s aesthetic dancing” also made an appearance during the year.

At this point in its history, OAC was committed to providing physical education opportunities for its entire matriculated population. All women enrolled at OAC were required to take a P.E. course, and the Physical Education for Men curriculum was among the most robust on the West Coast.

P.E. opportunities during the year were augmented by the completion of a new campus pool, chronicled as a welcome relief to students who had, according to The O.A.C. Alumnus, “swelter[ed]” through an “unprecedented heat wave.” The 50 x 100 foot facility, 9.5 feet deep at the diving well end, was among the larger pools around, and was surrounded by a grandstand that could seat 5,000 people.


The 1920-21 school year was also important for OAC’s academic programs, many of which were garnering more attention. The schools of Agriculture and Engineering could both boast of national rankings, while the School of Home Economics was the third largest in the country, featuring four departments (household science, household art, household administration, and home economics education) which all led to a four-year degree. The School of Forestry, which focused on logging engineering and technical forestry, was among the biggest in the country, and the School of Pharmacy – which offered pre-medicine courses in addition to a more standard pharmacy curriculum – was a leader on the West Coast. OAC did not house a School of the Liberal Arts, but its music program was, on the basis of size, ranked second among western colleges.

An OAC Chemistry lab, circa 1920

While not of national stature, OAC’s Chemistry program was quite large, owing to the fact that all students were required to take at least introductory chemistry as part of their studies. The Physics department created a research branch during the school year, and students in Botany and Plant Pathology could make use of the largest library of plant diseases and plant equipment west of the Mississippi River. Two new degree programs were added in Fall 1920, a B.S. in Vocational Education and a B.S. in Military Science and Tactics, “for the training of men for appointment as officers in the Regular Army.”

The school year also saw significant additions to the built environment on campus. The Health Services department moved into new quarters, where one full-time physician and two nurses cared for the OAC community. The School of Commerce broke ground on a new building, as did the School of Home Economics, and the School of Engineering moved forward on a series of projects with an estimated value of $360,000 (approximately $4.7 million today). Notably, present-day Kearney Hall – then known as Apperson Hall and home to the School of Engineering – received an extensive addition.

The 1920-21 academic year was also the last tuition-free year for OAC students. In previous years, students had been charged fees to help support specific classes, but they were not assessed a separate tuition charge.

Additional costs continued to come about for lodging. As before, all female students were required to live in the dorms, unless they received special permission from the dean or their parents lived in Corvallis. OAC’s women’s dorms featured “large air parlors” and cost $18 per term for a single or $9 for a double, plus $5 for deposit and incidentals. All rooms had access to “pure mountain water, both hot and cold,” lights, heat and “other modern conveniences” including a bed, pillows, linens, towels, sheets, and a wardrobe. Starting a few years prior, men also had the option to live in dorms. Their rooms were outfitted for between two to six people, but private rooms could be found around town for $4.00 – 5.50 per week, including meals.


Despite all of the positive momentum, many at OAC recognized that it was still hard for recent graduates to find work outside of the Pacific Northwest. (except for those from the School of Agriculture, who had “no such hardship.”) Some argued that a reason for this was the name of the school, and that if O.A.C. were to be rechristened as Oregon State College, a “handicap” that was “neither fair nor equitable” to graduates outside of the School of Agriculture would be removed. As it happened, the school was eventually renamed, but not until 1927 and even then as Oregon State Agricultural College. OSAC became Oregon State College in 1932 and, at long last, Oregon State University in 1961.

Pauling’s OAC: A Maturing Relationship with Chemistry

Linus Pauling, 1920.

[A look back at Linus Pauling’s undergraduate experience from 100 years ago; part 2 of 3.]

By the fall of 1920, Linus Pauling was connected to an academic trajectory that he would continue to pursue for the rest of his life. That said, during his years at Oregon Agricultural College, he was compelled to advance his studies in chemistry through rather unorthodox means. Because OAC was a land grant institution, the practical and applied sciences were the main point of emphasis within the college’s curriculum. Further, because the state of Oregon discouraged (and later mandated against) redundancy in the majors offered by its two largest institutions of higher learning, and because the University of Oregon already offered a degree in chemistry, Pauling’s only real option as a Beaver was to major in chemical engineering.

Partly as a result of these circumstances, much of the chemistry that Pauling had learned so far was fairly out of date. Not surprisingly, Pauling had found many of his classes to be dull and, at times, rote in their emphasis on solving problems of interest to engineers rather than academic chemists. But by the fall of 1920, having spent the previous year teaching, Pauling re-enrolled at OAC with a boost in confidence and a willingness to seek out opportunities in non-traditional ways. Fortunately, the school year reciprocated, offering key new acquaintances who broadened horizons for the precocious young student.


Throughout his studies in chemistry, the young Pauling often found himself questioning aspects of what he was learning and seeking to uncover more. For example, Pauling was intrigued by magnetism and puzzled over questions of why certain materials with similar physical structures varied in their degree of attraction to one another.

The courses that Pauling had taken to date were not providing answers to these questions. As a chemical engineer in training, he was learning that different substances expressed different levels of magnetism, but he had no insight into why. Prior to his junior year, Pauling may well have been resigned to the notion that these were unanswerable questions. However, more satisfactory solutions soon emerged with the help of a few influential professors.

OAC alumni inducted into Phi Kappa Phi, 1924. John Fulton stands in the back row, second from right.

Though he had saved up enough money to return to school, Pauling still needed to earn a wage to pay for on-going expenses, so he took up a job as an assistant to OAC Chemistry Professor Samuel Graf. Even though the job consisted mostly of working through computations, it also allocated time for Pauling to engage with the scientific literature. OAC’s Chemistry head, John Fulton, helped facilitate this by giving Pauling a few of his own chemical journals, and during his stint as Graf’s assistant, Pauling began to consume these journals with relish.

It was in this setting that Pauling first encountered the work of G.N. Lewis and Irving Langmuir, both of whom were exploring some of the most exciting questions in subatomic chemistry. While their publications did not answer all of Pauling’s questions, (many of which were in their earliest stages of formation) reading Lewis and Langmuir made Pauling realize that this new field of subatomic chemistry could solve problems, many of which he had not even realized existed.


While the history of the field of subatomic chemistry is quite complex, many of the ideas that Lewis and Langmuir were developing emerged because of headways that the Danish chemist, Niels Bohr, made with the formalization of his quantum theory in 1918. At OAC all of the chemical engineering courses were physical and practical in their orientation. The kind of theoretical work that Bohr, Lewis, and Langmuir were doing was novel – and not being taught at OAC – but making its acquaintance equipped Pauling with new tools to explore some of the questions that he was pondering as a nineteen-year-old undergraduate. This breakthrough renewed Pauling’s fervor for chemistry and his determination to pursue it for a career.

Pauling’s moment of insight was especially well-timed in that it corresponded with another interaction that he had with an OAC professor, one where he learned about the availability of graduate fellowships at the California Institute of Technology. The fellowship announcement bore the imprimatur of Caltech chemistry chief A.A. Noyes, among the country’s leading physical chemists and a mentor to several promising young scholars. It is no surprise then, that the flyer caught the eye of Pauling almost immediately and helped to steer him toward graduate studies in Pasadena.