Mary Jo Nye on Pauling’s Models

In the Fall of 2009 and Spring of 2010, Jane Nisselson, founder of the film and design studio Virtual Beauty, served at Resident Scholar in the OSU Libraries.  The intent of her visit was to gather research in support of a film which examines the importance of model building in science, both historically and in present day.

The above clip is a rough cut of part of this work.  It features OSU historian of science emeritus Dr. Mary Jo Nye discussing Pauling’s early breakthroughs in structural chemistry and the importance of model building within the chemical world.  “The iconography of chemistry,” she says, “is molecular representation.” All of the artifacts shown in the clip are held in the Ava Helen and Linus Pauling Papers.

Nisselson’s film, tentatively titled “Unseen Beauty: The Molecule Imagined,” is still in development. For more preliminary clips of this project, see the Virtual Beauty website and the Virtual Beauty channel on Vimeo.

An Electronegativity Breakthrough

Linus Pauling, ca. early 1930s.

Exciting news from the laboratories of Oregon State University: a group of researchers here have developed a method that simplifies the scientific understanding of electronegativity, a concept introduced and greatly advanced by Linus Pauling in the 1930s with his “electronegativity scale.”  We’ve written about Pauling’s electronegativity work before and since we’re in an interviewing mood lately, we thought we would catch up with the authors of this new breakthrough to find out what it’s all about.

Below are the fruits of our conversation with OSU’s Ram Ravichandran and Brian Pelatt, both doctoral candidates in electrical and computer engineering and co-authors of this important new paper, “Atomic Solid State Energy Scale.” (J. Am. Chem. Soc., 2011, 133 (42): 16852-16860.)

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Pauling Blog: In layman’s terms, what is electronegativity?

Brian Pelatt: Electronegativity as defined by Pauling is “the power of an atom in a molecule to attract electrons to itself.” When two atoms come together to form a molecule, the electronic charge will distribute itself so that one of them will be positively charged and the other negatively charged. Electronegativity is a way to explain that charge redistribution and quantify which atoms will more likely become negatively charged.

In the solid state energy framework, the electronegativity would be a measure of how good an atom is at taking, or “stealing”, electrons (negative charge) from other atoms. As an example, fluorine is the most electronegative element and has the largest solid state energy, so it will almost always take negative charge from another element when they bond.

Electronegativity calculations by Linus Pauling, ca. 1930s.

What was Linus Pauling’s contribution to our understanding of electronegativity?

Ram Ravichandran: The concept of electronegativity was first proposed by Pauling in 1932. In an effort to explain chemical bonding, Pauling looked at chemical bond energy data derived from thermochemical measurements. He noticed that the bond energy of dissimilar atoms was greater than the covalent bond of similar atoms. This difference, what he called the ionic character of the bond, gave rise to the electronegativity scale. However, his scale was arbitrary. He assigned a value of 4.0 to fluorine and then proceeded to use that value to calculate his electronegativity scale.

What has your group done now to further the concept of electronegativity?

Pelatt: Our group has furthered the concept of electronegativity by simplifying it in the Solid State Energy (SSE) framework. Now, instead of an element having an arbitrary number assigned to it as the electronegativity, it has a solid state energy that is based on experimental measurements and easy to understand. As can be seen from the electronegativity definition given earlier, electronegativity is a difficult concept both for students and teachers; this work simplifies the concept significantly. This approach can be used to simplify difficult chemical concepts such as the Hard-Soft Acid-Base Principle, chemical hardness, covalent vs. ionic bonding, and acidic/basic oxides.

The SSE approach is also a way to predict the band gap of a material, which is a fundamental property of materials and determines the behavior of the material, whether it behaves as a conductor, semiconductor, or insulator. Another advantage of the SSE approach is that it is based on solid state measurements, rather than gas phase measurements. Solid state electronics is the foundation of modern devices, having a scale that is tailored to the solid state is a huge advantage of the SSE concept.

How long have you been working on the problem?

Pelatt: We have been working on this for about a year, with most of the time collecting and interpreting data.

What methods did your group use to arrive at your conclusions?

Ravichandran: At the beginning, we did not anticipate working on electronegativity. We began looking at doping trends and energy band offsets of compounds to see if we could come up with an ability to predict properties of new materials.

A concept that connects chemistry and device engineering is the position of energy bands. In chemistry, these bands are known as HOMO (highest occupied molecular orbital) and LUMO (lowest unoccupied molecular orbital) which translate to the valence band and conduction band in device engineering. The separation between the bands is the energy band gap. The position of these bands is a measurable quantity known as ionization potential, IP (for HOMO, or valence band) and Electron Affinity, EA (for LUMO, or conduction band).

We started tabulating the values for EA and IP for about 69 compounds, and when plotted against the band gap, we noticed a curious trend. The values were centered around 4.5 eV, which corresponds to the standard hydrogen potential commonly used in electrochemistry and is also the bond strength of a hydrogen atom. This led us to conclude that hydrogen could serve as a universal energy reference position.

In chemistry, for a binary A-B compound, the HOMO is predominantly anion derived (mostly B), while the LUMO is cation derived (mostly A). Since the EA is related to the LUMO, by collecting and averaging the EA values for a cation such as aluminum, we were able to come up with a new “solid state energy” (SSE) value for aluminum. Similarly, by collecting and averaging IP values for nitride materials, we were able to calculate an SSE value for nitrogen.

When we then organized these SSE values for 40 elements in an increasing order along with the universal energy reference, common chemistry concepts such as oxide clarification, electronegativity, chemical hardness and ionicity became easier to interpret.

Where do you go from here?

Ravichandran and Pelatt: That is a very interesting question. Within our group, we are already utilizing this concept to be able to predict properties of new compounds. Dr. Michael Lerner and Dr. Richard Nafshun in the Chemistry department are both excited about incorporating this concept in graduate and undergraduate chemistry courses. We believe that the SSE concept has a wide reaching audience, from those trying to understand chemistry concepts to using this scale to investigate properties of new materials with applications in battery technology, water splitting and solar cells, amongst others.

Another exciting path is to see where the scientific community at large takes the SSE concept. The SSE values can be refined and this should spark a lot of conversation about measurement techniques. Doing so would provide a lot of subtle information about chemical bonding that is not readily apparent. The SSE scale helps to quickly estimate important properties of new materials, enabling rapid material development for applications. As a result, the most exciting aspect of SSE is how the community is going to use this scale in the search for materials with novel applications.

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For more on this breakthrough in our understanding of electronegativity, see this press release issued by OSU.

On Isosteric Isomers: An Important Early Paper

Ava Helen, Linus Jr. and Linus Pauling on a family hike, 1926.

In 1926 life was going well for twenty-five year old Linus Pauling – he had been married for a couple of years, had a healthy one year old son, and was quickly establishing himself as one of the top chemists in the United States. His primary research topic at the time was structural chemistry, and his hard work in the laboratory had already resulted in a good number of publications.

Nonetheless, Pauling had yet to publish a paper dedicated entirely to the subject that would soon become synonymous with his name: the chemical bond. This finally changed in 1926, when he and Sterling B. Hendricks  (who considered himself to be “Linus’ first student”) published a paper titled “The Prediction of the Relative Stabilities of Isosteric Isomeric Ions and Molecules.”

In this publication, Pauling and Hendricks calculated the potential energies of similar ions and similar molecules in order to predict the most favorable structure of a given ion or molecule. For example, the atoms in carbon dioxide (CO₂) can only be arranged in two configurations, OCO and COO. By comparing the potential energy values measured for each structure, it is possible to determine which is more likely to naturally form – the structure with the lower potential energy.

In their 1926 paper, Pauling and Hendricks first found it necessary to define the terms “isosteric” and “isomeric.” Isosteric refers to “molecules or ions that contain the same numbers of atomic nuclei and the same numbers of electrons, but differ in the positive charges on the nuclei.” Some examples of isosteric molecules are:

:N:::N: :C:::O: (:C:::N:)^- (:C:::C:)^-

Isomers, on the other hand, are molecules or ions that contain the same atoms in a different arrangement. In the carbon dioxide example given above, COO is one isomer and OCO is another. Other examples include the cyanate ion (NCO)^- and the fulminate ion (CNO)^- as well as the NON and NNO configurations of nitrous oxide (NO₂).

Although two molecules or ions that are isosteric aren’t always isomeric and vice versa, Pauling and Hendricks were only interested in substances that fulfill both conditions. Because isomeric and isosteric substances are so similar in many ways, Pauling and Hendricks argued that their differences in stability could be attributed almost exclusively to differences in potential energy. Using a rather complex equation that is explained in the paper, Pauling and Hendricks determined the most stable structures for a variety of isomeric and isoteric substances. Their results are displayed in the table reproduced below.

As the authors stressed in their paper, two points are particularly important to note: 1) the significance of relative values (as opposed to absolute values) of potential energy, and 2) that the configuration with the lowest potential energy is favored. Therefore, for carbon dioxide, the OCO configuration is favored over COO; for nitrous oxide, the NNO configuration is favored over NON, and so on. As it turned out, the two men’s results agreed very well with prior experimental and chemical evidence, with a few specific exceptions.

Although this paper is not generally counted among Pauling’s most important contributions, it does stand as an undoubtedly strong start to the chemical bond work that would win him the 1954 Nobel Prize in Chemistry. For more information on Pauling’s breakthroughs in structural chemistry, please visit the website Linus Pauling and the Nature of Chemical Bond: A Documentary History.

A Theory of the Structure of Ice

Linus Pauling, December 1935.

By 1935 – 75 years ago this year – Linus Pauling was in the thick of his scientific career. He had been at Caltech for over a decade, and had already conducted an impressive amount of important research on crystal structures and the chemical bond. During this year, Pauling also postulated a theory of the structure and entropy of ice, one of his lesser-known, yet still very significant, contributions.

Pauling’s work on this subject was encapsulated in a paper titled “The Structure and Entropy of Ice and of Other Crystals with Some Randomness of Atomic Arrangement,” which appeared in the Journal of the American Chemical Society in December 1935.

In order to better understand Pauling’s theory, it is necessary to discuss what was already known about water and ice at this time. First, as stated in Pauling’s paper, “it has been generally recognized since the discovery of the hydrogen bond that the unusual properties of water and ice (high melting and boiling points, low density, association, etc.) owe their existence to hydrogen bonds between water molecules.”

Second, the precise arrangement of the oxygen atoms – but not of the hydrogen atoms – in ice crystals was already known. Each single oxygen atom is tetrahedrally surrounded by four other oxygen atoms.  It was also generally understood that there is only one hydrogen atom located between each oxygen atom.

From this base of knowledge Pauling pondered “whether a given hydrogen atom is midway between the two oxygen atoms it connects or closer to one than to the other.”  In evaluating the experimental evidence, Pauling decided in his paper that the hydrogen atom is, in fact, closer to one oxygen that to the other.

At this point, Pauling made four assumptions, all of which he supported later in the paper. These assumptions are as follows:

1. In ice each oxygen atom has attached to it two hydrogen atoms affixed at distances of about 0.95 Å.  These atoms form a water molecule, with the H-O-H angle measured at about 105º, as in the gas molecule.
2. Each water molecule is oriented such that its two hydrogen atoms are directed approximately toward two of the four oxygen atoms which surround it tetrahedrally, forming hydrogen bonds.
3. The orientations of adjacent water molecules are arranged such that only one hydrogen atom lies approximately along each oxygen-oxygen axis.
4. Under ordinary conditions the interaction of non-adjacent molecules would not appreciably stabilize any one of the many configurations satisfying the preceding conditions with reference to the others.

Adhering to these assumptions, Pauling theorized that the water molecules in ice crystals can orient themselves in a number of different ways, and that the crystal can likewise change from one orientation to another, provided that it adheres to the four assumptions.

Pauling used his theory to calculate the number of possible configurations available to the crystal (as allowed by his theory), and in turn used this number to calculate the residual entropy of ice. He found this value to be in very good agreement with the experimentally observed entropy value.

Pauling felt good about his data, concluding that “the observed entropy of ice at low temperatures provides strong support for a particular structure of ice, and thus gives an answer to the question which has been extensively discussed during the past few years.”

Historians of science now concur with Pauling’s assessment. Writing on the matter in his Pauling biography Force of Nature, Thomas Hager notes

[I]n 1935, Pauling had a flash of insight that led to paper on ‘orientational disorder’ – a theory concerning water molecules that explained the residual entropy of ice at absolute zero. It was purely theoretical work, harkening back to his days with [Richard C.] Tolman. Thirty years later, when sophisticated computers were finally able to run the numbers thoroughly, Pauling’s theory was proven right. Now called ‘proton disorder,’ the idea became, as one student of the field say, ‘the most important American contribution to the modern crystallography of water.’

Julia Bursten, Resident Scholar

Peter Freeman, Julia Bursten and Judith Freeman.

Julia Bursten is the fifth individual to complete research as a Resident Scholar in the Oregon State University Libraries Special Collections.  Bursten is a doctoral candidate in the History and Philosophy of Science at the University of Pittsburgh.

[Update: transcribed video of Bursten's Resident Scholar presentation is now available.]

Bursten came to Corvallis to study a specific aspect of Linus Pauling’s valence theory of chemical bonds.  In particular, she is interested in the development of Pauling’s ideas on the bent equivalent double bond, or “banana bond” as it is sometimes called.

From his unique vantage point as a structural chemist immersed in contemporary work on both molecular architecture and quantum mechanical behavior, Pauling was well-positioned to make groundbreaking contributions to the scientific understanding of how atoms interact to form molecules.  In due course, he proposed that atoms assumed the form of tetrahedra, and that chemical bonds – including the double bond, a shorter and stronger bond that incorporates four bonding electrons rather than the usual two  – could be represented in a similar fashion, as tetrahedra that are, roughly speaking, pressed together.

Pauling wrote extensively in support of this theory, though focused much of his attention on single bonds.  Indeed, aside from a brief mention in a 1931 paper, Pauling’s only other early reference to the bent equivalent bond was a short passage in the first (1939) edition of The Nature of the Chemical Bond, in which Pauling again reiterated his position in support of the merged tetrahedra.

As with most of Pauling’s structural chemistry work, this picture of the double bond was generally accepted for several years.  Gradually though, Pauling’s valence approach began to come under attack by a group of scientists supporting the molecular orbital model of chemical bonds.  While confusing to most non-scientists, the differences between the valence bond theory and the molecular orbital approach are ably described as follows by Pauling biographer Thomas Hager.

In Pauling’s approach, derived from the electron-interchange idea of Heitler and London, molecules were aggregates of individual atoms, each linked to its neighbors by bonds formed by electrons localized between two nuclei.  The number of bonds equaled the element’s valence, or bonding capacity…. In theory, the total quantum-mechanical state of a molecule could be calculated by adding together the wave functions that were involved in each bond, with appropriate adjustments for the effect of each bond upon its neighbors….

[The] molecular orbital theory [is] an approach predicated on a belief that molecules were not what valence bond advocates thought they were.  Molecules to [molecular orbital proponents] were not aggregates of distinct atoms connected by distinct bonds but things unto themselves, with their own odd behavior explicable only in molecular terms….  [The theory posits] that molecules could be more profitably viewed as if their binding electrons were somewhat delocalized and spread across the surface.

According to Bursten, the molecular orbital supporters suggested that both their approach and Pauling’s valency approach yielded the same results in the explanations that they gave for molecular behavior.  This noted, the mathematics underlying the molecular orbital techniques were much simpler to apply and, as a result, an improvement on Pauling’s work.

Notes prepared by Linus Pauling for his 1958 defense of the valence bond theory. Note in particular his statement: "Bent Bonds Are Better!"

Pauling did not respond to these developments until 1958, when he issued a series of three speeches in support of the valency approach and, specifically, his model of the double bond.  These presentations were followed by a detailed technical rebuttal of the molecular orbital school in Pauling’s third (1960) edition of The Nature of the Chemical Bond.  As Bursten points out, Pauling had not published this defense in any formal channels before including it in his 1960 text – an approach viewed by many as highly unorthodox.

As it turned out, Pauling’s writings in 1960 marked both his last major defense of the bent equivalent double bond as well as the beginning of the end for Pauling’s valence theory.  Over time the molecular orbital approach gained the favor of the scientific community.  Indeed, Bursten’s research indicates that Pauling was denied several later grant requests for work on theoretical structural chemistry, precisely because he sought to conduct further research grounded in his valence bond model.  As with his and Robert Core’s triple-helix structure for DNA, Pauling’s valence bond approach was swept aside by the research of others.

The OSU Libraries Special Collections Resident Scholar Program is generously supported by the Peter and Judith Freeman Fund. Past recipients have included Dr. Burtron Davis of the University of Kentucky’s Center for Applied Energy Research, Toshihiro Higuchi of Georgetown University, Dr. Mina Carson, professor of history at Oregon State University and Jane Nisselson, a documentary filmmaker based in New York City.

Resonance in Benzene and Beyond

Introductory sentence and diagram from Pauling and Wheland's paper on resonance in benzene and naphthalene, June 1933.

[Part 2 of 2]

“Suppose that we ask: is it necessary that a molecule such as CO have a definite valence-bond structure? The answer, which is part of the new idea, is no; instead the CO molecule may have (and does have) a structure which is neither C=O or C≡O, but is somewhere between them, or which rather has some aspects of both. It is customary now to speak of the molecules as resonating between these two structures.“

-Linus Pauling, 1936.

As with the structure of carbon monoxide, the principle of resonance also explains what was once a chemical enigma – the true molecular structure of benzene.

Benzene, a double-bond conjugated six member hydrocarbon ring, can be represented by two structures that are equivalent in energy.

A simple model representing oscillation between the two primary structures is, however, insufficient as it does not explain one of the principle chemical properties of the molecule – its inability to saturate.

The application of the theory of resonance permitted the determination of a more accurate model. In the resonance model, the molecular configuration of benzene is such that all possible structures (including ones not shown above) contribute to the true structure – a combination of all structures at once, with each carbon-carbon bond energetically equivalent.  As Pauling wrote in his 1946 Encyclopedia Britannica entry on resonance

It is sometimes found…that a choice cannot be made between two or more structures which are about equally stable, and of which no one accounts in a completely satisfactory way for the properties of the substance. The concept of quantum-mechanical resonance has provided the solution to this problem: namely, the actual normal state of such a molecule does not correspond to any one of the alternative reasonable structures, but rather to a combination of them, their individual contributions being determined by their nature and stability….Just as it is customary to speak of the electrons in an atom in its normal state as moving around the nucleus in roughly the way described by the old quantum theory…so is it customary, and for some purposes useful, to speak of the resonance of a molecule in its normal state between two or more structures.

Based on this theory, the benzene molecule is now often depicted as a hexagon with a circle in the middle representing the true resonating nature of the molecule.

This structure has since been verified by multiple experimental techniques such as electron diffraction, x-ray diffraction, and molecular spectroscopy.

Having recognized resonance as an important missing link in understanding molecular bond structure, Pauling applied his theory to a large collection of empirical results.  (For one, he identified electronic resonance as the principle that allows for the formation of four equivalent bonds to be formed by carbon.)  His analysis consistently explained gaps in the classical models of bond theory and aligned with the quantitative data available.   Pauling’s theory of resonance has since contributed fundamentally to the scientific understanding of molecular shape and stability, and has permitted insight into the true nature of the chemical bond.

Developing the Theory of Resonance

Linus Pauling, 1930.

[Part 1 of 2]

“I think my work on the chemical bond probably has been most important in changing the activities of chemists all over the world – changing their ways of thinking and affecting the progress of the science.”

Linus Pauling, 1977.

In early 1932, Linus Pauling spent several months visiting the University of California, Berkeley and the Massachusetts Institute of Technology to present two different lecture series on the theory of resonance and its implications for molecular structure and function. This topic, a product of Pauling’s adventures in Europe as a Guggenheim fellow, would profoundly impact the ways in which twentieth-century chemists ultimately understood the chemical bond and predicted molecular structures.

Throughout his long career, Pauling sought to improve his understanding of molecular structure in order to better predict chemical function. As Pauling saw it, molecular structure dictates function and should thus be considered accordingly.  As he wrote in 1946

…I am confident that, as our knowledge of the structure not only of simple molecules but also of proteins and other complex constituents of organisms increases, we shall in time achieve an insight into physiological phenomena which will serve as an effective guide in biological and medical research, and will contribute to the solution of such great practical problems as those presented by cancer and cardiovascular disease.

Indeed, much of Pauling’s work sought to develop the tools necessary to enable chemists to bridge the gap between structure and function.  Pauling spoke somewhat literally of this quest in 1936, in a speech where he compared the tools of the chemist to those of an architect.

The structural chemist of the past and present has been an architect working with materials of whose nature he is largely ignorant – an architect who does not know what an I beam is, but only that it can be used in his construction, and who must proceed to design structure after structure, to find ultimately that certain designs lead to satisfactory results – to a building with rooms adapted to the use of certain visitors, to bridges strong enough to hold their load, and so on. The structural chemist of the future will be able to plan his structures and forecast their properties in the same definite way that the architect and engineer plan the macroscopic, even gargantuan, structures of modern civilization.

In the years just prior to, and at the start of, Pauling’s career, great strides had been made in  molecular structure determinations.

G.N. Lewis, ca. 1930s.

In 1916, for example, Gilbert Newton Lewis developed a theory of molecular diagrams based on valence electrons, now referred to as Lewis-dot structures. The subsequent application of spectroscopic methods to molecular chemistry allowed for more direct quantitative studies of atomic and molecular structure. Later, advancements in quantum mechanics increased chemists’ and physicists’ understanding of the detailed interactions that occur between nuclei and electrons that ultimately determine atomic and molecular structure.  Meanwhile, various valence bond theories had been developed but were not applicable to all matter uniformly.

While these and other contributions were significant, many questions still remained as not all quantitative data aligned with current theories. To provide an explanation for the many apparent holes in understanding, Pauling developed his theory of resonance – an idea which became the central concept of Pauling’s valence theory.

Resonance, or electron exchange, is a property integral to the formation and maintenance of chemical bonds, as it accounts for the formation of hybrid structures that cannot be explained by the classical models of molecular structure alone.

Pauling used his theory of resonance to explain why many molecules can be drawn in various forms according to Lewis’s scheme even though no single structure could be differentiated as the “correct” structure based on energy theories and quantum mechanics.

According to Pauling’s theory, these structures could not be differentiated quantitatively because the electrons exchanged between atoms caused the molecule to resonate between multiple structures. Thus the structure of a molecule is not made up of one single structure, but in some cases, such as in carbon monoxide (CO), the true nature of the molecule resonates between multiple structures. Pauling therein predicted that the CO molecule fluctuates rapidly between multiple conformations thus creating a more stable structure, known as a “resonance hybrid.”

By Pauling’s way of thinking, the theory of resonance explained many of the obvious inconsistencies in the understanding of specific molecules at that time, and he argued that the theory should be applied when predicting new molecular structures and functions.  In our next post, we’ll talk more about the impact of the theory of resonance by examining its application to the study of the enigmatic structure of benzene.

Project Adrift: The Second Edition Fizzles Out

[Part 4 of 4]

The cordial disagreements over the shape of the second edition of Introduction to Quantum Mechanics began in August 1955 when Martin Karplus sent to Linus Pauling his first revision of the book. Many of the revisions that Karplus was making did not fall in line with those that Pauling and E. Bright Wilson, Jr. had in mind. Pauling and Wilson expressed their concerns to Karplus, Pauling writing that “I think that the revision that you propose is more extensive than we want” and Wilson suggesting that “I think in general I stand part way between the two of you with respect to the extent of revision necessary.”

Clearly there existed significant gaps in agreement; gaps which steadily grew into crevasses.

In October 1956, Pauling wrote to McGraw-Hill requesting that the previous agreement concerning royalties for the second edition be revised to that of an equal division among the three co-authors, as “circumstances have changed since 1935.” This was an early sign of what would soon become an apparent lack of investment in the second edition on Pauling’s part.

In November Karplus met with McGraw-Hill and agreed on a new deadline of Fall 1957 for publication in 1958. Nearly a year later, in October 1957, the deadline was extended again to the summer of 1958. In a September 1957 letter, Karplus gave some insight into the reasons for the delays:

I am very sorry that I have not been able to accomplish more on the revision. Other obligations as well as some personal problems, have prevented me from devoting as much time as I should have liked to give to the work.

As time passed, the disagreements between the co-authors and the lack of organization were becoming more and more apparent. In a letter to Karplus, Pauling’s frustration with the situation was evident

I wish that you would send me a copy of some of your material on the revision of Introduction to Quantum Mechanics…I trust that you are not changing the book completely.

In early March 1959, Karplus wrote to Pauling and offered a glimmer of hope that the end of the revisions was near, suggesting that “in terms of the present rate of progress, it is perhaps not completely unrealistic to hope that the new edition will appear in the spring of 1960.”

Yet the publication delay continued.  When Pauling was asked to compose a recommendation for Karplus for a position opening at Tufts University in October of 1959, Pauling shed further insight into the issues plaguing the revision.

The work that Karplus has been doing in revising the book seems to me to be of the highest grade. My suggestions in the main have dealt with a simplification of what he has written. I have not found any errors, not even in judgment, except that I am afraid that he tends to be interested in the more complex aspects of the subject….I may point out that he was, I think, somewhat disturbed in his research and his revision of Introduction to Quantum Mechanics by personal difficulties.

Clearly reasons other than lack of cooperation between co-authors were causing the delays.

Two years later, in 1961, Pauling wrote to a Dr. T. Katsurai in Tokyo, who had inquired about the publication date of the new edition. In this reply, Pauling confessed of his own contributions to the slowness of the project.

Professor Wilson and I have got to the stage in life when we have many duties and less energy than formerly. Professor Karplus is hard at work, but the job is a big one, and I surmise that it will still be a year or more before it is finished.

Later that year, Pauling received a letter from a McGraw-Hill editor informing him that “Professor Karplus assured me three weeks ago that he and Professor Wilson are actively revising, and that they anticipate completion in the summer of 1962.” And yet, 1962 passed without publication.

In 1963 Pauling cut himself off from the project, in large part because of the growing press of work, especially peace work, that now defined his every waking moment.  In a letter to Karplus informing him of his decision, Pauling wrote

I feel that I should not be a co-author with you and Bright. I have decided that my many activities, combined I think, with some lack of interest in the details of modern quantum mechanics, will prevent me from making any contribution to the new book.

Wilson responded curtly to Pauling’s withdrawal, insisting that it would reflect poorly on Karplus as well as the second edition.

I was very disturbed to receive the copy of your letter of May 22 to Martin Karplus. I don’t really see how you can do this to him. After all, it was you who chose him and persuaded him to undertake the job of revising our book, a job which has turned out to be enormously more burdensome and worrying than he could possibly have realized, in good part because of his own success in research. I fear that he (and the public) will infer that you have no confidence in him and prefer the old edition even though you haven’t examined the new.

Pauling responded to his colleagues apologetically, explaining that his request to withdraw was not based on a lack of confidence in the project, but on his own inability to participate.

I[f] you and Bright are willing that I be a co-author, I would be happy to be. My conscience has been bothering me because I have not contributed anything significant to the revision. I do not like to accept anything that I do not deserve, and I have felt doubtful that I deserve to be a co-author of the new edition.

Nine years later, in 1972, Karplus sent Pauling a letter informing him that the book was near completion. But by 1974, nearly twenty years after Pauling and Wilson made their decision to publish a second edition, McGraw-Hill still had not received a manuscript. At this time, Pauling sent a final letter to Wilson and Karplus officially withdrawing as co-author.  The second edition never made it to print.

Introduction to Quantum Mechanics: A Second Edition?

Martin Karplus and Linus Pauling, 1960s.

[Part 3 of 4]

During its first eighteen years in print, Linus Pauling and E. Bright Wilson, Jr.’s Introduction to Quantum Mechanics sold over 17,000 copies.  Heartened by the success of the first edition, Pauling wrote to his co-author in November 1953,

It seems to me that the book has been successful enough to justify a second edition, and I do not think that any of the newer books takes its place.

In December, Wilson replied in agreement but with slight trepidation.

I should be quite willing to have a second edition of Introduction to Quantum Mechanics prepared if it didn’t involve too much work for me and if we could agree on the general principles which we were going to use in carrying out the revision.

Wilson’s reply in hand, Pauling wrote to their editor at McGraw-Hill, Hugh W. Handsfield, and informed him of their decision to revise the 1935 text. Handsfield replied with a deadline of winter 1956 in preparation for a publication date of January 1957.

To alleviate some of the work load, Pauling suggested that the two authors collaborate with a young Ph. D. named Martin Karplus, a former student of Pauling’s and a recent graduate of Caltech. Wilson agreed, Pauling extended the offer and Karplus accepted, if warily, noting that “though I am not certain that I am qualified for the task, I should like to attempt it.”

Pauling, Wilson, and Karplus agreed to divide the royalties according to estimated contribution levels, allotting a quarter of both Pauling’s and Wilson’s royalties to Karplus. Having delegated a significant portion of the revision work to Karplus, Pauling did not foresee much difficulty in the development of a second edition.

Unfortunately, there emerged significant flaws in their three-way collaboration, especially that of geography. When the authors began making preparations for the revision, Karplus was in England at Oxford completing a National Science Foundation Fellowship, Pauling was in Pasadena at Caltech, and Wilson was in Massachusetts at Harvard. This distance presented obvious complications in revision efficiency, appreciably slowing down the process.

The problems, however, did not end with logistics, but extended to revision philosophy. When Pauling and Wilson asked Karplus to participate, they both made it very clear that they wished to maintain the integrity of the first edition by keeping the focus of the text on applications to chemistry for students that were less mathematically inclined.  As Bright Wilson wrote in a letter to Karplus

I think Professor Pauling agrees with me that we are very anxious to keep the book at a level which can be understood by first-year graduate students in chemistry. It has always been the great feature, in my opinion, which has made the book so successful in its first edition. My main contribution to it was to bring an adequate degree of ignorance into the authorship, and I therefore claim a lot of credit for the success of the book because I was not able to understand anything highbrow and therefore there was not very much highbrow put into it.

Their reasons for keeping the book at such a level extended beyond academic intentions. Both original authors recognized the market potential for a text on quantum mechanics applied to chemistry as there existed many quantum mechanics books available for physicists and none, other than theirs, for chemists.  Wilson summed it up succinctly

I don’t think we can hope to compete with the books designed specifically for physicists and that we should try very hard not to increase the level of difficulty because otherwise we will lose our principal attractive feature.

Karplus found such revisions difficult as he struggled to incorporate what he believed to be important while maintaining the original format developed by Pauling and Wilson. These disagreements in revision philosophy ultimately amounted to yet another considerable hurdle in the revision process.  Time passed quickly over the next year and it soon became apparent that there was no conceivable way that the co-authors were going to meet their first deadline.

Pauling and Wilson

[Part 2 of 4]

In 1926, while still in Europe completing his Guggenheim fellowship, Pauling attended history’s first full-term lecture on the new concept of wave mechanics as applied to quantum theory. This course, taught by Arnold Johannes Willhelm Sommerfeld, a renowned German theoretical physicist and a pioneer of quantum mechanics, was historically significant as the first of its kind.  Sommerfeld would later write of the classes,  “My first lectures on this theory were heard by Linus Pauling, who learned as much from them as I did myself.”

Upon returning from Europe to Caltech, Pauling used the knowledge gleaned from his Guggenheim experience to develop his own lecture series on quantum mechanics. Among those who attended these was none other than Albert Einstein who sat in on one of Pauling’s talks in 1930.

The content of this course became the foundation for Pauling’s first textbook, Introduction to Quantum Mechanics with Applications to Chemistry, which he developed with a former Ph. D. student named Edgar Bright Wilson, Jr.

E. B. Wilson, Jr., known to many as Bright, was born in Gallatin, Tennessee in 1908. After graduating from Princeton in 1930, Wilson attended Caltech and, under Pauling’s direction, received his doctorate in 1933. Wilson then became a fellow at Caltech until accepting a position at Harvard in 1934.

E. Bright Wilson, 1970.

In 1935 Wilson and Pauling published their co-authored text, which took the duo over two years to transform from Pauling’s original lecture notes.  The primary goal in writing the volume was to “produce a textbook of practical quantum mechanics for the chemist, the experimental physicist, and the beginning student of theoretical physics,” for the authors firmly believed that quantum mechanics had applications to nearly all scientific disciplines.

Cognizant of the need to guide the less mathematically adept reader “through the usually straightforward but sometimes rather complicated derivations of quantum mechanics,” Pauling and Wilson formatted their content such that it could be understood by those with mathematics training up through calculus, with some limited additional background on complex numbers, differential equations, and partial differentiation.  Pauling and Wilson wrote that

The book is particularly designed for study by men without extensive previous experience with advanced mathematics, such as chemists interested in the subject because of its chemical applications.

In completing the text, the authors acknowledged a number of mentors and colleagues – many of them Caltech contemporaries – for their contributions to both the authors’ own personal knowledge and to the field of quantum mechanics: Arnold Sommerfeld, Edward U. Condon, Howard Percy Robertson, Richard C. Tolman, Philip M. Morse, Leslie E. Sutton, George W. Wheland, Lawrence O. Brockway, Jack Sherman and Sidney Weinbaum. And last, but certainly not least, the authors acknowledged their wives, Emily Buckingham Wilson and Ava Helen Pauling.

In the years following publication, Wilson built a career as a highly successful chemist and an esteemed member of the scientific community. In 1949 Wilson too received a Guggenheim Fellowship, with another to follow in 1970. And in 1975 Wilson was awarded the prestigious National Medal of Science for physical sciences, just one year after Pauling.