Chemistry's Decision Point - ACS Publications - American Chemical

Chapter 7

Chemistry’s Decision Point: Isotopes

Downloaded by UNIV OF FLORIDA on November 26, 2017 | http://pubs.acs.org Publication Date (Web): October 30, 2017 | doi: 10.1021/bk-2017-1263.ch007

Brett F. Thornton1 and Shawn C. Burdette*,2 1Department

of Geological Sciences and Bolin Centre for Climate Research, Stockholm University, 106 91 Stockholm, Sweden 2Department of Chemistry and Biochemistry, Worcester Polytechnic Institute, Worcester, Massachusetts 01609-2280, United States *E-mail: [email protected].

Although the modern periodic table barely resembles the one constructed by Dmitri Mendeleev, every chemistry student learns that the placement of missing elements in the open slots of Mendeleev’s table was a scientific triumph. The discovery of isotopes in the early 1900s was an inflection point in periodicity, and chemistry as a discipline. Chemists once characterized each new isotope as a unique element—but as isotopes proliferated, fitting them into the existing periodic table became impossible. Several decades passed before the concept of isotopy fully developed. At that point, scientists seemingly concluded that chemistry occurred at the atomic level, and isotopic differences were the purview of physics. Had a different understanding of isotopy prevailed, the direction of chemistry could have changed dramatically. While the trajectory of synthetic chemistry might have remained constant, ‘chemists’ may have dominated the discovery of new superheavy elements by appropriating the modern conventional definition of ‘nuclear physicist’.

© 2017 American Chemical Society Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.

Isotopes and the Edge of Chemistry


In 1959, a rather innocuously named annotated bibliography of the previous half-century of efforts to separate isotopes appeared as a government report issued by Oak Ridge National Laboratory (1). Isotope Separation and Isotope Exchange, by G. M. Begun, lists nearly 2500 papers published between 1907 and 1957 on the problems and solutions of separating isotopes. As a bibliography, it is not a narrative work, but a short foreword begins: The simple building blocks of the nineteenth century scientist, the ninetytwo chemical elements have been multiplied today to well over thirteen hundred different nuclides of which some three hundred and thirty occur in nature while the remaining number are artificially radioactive. We know now that isotopes are not identical, even chemically, and that their nuclear properties may be vastly different. In those four sentences, one finds the quintessential struggle to define the limits of chemistry, or to ask the question very bluntly: what constitutes chemistry? The question has proven decidedly difficult to answer, especially where physics and chemistry collide. Although some have posed the question as “is chemistry a branch of physics?” (2), we will take as axiomatic that chemistry is not a branch of physics. Instead, we propose that the natural breakpoint between chemistry and physics is best seen, described, and understood through the discovery and elucidation of isotopy in the early 20th century. The “simple building blocks of the nineteenth century scientist”, the chemical elements, today have expanded to include the approximately 3000 known nuclides (3). Each of these nuclides is a unique building block of matter, but to most chemists, all the isotopes of an element are squeezed into a single spot on the periodic table. The word isotope even derives quite elegantly from the Greek words for ‘same place’. It is almost a scientific cliché to mock the supposed conventional wisdom at the end of the 19th century, when many confidently believed that all the major discoveries in science had been revealed. Physics had a working model of the world in Newtonian mechanics, and all that remained was to fill in the minutia. With the state of knowledge at that time, scientists sought no, or very few, new physical laws to model the natural world. Failings started appearing in the late 19th century. The Michelson-Morley experiment famously showed there was no luminiferous aether (4). Any remaining scientific hubris was expunged by the end of 19th and beginning of the 20th century with discoveries including relativity, radioactivity, subatomic particles, and isotopy. Suddenly, science was not finished after all. Often, the evolution of modern 20th century science is told from the perspective of physics (5); however, chemistry and physics were intimately intertwined during the first few decades of the 20th century. Many would-be physicists became chemists, and vice versa, amalgamating the fields with neither side immediately taking full ownership of isotopes nor other discoveries. In modern times, such work undoubtedly would have been characterized as 120 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


interdisciplinary, and pursued as a joint venture. In these early days, however, research on concepts such as isotopes almost inevitably was partitioned between chemists and physicists. An effect of these decades of confusion about the course and content of chemistry was linguistic flux. Even into the 1920s and 1930s, the scientific language surrounding chemical elements was not fully resolved. A 1920 paper in Science announced the successful separation of chlorine into chlorine (35Cl in modern notation) and meta-chlorine (37Cl) (6). At the time, one might have concluded that 35Cl was an actual chlorine atom, and 37Cl was just 35Cl with two strongly bonded hydrogen atoms since the neutron was not discovered until 1932 (7). Since the masses of 35Cl and 37Cl differed, how could they have the same chemical properties? Once the separation of isotopes of various elements was demonstrated, they seemed even more like fundamental bits of matter. 35Cl was plainly not the same thing as 37Cl, but frustratingly, their properties other than mass seemed identical. The linguistic question lingered far into the 1920s—should meta-chlorine be characterized as an element? Isotopy emerged shortly after a timespan when 21 elements were discovered between the publication of the periodic table in 1869 and the start of the 20th century (8). Some of these new elements fit nicely into the vacant slots left by Mendeleev, but there were a few non-trivial puzzling exceptions for 19th century chemists. The proliferation of rare earth elements was problematic to the fledging system of element organization. Cerium, lanthanum, erbium and holmium, had been identified by 1869, and the seven similar elements that were discovered by the turn of the century needed positions that were not obvious in Mendeleev’s system. Of course, this family eventually would be identified as the lanthanides. Of the modern actinide family, only thorium and uranium were known when Mendeleev drew his periodic table. The next actinide, actinium, would not be glimpsed until 1899 (9), but not fully characterized until 1902 (10). The placement of elements in group 3 of the periodic table—scandium and yttrium, plus lanthanum and actinium or lutetium and lawrencium—remains unresolved (11). At first, there was not even a column in the periodic table that might host the noble gases, but these elements eventually would be placed in logical, systematic locations (12).

Radioactivity: Another Challenge to the Periodic Table Like the challenges of placing the lanthanides, actinides, and noble gases on the periodic table, radioactivity created a similar conundrum for the expanding system. The report of polonium and radium in 1898 was the beginning of a 20th century preoccupied with radioactivity (13, 14). Within a few short years, a whole menagerie of new radioelements was discovered to be emanating from, or forming in, heavy elements like uranium. Unlike almost all the previously known elements, these radioelements had finite lifetimes—sometimes measured in seconds. Scientists were confronted with unexpected questions. Where did these radioelements come from, and where did they go when they vanished from existence? Other than radioactivity, the radioelements seemingly exhibited identical chemical reactivity, like previously known elements. Plus, they could 121 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


not be broken down into more elementary components by chemical processes. Being indecomposable, radioelements were clearly elementary substances, but how did they fit into the periodic table? There were clearly too many to fit into the few remaining gaps on the periodic table? These early radioactivity discoveries again illustrate both the convergence and divergence of chemistry and physics in the history of the periodic table. Henri Becquerel along with Pierre and Marie Curie were collectively awarded the 1903 Nobel Prize in Physics for the discovery of spontaneous radiation emanating from uranium ores (15, 16). Eight years later, Marie Curie would receive the Nobel Prize in Chemistry for the isolation of polonium and radium from those same minerals (17, 18). From this one might conclude that chemistry prioritized the discovery of new substances (radioelements), and physics focused on the general phenomenon (radioactivity), even though the two were inexorably intertwined. By 1905, some chemists, including Kazimierz Fajans, began to hypothesize that the old periodic table had outlived its usefulness for organizing elements. A new arrangement of the table, which assumed the recently discovered α-particle was the basic building block of atoms (19), had been devised by Antonius van den Broek with spaces to accommodate the proliferating number of proposed radioelements (20). Fajans was a remarkably versatile chemist whose career spanned most of the 20th century. He was a physical organic chemist by training, having obtained his PhD under the direction of Georg Bredig at the University of Heidelberg studying chiral resolutions. Fajans became interested in binding forces in carbon during this time, but concluded these topics would be better addressed by physics than chemistry. Thus, a chemist chose to change fields, become a physicist, and investigate the new radioelements from that viewpoint. After receiving his PhD in 1909, Fajans moved to Manchester, England to work with New Zealander Ernest Rutherford during this time of rapid and exciting advancements in atomic theory. Upon returning to Germany, he discovered element 91 in 1913 with Oswald Göhring (21). They had discovered a short-lived isotope of element 91, which they named brevium. A much longer-lived isotope was discovered years later, and through a ‘linguistic loophole’ rooted in conflating the concepts of isotopes and elements, element 91 was renamed for this true ‘parent of actinium’, protactinium (22). Along with radon (23), protactinium is one of only two times when an isotope name displaced the discoverers’ preferred element name. Fajans had a busy 1913 in addition to discovering element 91. He published the radioactive displacement laws that explained the effects of α- and β-decay in the radioactive elements (24). Fajans’ explanations were based on electrochemical results showing radioelement similarities. A few weeks later, similar radioactive displacement theories from Englishman Frederick Soddy appeared (25). Soddy’s conclusions were based on chemical, not electrochemical, similarities of radioelements. Soddy reportedly had read Fajans’ paper on the displacement laws before publishing his study, and became more associated with the laws in the English-speaking world. Both Fajans and Soddy noted that radioelements that fell onto the same spot on the periodic table would be chemically inseparable. Soddy’s presentation of his work on the displacement laws as distinct from 122

Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


Fajans’ work has been criticized (26), but without a resolution to the questioned possibility of intellectual infringement. Frederick Soddy had a remarkably similar career path to Fajans, albeit he was Fajans’ senior by a decade. Soddy was trained at Merton College at the University of Oxford. Following his studies, he moved to McGill University in Canada where he met Rutherford (27), who was making the earliest observations of radioactive emanations of thorium (28, 29). In 1900, Rutherford and Soddy began to work on various aspects of radioactivity and radioelements. Returning to England, Soddy worked with William Ramsay on production of helium by radioelements (30). Soddy then became a lecturer at the University of Glasgow in 1904. In 1911, Soddy reported that mesothorium (228Ra) and radium (226Ra) had different radioactivities, but the same chemical properties (31). In addition to the controversy with Fajans about recognition for establishing the radioactive displacement laws, Soddy, along with John Cranston, was also the aforementioned discoverer of protactinium—five years after Fajans reported on brevium (22). Otto Hahn and Lise Meitner also independently identified the same isotope of element 91 (protactinium) nearly simultaneously (32). At the time of Fajans’ and Göhring’s discovery of brevium, the proposal on the nature of isotopes was more than six months in the future. In 1919, to the happy surprise of Meitner and others, Fajans agreed that the longest-lived isotope should be used to name the element (33). Since the definitions of both element and isotope were still evolving in 1919, his stance was not surprising. Fajans originally had held that isotopes were different elements, since an elementary substance by definition could not come in different forms; however, by the 1920s he had adopted the current definition (34). In retrospect, however, the choice of naming element 91 protactinium vastly diminished the recognition of Fajans’ original discovery. The element name change was rationalized by claiming that discovering the parent of actinium was the major finding at the time. Although actinium was known to be associated with uranium, the decay series by which actinium formed was unknown. Once the isotope corresponding to the true mother of actinium was found (22, 35), brevium, which had no such cachet, was discarded as a name for the element (36). Otto Hahn was more sanguine, later stating clearly that protactinium was an isotope of brevium. Fajans would later regret this course of events, and his apparent loss of discovery priority for element 91 (37). At the time, and even now, Soddy’s protactinium isotope often is described erroneously as the discovery of element 91. Soddy published a letter to Nature in December 1913 (38), which took several contemporary developments to the logical conclusion, and coined the term isotope. In his 1922 Nobel Prize acceptance speech, Soddy described isotopes as atoms that "have identical outsides but different insides" (39). As every chemist knows, the ‘outsides’—the electrons—impart the same chemistry to all the isotopes of a given element; the ‘insides’—the protons and neutrons—don’t change the element’s chemistry, right? The importance of discovering isotopy to understanding atomic structure cannot be overstated. If Soddy had not expressed this first, others would certainly have formulated the same conclusion eventually, as several of his contemporaries were tantalizingly close. If the rapid advances in understanding of atomic structure in 1913 seem impressive today, they were even 123


more so then. Fifty years later (40), Neils Bohr commented on the events of the year: “1913 was a very curious time.”


Soddy, Aston, and the Mass Spectrometer In the late 19th century, electrical discharge vacuum tubes, primitive cathode ray tubes (also known as Crookes tubes), were invented, which allowed the production of positive rays. These rays were the heavy positively charged ions produced by stripping one or more electrons from an atom. Since the mass of the extracted electron is always an infinitesimal portion of the atom’s weight, the electron is ejected at high speed, while the much heavier positive ion moves slowly. By introducing electric fields, the positive ions can be accelerated and separated based on their mass/charge ratio (41). The spatially separated positive ions then strike a photographic plate, fluorescent screen, or other detector, and thus a crude precursor to a mass spectrometer could be fabricated. Using this method, scientists finally showed that atoms of the same element had the same atomic mass. This had been an unchallenged, but unproven assumption of atomic theory since Dalton in the early 1800s. In 1912, physicist J. J. Thomson, who had been Rutherford’s research advisor a few years earlier, subjected neon to similar investigations in such discharge tubes with his research assistant F. W. Aston at Cambridge University (42). Contrary to expectations, neon produced two mass peaks. With the crude technology of the time, these two signals appeared as parabolas around mass 20 and 22 with the 22 signal being much fainter. The ratio of intensity of the two peaks roughly corresponded to neon’s known atomic mass of 20.2 amu (43, 44). Suddenly, there appeared to be two types of neon atoms, but neon occupied only one spot on the periodic table. Things were not yet certain—it was conceivable that neon’s actual mass was about 20.2, and the parabola at 22 was not neon (45). Isotopy was being approached from the heavy end of the periodic table as well. In 1909 in Berlin, Bruno Keetman confirmed the apparent chemical and physical inseparability of thorium (232Th) and ionium (230Th) (46). These two substances were also shown to have the same spectra (47). Rutherford, who was seeking to use radioactive elements for tracer studies, was fascinated by the problem. He asked Bertram Boltwood, an American who had reported on ionium previously (48), to attempt a separation, which was unsuccessful (49). In the early 1900s, such a separation failure was an unexpected result (26). Only a few years earlier in 1898, the Curies had discovered polonium and radium by methodical and laborious purification of uranium ore. The difficult separation of the rare earth elements, which was a significant challenge during the second half of the 19th century, was also well known. Quite recently in 1907, lutetium had been split from ytterbium (50). So the seemingly intractable separations like thorium, radiothorium (228Th), and ionium were widely expected to eventually succumb to meticulous wet chemical techniques. Additional results of physically inseparable yet chemically similar substances were obtained for radium (226Ra), thorium X (224Ra), and actinium X (223Ra) in crystallization experiments by Swedish chemists Daniel Strömholm and Theodore Svedberg in 1909 (51). 124 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


By the end of 1912, scientists in England, Scotland, Sweden and Germany were approaching a resolution of where the radioelements belonged on the periodic table; the only question was who would reach the correct answer first (26). While possibly sounding incongruous to chemists today, the idea that different substances could appear to be chemically identical was nearly heretical in the early 20th century. Strömholm and Svedberg performed crystallization experiments using solutions of radioelements to ascertain chemical characteristics. Using these methods, they placed radium, thorium X, and actinium X in a single ‘spot’ on the periodic table (51, 52). They were the first to suggest this possibility, which was a startling conclusion. At least some of the radioelements appeared to be identical chemical elements, which seemingly could not be true since they had differing half-lives and sources. The three radioelements Strömholm and Svedberg studied were plainly physically different, yet with apparently identical chemistry. In December 1908, they discussed their idea with Rutherford while he was in Sweden to receive his Nobel Prize. Rutherford was unenthusiastic about the idea. Strömholm and Svedberg published their results and speculations, but Svedberg later noted that they were “very careful” about their predictions due to Rutherford’s negativity; Strömholm felt similarly (53). Soddy was reminded of Strömholm’s and Svedberg’s work in 1922. In his Nobel speech (39), Soddy translated their 1909 work as “the elements of the scheme were mixtures of several homogeneous elements of similar but not completely identical atomic weight.” By substituting ‘periodic table’ for ‘scheme’ in Soddy’s statement, one arrives at a paraphrasing of the modern definition of an isotope. Within this story resides an important lesson—never let another scientist, no matter how prominent, discourage you from advocating for your exciting results. Svedberg would win a Nobel Prize in 1926, but was recognized for his contributions to disperse systems including colloids, not his early work on isotopes. The accumulated evidence on chemically inseparable elements finally led to Soddy’s proposal that isotopes are “chemically identical elements of the same nuclear charge”. The name isotope, or ‘same place’, had been suggested at a dinner party by Margaret Todd, a Scottish medical doctor, distant relative, and friend of Soddy’s (54). Soddy’s Nobel Prize speech provides an excellent, and very readable, summary of the events leading up to the realization of isotopy. Neither Soddy nor the two Swedes used Henry G. J. Moseley’s famous, but then in 1913 new results, which established that atomic numbers mathematically corresponded to x-ray emission wavelengths (55), and thus each element has a unique atomic number. We now recognize that number as the number of protons in the nucleus. Yet all the early conclusions on isotopy were derived from chemical similarities and atomic masses without relying on knowledge of the internal structure of the atom. In a series of letters to Nature in 1913, Soddy explained isotopy and provided the name isotope. He also defended the new theory against Rutherford’s assertion that isotopy depended on Moseley’s atomic numbers, though the two properties eventually would be shown to be intimately related.

125 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


The Unanswered Question: The Physical Reality of Isotopes The isotope proposal by Soddy left a fundamental question unanswered: how could objects that were physically different have the same chemical properties? This became a fiercely debated topic and research focus for the next decade. Georg von Hevesy’s work using ion mobility and valences as a proxy for atomic mass helped to show that many of the then-new radioelements were likely variants of the same elements (56), which supported Soddy’s isotopy hypothesis. Von Hevesy was from a wealthy Hungarian family, and studied at several institutes in Budapest, Berlin and Freiberg. After obtaining his PhD, he worked with Fritz Haber in Karlsruhe. Although today Haber is mostly known for his work with catalysts, he wanted von Hevesy to work on new projects. Having found a deficiency of electrochemistry experts in Germany, Haber enthusiastically agreed that von Hevesy should go work with Rutherford in Manchester where he arrived January 1911. Rutherford put von Hevesy to work separating radium-D (210Pb) from lead as Rutherford was interested in isolating radioelements for use as tracers. Von Hevesy worked on the project for a year before realizing the futility of his efforts (57). Kazimierz Fajans’ and Frederick Soddy’s aforementioned radioactive displacement law had explained the effect of α and β decay of nuclides correctly in 1913 (24, 58). Soon afterwards, von Hevesy took a 3-month break from Manchester to visit Friedrich Paneth in Vienna (59). They published a study showing that radium-D (210Pb) and thorium-B (212Pb) were identical to ordinary lead by demonstrating that all these nuclides deposited electrochemically identically (60). Paneth had abandoned a PhD in organic chemistry to focus on radioactivity, joining the Radium Institute in Vienna in 1912. He became a major figure in chemistry in the first half of the 20th century working in organic and inorganic chemistry, radiochemistry, electrochemistry, and eventually studying meteorites. In 1927, he briefly believed he had discovered what we would label now as cold fusion (61). A Protestant of Jewish ancestry, he did not return to Germany after 1933. Living in the UK during and after World War II (62), he remained influential, notably proposing guidelines on discovery priority and naming protocols for the many new synthetic chemical elements produced during and after the war (63, 64). Paneth and von Hevesy took Soddy’s position that the isotopes of the same element were chemically identical and interchangeable in their chemistry (‘vertretbarkeit’), a decidedly physics standpoint with which not all chemists concurred (65–67). Perhaps recognizing that Soddy’s publication of the word isotope constituted a strong claim to discovery, Fajans began referring to the set of isotopes of the same elements as a pleiad, a word derived from the Pleiades star cluster (68, 69). The term pleiad did not spread through the radioelement community, which is unfortunate since even now we still do not have a single designation for ‘all the isotopes of a single element’ (e.g. the ‘carbon pleiad’ or the ‘tin pleiad’). Fajans emphatically asserted that isotopes could not be chemically identical; he argued from a thermodynamic point of view that two substances of different masses could not have the same chemical properties (70, 71), but his view was roundly rejected at the time by Soddy, Paneth, von Hevesy and others (71). 126 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


With the non-radioactive elements, the concept of isotopes also continued to coalesce, but at a slower pace. After identifying the two mass signals in neon, Aston attempted a physical separation of the two substances with Thomson (43). He worked on this until the outbreak of World War I. Thomson was confident of Aston’s eventual triumph (72), who achieved a very small, but measurable, success using a fractional distillation technique. After the war, Aston published a paper with Frederick Lindemann reviewing possible methods of separating isotopes, and concluded that theoretically many methods should work (73, 74). Aston developed a ‘mass spectrograph’ and attempted to separate atoms by mass (75), since the earlier observations by Aston’s mentor Thomson were insufficient to prove that the 20 and 22 mass peaks observed were not two separate elements, instead of neon isotopes (76). In 1920, Aston asserted that the “practical chemistry” of an element was not affected by isotopy—that is, there was no difference in the chemical behavior (77). This put his views in alignment with the physics-first group of Paneth, von Hevesy, and Soddy, but in opposition to Fajans. He added that isotopy had provided “a very desirable simplification into the theoretical aspects of mass.” Aston was referring to chemistry finally having an explanation of atomic weights, but the statement reflected the feeling at the time that isotopy was already on the edge of chemistry. Isotopy was an element of physics that could explain a fundamental concept in chemistry, atomic weights. Aston’s expertise in mass led to his involvement in the International Committee on Chemical Elements, which was known as the Committee on Atomic Weights before World War I. The renaming reflected a greater domain of responsibility. The new postwar committee would review not only atomic weights, but also make tables of isotopes and radioactive elements. The effect of Aston’s physics-first interpretation was quickly seen in the Committee’s first post-war report in 1923 that included isotopes (78), whereas, the previous report from 1918 had not (79). Aston received the 1922 Nobel Prize in Chemistry, but his Nobel lecture was dominated by topics that might be seen as physics, not chemistry by modern interpretation (80). The topics presented included the design and implementation of his mass spectrograph, the resolution of isotopes, and the interpretation of mass spectra on photographic plates as an independent confirmation of the atomic weights of elements. Aston’s speech continued to diverge from classic chemistry. His mass spectra unequivocally showed that that the mass of four hydrogen atoms was greater than the mass of a helium atom. Aston rationalized that since hydrogen contained only one proton and electron, helium must contain four of each (neither deuterium, nor the neutron had been identified yet). The masses of helium and hydrogen were different, but where did the extra mass go? Aston realized that the mass was converted into energy according to Einstein’s E = mc2. The mass spectrometer had shown the relative energy that could be created by fusing hydrogen into helium. The implications of this observation were profound as astronomer Arthur Eddington had recently pointed out that this could account for a very long-lived, stable sun, which was fueled by fusing hydrogen into helium (81). Aston’s reaction to this revelation of the energy locked in atoms was mixed. In his Nobel lecture, he spoke of the near limitless energy 127



that might one day be obtained; however, he closed darkly, expressing worry that releasing that energy might literally turn the Earth into a star by igniting all available hydrogen (80). In similar lectures given in subsequent years, he was sometimes more positive, pointing out that fusing the hydrogen in a cup of water could power a large ocean liner back and forth across the Atlantic Ocean (82). This was largely where chemistry stopped. The energy inside the atom was below the size domain of chemistry. Unlocking the power contained inside the atom, and doing so in a controllable way became the domain of physics. Nuclear energy research largely remains under the umbrella of physics today. After a generation of chemists that included Rutherford, Aston, the Curies, Soddy, etc., traditional chemistry displaced studies of radioactivity. Even now, there are many seeking to create new elements and push the boundaries of the periodic table, but most of those scientists trace their lineage to physics, not chemistry.

Philosophical Interregnum Apparently, by historical accident, high, wide-ranging praise is given to those who discover or create a new element, but not a new isotope. Praising element discoveries prioritizes the chemistry over the physics (83), but the technical and practical challenges of creating a new element are not necessarily greater than the challenges in creating an exotic new isotope of a known element. So what can chemists do with the proliferation of isotopes? Without the status of elements, are isotopes even a topic chemists should consider? This philosophical debate that is almost forgotten today, defined chemistry in the early 20th century, and drew a line separating chemistry and physics. Today, isotopes are commonplace in chemistry as tools; whether elucidating reaction mechanisms with isotopically labeled substances (84), or in the history and source encoded in isotopic ratios used in geosciences. Kazimierz Fajans was the leading proponent of the belief that the periodic system would not survive the discovery of isotopes. A view that unsurprisingly conflicted with influential supporters of the traditional periodic table, where each element was a substance that could not be further separated by chemical means—the view held by Paneth. Fajans insisted that if isotopes were not physically identical, the chemical properties could not be identical. To suggest otherwise was contrary to physics. Fajans’ belief that the periodic table had become obsolete, though nearly forgotten now, was a significant debate in chemical history. If Fajans’ arguments had been successful, chemistry might have evolved into something like modern nuclear chemistry or nuclear physics. The past century of chemical science advancement would have happened, but may have done so under a different banner, with chemistry relegated to chasing isotopes, and the elementary particles that make up protons and neutrons. If history had followed this course, the awarding of the 1908 Nobel Prize in Chemistry to Rutherford, an ardent physicist, would have seemed remarkably prescient. Paneth viewed Fajans’ position as dangerous. Paneth believed that treating the ever-increasing number of isotopes as elements threatened the foundations of 128



chemistry. He wrote in 1925 that an “element is a substance of which all atoms have the same nuclear charge", which by then was becoming a more common view. Ultimately, Aston and the Atomic Weights committee implicitly, and sometimes explicitly, promoted Paneth’s position that all isotopes of an element were chemically identical (85). Fajans’ viewpoint, already fading from mainstream chemistry, had never gained significant traction. Paneth and von Hevesy’s 1913 electrochemistry experiments with lead isotopes had helped establish the field of radioactive tracers, which they extended to other elements over the next few years (86–88). As far as they could measure, the two isotopes of an element behaved identically, and so they concluded that chemistry does not depend on isotopy. Paneth applied this knowledge to a variety of experiments such as the detection of polonium and bismuth hydrides using tracers (89). Much of the instinct of chemists to ignore differences in isotopic chemistry traces back to these early experiments. In this view, only the nuclear charge, the atomic number, impacted chemical properties. Subtle problems began to emerge a few years later, particularly when Harold Urey discovered deuterium in 1931 with George Murphy, and Ferdinand Brickwedde (90, 91).

Deuterium: An Opportunity To Reevaluate Isotope Chemistry Arthur Lamb and Richard E Lee of New York University already had reported that the density of water varied from sample to sample in 1913, another key event in the realization of isotopy in 1913 (92). Unbeknownst to Lamb and Lee, they had observed changes caused by varying amounts of a heavier isotope of hydrogen—deuterium—in the water. The actual discovery of the heavy isotope of hydrogen would take much longer. Although many isotopes were discovered in the 1920s using mass spectrometric methods (93), the most common elements in living things—carbon, oxygen, nitrogen, and hydrogen—remained apparently monoisotopic (45). This began to change in 1929 with the report of 18O, found by atmospheric absorption spectroscopy (94), and the discovery of the much less abundant 17O shortly afterwards (95). Harold Urey earned his PhD in physical chemistry under the direction of Gilbert Lewis, and worked at the Niels Bohr Institute in Copenhagen before undertaking an independent career at Columbia. Urey and Murphy first photographed emission lines from deuterium in the expected locations, which they had recalculated for the position of the Balmer emission lines of hydrogen, if hydrogen had a mass of two. To confirm the observation, they needed a sample of hydrogen containing a higher concentration of deuterium. In practice, they needed to separate some of the light hydrogen from the heavier hydrogen—something that the original definition of isotopes had presumed to be impossible. Urey contacted Brickwedde to provide 5-6 L samples of hydrogen that had been evaporated down to 2 mL. The assumption was that the lighter hydrogen isotope would evaporate faster, increasing the concentration in the residue. The assumption was correct, and spectra of the residual hydrogen showed much stronger emission 129



lines of deuterium. These experiments, followed by realization of its different properties from protium, should have nullified the assertions of Soddy and Aston that the chemistry of all isotopes is identical. Fajans’ thermodynamic arguments from 1914, which implied that deuterium could not have the same chemistry as protium, were closer to reality. Urey would receive the Nobel Prize in Chemistry for his discovery of deuterium in 1934 (96). Deuterium’s usefulness as a tracer was immediately obvious. Following the discovery, Urey gave von Hevesy a few liters of deuterium-enriched water, allowing the latter to test an idea he had discussed with Moseley in April 1913: what happens to a cup of tea after it is drunk? There is little excreted for some hours, but because the admixture of a known amount of heavy water will disperse throughout the body relatively quickly, it allows for an easy calculation of the total water in a person’s body by measuring the D2O concentration in blood plasma. A cringe-worthy example that deuterium does not behave like hydrogen from a chemistry point of view is the 1961 paper, “Physiological effects of deuterium on dogs” (97). The term ‘deuterated dog’ did not however become commonplace. The notion that isotopes are chemically inseparable, a convenient approximation that sometimes still appears in textbooks, was championed for many years by Frederick Soddy. That position faced its biggest test with the discovery of deuterium. Soddy vehemently maintained that deuterium could not be an isotope of hydrogen. Soddy stuck to chemical inseparability as a criterion for isotopes, and therefore refused to recognize deuterium as an isotope of hydrogen. For Soddy, deuterium was a species of hydrogen with different atomic weight, but not an isotope of hydrogen (98). Soddy’s original definition of an isotope included chemical inseparability from the other isotopes of the same element, but deuterium was separable. As additional studies accumulated, other isotopes were separated with varying degrees of difficulty from pleiads. This fundamentally changed the orthodoxy of isotopy. As Urey’s collaborator Brickwedde wrote “before the discovery of deuterium, chemical properties were generally supposed to be determined by the […] extranuclear electrons, quantities that are identical for isotopes of the same element. It had not been realized that chemical properties are also affected—but to a lesser degree—by the mass of the nucleus.” Perhaps Soddy’s insistence that isotopes in a pleiad could not have different chemistry, as well as the influence of power brokers like Paneth, Aston, von Hevesy and others, led to the concept of all isotopes having identical chemistry only being abandoned tacitly (99). The deuterium case immediately demonstrated the different properties of light and heavy hydrogen. Of course, deuterium is an extreme case as the mass is doubled compared to protium. Urey’s experiments showed that deuterium did not behave the same as protium in chemical reactions. So, although Fajans’ hypothesis on isotopes exhibiting different chemistry was vindicated, we still teach that all isotopes of an element have the same chemistry with sometimes glibly adding that “the effects are largest and important for light atoms, like hydrogen”. Yet the tiny variations in isotope chemistry lead to tiny variations in relative abundances in a compound. These changes can sometimes tell the history and source of samples, and form the basis of huge swaths of modern geosciences. If Paneth were strictly correct, entire fields of research would not exist. In philosophical terms, Paneth’s 130


viewpoint is realism, where practical chemical similarities of isotopes of a single element control its chemistry. Fajans’ interpretation is reductionist, where physical differences between isotopes of an element must inform its chemical behavior (100).


Geochemistry: Where Isotopes Are Chemically Distinct In contrast to classical chemists, geochemists have been more accepting of isotopes having different chemistry. The uses are too myriad to describe here comprehensively. The potential use of isotopes as geochemical process tracers was realized early, but practical application of the differing chemistry of isotopes remained elusive for decades. Theodore Richards received the 1914 Nobel Prize in Chemistry for his work on atomic weights. He quickly realized that isotopy meant that atomic weights could vary. In his Nobel lecture, he noted “Because of this recent knowledge [isotopy], some investigators seem to believe that these so-called ‘constants of nature’ are of less significance than of old, since some atoms may be transitory, and since a given substance may have different atomic weights according to the circumstances of its life history (101).” Herein, Richards anticipates the existence of isotopic fractionations in nature. In 1921, J. J. Thomson hypothesized that chlorine samples formed in other times would have abnormal densities (alternate atomic weights). Abnormal densities would arise from varying ratios of 35Cl and 37Cl that were different from what Aston’s mass spectrograph had revealed (73). When Thompson made this proposal, there was still some doubt about the purity of chlorine with the aforementioned proposal that 37Cl was actually 35Cl with two really strongly bonded hydrogen atoms. This sounds like an absurd notion today, but was a legitimate concern in 1921. Thomson thought that it “ought not to be a very difficult matter to detect a change of 0.5 % in the atomic weight. So that it seems to me that an experiment which apparently does not present prohibitive difficulties—though you never can be sure until you start these things—would enable one to definitely to settle the question of whether chlorine was a mixture or not.” He further pondered “Is it likely that various samples of chlorine should be of such invariable composition [worldwide]?” Thomson imagined processes that could cause such fractionation of chlorine (102). In modern terminology, Thomson seems to be predicting the existence of and describing the kinetic isotope effect (KIE). The KIE measures the reaction rate difference between a molecule containing a lighter isotope versus one containing a heavier isotope, for example the ratio of the reaction rate of 12CO2 versus 13CO2. Thomson vastly underestimated the difficulty of his proposed experiment, because the KIE fractionations are much smaller than he anticipated. The differences in chlorine composition in natural materials are on the per mille scale rather than the percent scale (58). Finally in 1946, Harold Urey presented a brief review of instances where isotopes other than deuterium behaved differently (60). This pivotal presentation marked the beginning of stable isotope studies, and reflected a shift from seeing 131 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


isotopes as chemically identical to chemically distinct. Although the chemical distinction of deuterium and protium had been known for years by then, now the whole periodic table was in play. Isotopes of an element no longer had to be relegated to being mere tracers that behaved identically; the chemical differences could also be exploited. In natural and artificial systems, a tiny preference for one isotope or the other in a reaction that is repeated many times will eventually lead to isotope fractionation from reactants to the product. Urey was involved in some of the greatest early breakthroughs, such as in 1951 when he led a paper describing paleotemperatures during the Jurassic and Cretaceous using the 18O isotope in calcium carbonate of marine belemnite fossils. These fossils also became the basis of the 13C reference scale (now based on an updated standard known as Vienna Pee-Dee Belemnite) (59). The most common isotope effect observed is the aforementioned KIE, where the outcome of a reaction is impacted by the mass of the isotopes involved. For instance, carbon compounds produced via photosynthetic processes generally are depleted in the heavier 13C compared to 12C, because plants preferentially use 12CO2. Thus, while present-day (2017) northern hemisphere background atmospheric CO2 has a δ13C ≈ −8.4 ‰ (103), δ13C in plants have δ13C < −8.4 ‰ because they take up 12CO2 at a slightly higher rate than 13CO2. Additionally, the atmospheric value is slowly becoming more depleted in 13C due to the burning of fossil fuels, which add CO2 with δ13C ~−30‰ to the atmosphere. The δ13C in plant material also depends on the particular metabolic pathway the particular plant uses (104). The origin of the Earth’s moon has been a longstanding problem in planetary science with no shortage of explanations and hypotheses. Isotopic studies comparing lunar and terrestrial materials have become a favorite method of supporting—or disproving—lunar formation hypotheses. Similarities in oxygen isotopes in Apollo lunar samples compared with terrestrial samples have been used as arguments for a giant impact formation explanation of the moon since the theory gained favor in the 1970s (105). More recently, similar arguments were made using hydrogen isotopes (106). Of course, isotopes have not resolved everything, and lingering doubts remain since some results show the isotopic distributions of the earth and moon to be too similar. How much did the proto-Earth mix with the impacting protoplanet Theia, which is believed to have been isotopically distinct from the proto-Earth (107, 108)? Some recent studies have attempted to explain the isotopic similarities by invoking a series of smaller moonlet-forming impacts (109). Few isotopic corners have been left unturned in the search. Even titanium isotopes have been invoked (110, 111). The KIE is not the only effect acting on isotopic distributions in nature. In the early 1980s, photochemical reactions on oxygen were demonstrated to have a mass-independent effect on the fractionation of oxygen isotopes. This was suggested to potentially explain the different oxygen isotopic ratios seen in the earth and moon, compared to carbonaceous chondrite meteorites, which were thought to represent the pre-solar system nebula where the planets formed (112). Later studies of solar wind particles have shown that the solar oxygen isotope distribution is also different, which gives theorists more data to analyze (113). 132


The aforementioned, and more obvious to early workers, use of isotopes is tracers—based on the implicit assumption that all isotopes of an element have identical chemistry—dates to 1913 (114), and has remained an important tool ever since. Thus, chemists are in the slightly awkward position of explaining the treatment of the isotopes in a pleiad as chemically identical when using them as tracers, and chemically distinct when using them as sources of geochemical history or information (115).


Epilogue: Isotopes No Longer Hidden from Chemistry Many what-if scenarios can be envisioned around the place of isotopy in chemistry or physics. The early 20th century saw many physicists working on fundamental problems of chemistry—like the origin of atomic weight and radioactivity. This was a remarkably fruitful period of scientific advancement. Given the sheer number of isotopes that would eventually be discovered, the naming of individual isotopes almost certainly would never have lasted more than a few decades. Still we wonder if some of the names of elements would be different. Internecine politics gave us proactinium instead of brevium, radon-220 instead of thoron, and radon instead of emanation; however, ionium was never going to displace the name thorium. These are the smallest of potential speculations. The greater question is what would chemistry have been, had chemistry’s domain been extended to the subatomic. Instead, the subatomic became part of physics, and the early 20th century became an anomalous time in the history of chemistry, when many of the dominant figures were physicists, and physicists flocked to study chemistry problems. Chemists, like Paneth, moved from traditional synthetic chemistry research to follow the new excitement of radioactivity and later, isotopes. For a time, Nobel Prize decisions reflected this uncertainty about chemistry’s direction, but after the excitement passed, the dividing line between physics and chemistry was re-clarified by the mid-1930s. Although isotopes quietly have lain beneath the periodic table for nearly a century, better understanding of isotopes has increasingly changed one of the table’s most visible parts, the atomic weights. By 1951, IUPAC had added a disclaimer about the accuracy of the atomic weight of sulfur because of isotope variations. The advancement of analytical and separation technologies slowly added more footnotes to the weights of additional elements (116). The last three IUPAC atomic weight revisions in 2009, 2011, and 2013 (117–119), have updated the atomic weights of a dozen elements from a constant single value to a range. The weight range reflects the variations in element isotopic compositions found in in terrestrial materials. A commonly overlooked distinction established by Aston in 1923 is that the atomic weight is the average atomic mass of the distribution of an element’s isotopes on Earth. The atomic mass is a precise mass of a specific nuclide (78). This distinction limits the accuracy at which the atomic weight of a non-monoisotopic element can be known. It does not appear that Aston, or anyone working in the early 20th century, anticipated atomic weights becoming ranges, rather than a specific value. Instead, they emphasized finding average Earth 133 Benvenuto and Williamson; Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications ACS Symposium Series; American Chemical Society: Washington, DC, 2017.


materials from which to determine the correct atomic weight. To use Fajans’ terminology, the average mass of an element’s pleiad is the atomic weight. Only 26 elements are monoisotopic, having only one stable or long-lived isotope. As was succinctly stated in 2011, shifting knowledge of isotope compositions mean that the atomic weights are no longer constants of the universe (120), a callback to F. W. Clarke’s 1882 description of atomic weights as “constants of nature” (121). Although chemistry and physics seem to have drifted apart since the announcement of isotopy in 1913, the “mutual dependence of chemistry and physics is clearly visible” in the creation of new, superheavy isotopes and elements (122). As various isotopes were investigated more thoroughly with new and more sensitive techniques and instruments, Fajans’ reductionist assertion that isotopy must affect chemistry has proven correct. Fajans was considered several times for a Nobel Prize, and in 1924 he was such a favorite that the Stockholm newspaper Svenska Dagbladet confirmed his win just before the announcement that no prize would be awarded in 1924. Having missed out on the most prestigious prize a scientist can receive, as well as having his contributions to the discovery of element 91 and the radioactive displacement laws diminished, perhaps Fajans would find solace that his views on isotopy have entered mainstream chemistry orthodoxy. Isotopes do influence chemistry—even on its most central icon, the periodic table.

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