Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications 9780841232556, 0841232555, 9780841232525

Table of contents :
Content: Lost elements: the all-American errors --
The periodic table of the elements: a review of the future --
How the periodic table of available elements shaped natural history --
An elemental history of the early universe --
Building classes of similar chemical elements from binary compounds and their stoichiometries --
It's all in the sludge: elements that are always by-products --
Chemistry's decision point: isotopes --
Analytical methodologies for arsenic, selenium, and mercury: a historical perspective --
Recent advancements in the radiochemistry of elements pertaining to select nuclear materials and wastes --
Element 118: teaching a new element to new students --
Chemicals, their element names, and their place in society
also known as "Why did I choose to name my organization Palladium Science Academy?".

Citation preview

Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications

ACS SYMPOSIUM SERIES 1263

Elements Old and New: Discoveries, Developments, Challenges, and Environmental Implications Mark A. Benvenuto, Editor University of Detroit Mercy Detroit, Michigan

Tracy Williamson, Editor ACS Division of Environmental Chemistry Crofton, Maryland

Sponsored by the ACS Division of Environmental Chemistry, Inc.

American Chemical Society, Washington, DC Distributed in print by Oxford University Press

Library of Congress Cataloging-in-Publication Data Names: Benvenuto, Mark A. (Mark Anthony), editor. | Williamson, Tracy C., 1963- editor. | American Chemical Society. Division of Environmental Chemistry. Title: Elements old and new : discoveries, developments, challenges, and environmental implications / Mark A. Benvenuto, editor, (University of Detroit Mercy, Detroit, Michigan), Tracy Williamson, editor (ACS Division of Environmental Chemistry, Crofton, Maryland) ; sponsored by the ACS Division of Environmental Chemistry, Inc. Description: Washington, DC : American Chemical Society, [2017] | Series: ACS symposium series ; 1263 | Includes bibliographical references and index. Identifiers: LCCN 2017048995 (print) | LCCN 2017052125 (ebook) | ISBN 9780841232525 (ebook) | ISBN 9780841232556 Subjects: LCSH: Chemical elements. | Chemical elements--History. | Chemistry--History. | Periodic table of the elements. Classification: LCC QD466 (ebook) | LCC QD466 .E455 2017 (print) | DDC 546--dc23 LC record available at https://lccn.loc.gov/2017048995

The paper used in this publication meets the minimum requirements of American National Standard for Information Sciences—Permanence of Paper for Printed Library Materials, ANSI Z39.48n1984. Copyright © 2017 American Chemical Society Distributed in print by Oxford University Press All Rights Reserved. Reprographic copying beyond that permitted by Sections 107 or 108 of the U.S. Copyright Act is allowed for internal use only, provided that a per-chapter fee of $40.25 plus $0.75 per page is paid to the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, USA. Republication or reproduction for sale of pages in this book is permitted only under license from ACS. Direct these and other permission requests to ACS Copyright Office, Publications Division, 1155 16th Street, N.W., Washington, DC 20036. The citation of trade names and/or names of manufacturers in this publication is not to be construed as an endorsement or as approval by ACS of the commercial products or services referenced herein; nor should the mere reference herein to any drawing, specification, chemical process, or other data be regarded as a license or as a conveyance of any right or permission to the holder, reader, or any other person or corporation, to manufacture, reproduce, use, or sell any patented invention or copyrighted work that may in any way be related thereto. Registered names, trademarks, etc., used in this publication, even without specific indication thereof, are not to be considered unprotected by law. PRINTED IN THE UNITED STATES OF AMERICA

Foreword The ACS Symposium Series was first published in 1974 to provide a mechanism for publishing symposia quickly in book form. The purpose of the series is to publish timely, comprehensive books developed from the ACS sponsored symposia based on current scientific research. Occasionally, books are developed from symposia sponsored by other organizations when the topic is of keen interest to the chemistry audience. Before agreeing to publish a book, the proposed table of contents is reviewed for appropriate and comprehensive coverage and for interest to the audience. Some papers may be excluded to better focus the book; others may be added to provide comprehensiveness. When appropriate, overview or introductory chapters are added. Drafts of chapters are peer-reviewed prior to final acceptance or rejection, and manuscripts are prepared in camera-ready format. As a rule, only original research papers and original review papers are included in the volumes. Verbatim reproductions of previous published papers are not accepted.

ACS Books Department

Contents Preface .............................................................................................................................. ix 1.

Lost Elements: The All-American Errors ............................................................. 1 Mary Virginia Orna, Marco Fontani, and Mariagrazia Costa

2.

The Periodic Table of the Elements: A Review of the Future ............................ 41 Paul J. Karol

3.

How the Periodic Table of Available Elements Shaped Natural History .......... 67 Benjamin J. McFarland

4.

An Elemental History of the Early Universe ....................................................... 83 E. Prasad Venugopal

5.

Building Classes of Similar Chemical Elements from Binary Compounds and Their Stoichiometries ..................................................................................... 95 Guillermo Restrepo

6.

It’s All in the Sludge: Elements That Are Always By-Products ....................... 111 Justin Pothoof, Grace Nguyen, and Mark A. Benvenuto

7.

Chemistry’s Decision Point: Isotopes ................................................................. 119 Brett F. Thornton and Shawn C. Burdette

8.

Analytical Methodologies for Arsenic, Selenium, and Mercury: A Historical Perspective ............................................................................................................ 141 Larry Kolopajlo

9.

Recent Advancements in the Radiochemistry of Elements Pertaining to Select Nuclear Materials and Wastes ................................................................. 173 Eric S. Eitrheim, Andrew W. Knight, Michael K. Schultz, Tori Z. Forbes, and Andrew W. Nelson

10. Element 118: Teaching A New Element to New Students ................................ 195 Justin Pothoof, Grace Nguyen, Dawn Archey, E. Prasad Venugopal, and Mark A. Benvenuto 11. Chemicals, Their Element Names, and Their Place in Society. Also Known as “Why Did I Choose To Name My Organization Palladium Science Academy?” ............................................................................................................. 203 George W. Ruger Jr.

vii

Editors’ Biographies .................................................................................................... 217

Indexes Author Index ................................................................................................................ 221 Subject Index ................................................................................................................ 223

viii

Preface “The time has arrived when a knowledge of physics and chemistry forms as important a part of education as that of the classics did two centuries ago.” Dimitri Mendeléeff, -from The Principles of Chemistry, P.F. Collier & Son, 1902. This current volume is the result of a symposium held during the 252nd American Chemical Society National Meeting, in August 2016 at Philadelphia, Pennsylvania. The symposium provided an excellent venue to discuss the heavy elements that had been recently verified and named, and that had thus completed the seventh row of the Periodic Table – with oganesson being the name of the final element, which had formerly been known as Element 118 or Eka-radon. One can say honestly that it is a rare moment in time when a row of the Periodic Table of the Elements is completed. The Periodic Table of the Elements remains a living, growing document that attempts to map out all of the most primal matter known to humankind. Interestingly, Mendeleev’s statement, above, is as important today as it was when he penned it over a century ago, certainly when applied to the periodic table. Accordingly, we felt a volume from this symposium would help preserve our knowledge and understanding of the Periodic Table of the Elements as it exists at this moment in time, a moment in which the seventh row of the table had just been completed in terms of verified syntheses of the super heavy elements, and a point in time at which IUPAC had proposed names for them. While all the seminar speakers at the symposium were not able to convert their material into chapters, we are flattered and honored to have excellent chapters in this volume. Authors and topics include the following: Professor Mary Virginia Orna presents an excellent chapter on how elements are discovered and named – specifically some of the errors that have occurred (Chapter 1). This is hardly a litany of failure; on the contrary, it brings to life the struggle scientists go through when peering into the unknown, and trying to determine whether a substance is pure or mixed, or new or already known. It shows how science and scientists self-correct as knowledge is gained. Just as importantly, it is a fascinating read. We also have in this volume a chapter by Professor Paul Karol with the colorful name “The Periodic Table of the Elements: A Review of the Future” (Chapter 2. This chapter does an excellent job of making predictions based on how the table has developed in the past. Professor Karol tackles questions on if ix

and how the table will continue to evolve and grow, and questions what sort of new chemistry will emerge as we continue our quest to push the limits of the table. Dr. Benjamin McFarland provides an interesting and informative chapter that explains the relationship between the periodic table and natural history (Chapter 3). It may be fair to say that many of us who feel quite familiar with the periodic table do not possess an equal familiarity with how life evolved in the presence of concentrations of specific elements. Professor McFarland takes a highly holistic view, and makes numerous connections. It is a physicist, Professor Prasad Venugopal, who provides us with a chapter that could glibly be called a physicist’s view of the elements and the universe, although that would be both a disservice and an extreme oversimplification (Chapter 4). His chapter looks at chemistry far beyond that which occurs on our planet, provides a much greater context for how the elements exist through nucleosynthetic pathways, and discusses current theories on the existence of known matter and postulated dark matter. Dr. Guillermo Restrepo has contributed an excellent chapter on compounds, stoichiometries, and chemical similarities, and presents a method of study that should continue to be useful as the periodic table continues to expand and evolve (Chapter 5). This chapter should be especially useful to educators who are trying to teach an understanding of the periodic table of the elements, and of chemical reactions. One of our editorial duo (Benvenuto) and his researchers have produced a chapter discussing the distribution of elements, and their usefulness in modern life (Chapter 6). Titled: “It’s all in the sludge: Elements that are always byproducts,” the aim of the chapter is to connect the fact that certain elements are widely distributed throughout the planet in low concentrations, yet have become important in modern life. Overall the aim is to increase the understanding of what might be called the uneven-ness of elemental abundance, something not shown or implied by current representations of the periodic table. Professors Burdette and Thorton have contributed a discussion of the understanding of isotopes, and placed their development and nomenclature in the broader context of history and the development of chemistry and physics (Chapter 7). This chapter is an excellent treatment of how the understanding of isotopes grew, and how the two fields, chemistry and physics, met, co-mingled, and treated this new knowledge. The history of how arsenic, selenium, and mercury have been detected is the subject of the chapter authored by Professor Larry Kolopajlo (Chapter 8). This chapter is especially compelling, as arsenic and mercury have been known in some form for centuries, often-times for their use as some type of poison. The chapter does an excellent job of exploring how the identification and understanding of these elements throughout history has been linked to effects on humans and our larger environment. Professor Eitrheim and his colleagues have presented in their chapter a discussion of the heaviest, radioactive elements and of the knowledge that radiochemists bring to their uses in society (Chapter 9). They discuss the need for people trained in this discipline, as well as the fate of heavy, often radioactive x

elements over long periods of time. The chapter also presents thoughts on possible future uses of such elements and the materials, usually alloys, made from them. In a further chapter, Benvenuto and his group members have also weighed in with a discussion of how the newest, heaviest elements might behave if made in macroscopic amounts (Chapter 10). Students are often called on to memorize apparently disconnected sets of facts, such as the names of all the elements, but in the process are not told much about them. This chapter presents a method whereby students can add depth to their understanding of the elements by making predictions of their physical properties, even if the element in question only exists in minuscule amounts. Finally, George Ruger has provided us with a very different yet equally interesting chapter on using an element name in naming a company, and the public’s perception of it (Chapter 11). His insight into the public’s perception of chemistry and science throughout the ages is both a very good read and food for thought as we live our lives, develop careers, and interact with a public that is not always well informed about how chemistry and science improves the overall quality of life on Earth, or that has pre-conceived notions of how the science negatively impacts them. As far as the production of these chapters and the contributions of the authors, there are many people to thank. It is not a trivial matter to produce a chapter for such a volume, and we appreciate the effort our authors have dedicated to pursuing this project. As well, we appreciate the efforts of the reviewers, who give liberally of their time with no reward beyond the knowledge that they are contributing to what we believe will become an important, timely volume. And we also owe a debt to the ACS Symposium Books staff, especially Rachel Deary, Elizabeth Hernandez, and Arlene Furman, all of whom have helped ensure what we have produced is to the ACS’ high standard.

Mark A. Benvenuto Department of Chemistry & Biochemistry University of Detroit Mercy Detroit, Michigan 48221

Tracy C. Williamson Immediate Past Chair ACS Division of Environmental Chemistry 2527 Dog Leg Drive Crofton, Maryland 21114

xi

Chapter 1

Lost Elements: The All-American Errors Mary Virginia Orna,*,1 Marco Fontani,2 and Mariagrazia Costa2 1Department

of Chemistry, The College of New Rochelle, New Rochelle, New York 10805, United States 2Department of Organic Chemistry, University of Florence, Via della Lastruccia 13, Sesto Fiorentino, Firenze, Italy, 50019 *E-mail: [email protected]. E-mail: [email protected].

In the development of the periodic table, many wrong turns were taken and falsely claimed element discoveries discredited. Among these were the “American” elements, those “discovered” by three Presidents of the American Chemical Society, and those named after states of the Union: North Carolina, Illinois, Alabama, and Virginia. This chapter narrates their detailed stories.

Introduction Probing the nature of an elemental substance has never been an easy task. Getting down to the stuff that the material universe is made of, with respect to number, nature, and nomenclature, has baffled the best of minds through the millennia. Aristotle’s (384-22 BCE) philosophical framework dominated Western thought for over two thousand years, a system that banished number, weight, and measure to insignificance. Though the Aristotelians were always intent on tidying up the world, the world as they knew it came to an untidy end, in western culture, around the year 1750. Prior to that year, much occurred to demolish Empedocles’ (495-30 BCE) four-element hypothesis. Fast forward to the sixteenth century: Paracelsus (1493-1541) claimed that the four earthly substances (earth, air, fire, water) had three spiritual sources, namely, mercury, salt, and sulfur. To complicate matters, Robert Boyle (1627-91), in his book The Sceptical Chymist (1) contended that our knowledge of the elements must be enlarged upon by experiment, not mere speculation. He also attacked the four-element hypothesis ((1), pp. 33-38, passim), opening up the possibility that the number of elements © 2017 American Chemical Society

might not be so limited. Thus, by the dawn of the seventeenth century, practicing scientists began to think in terms of discovering new elements. A century later, Antoine Laurent Lavoisier (1743-94) advanced the germinating modern idea of “element” still further by casting doubt on the Greek four-element hypothesis in his famous Traité élémentaire de chimie (2) (Figure 1): We would be surprised not to find in an elementary treatise of chemistry a chapter on the constituent and elementary parts of bodies; but I note here that our tendency to want that all bodies in nature are composed of only three or four elements is due to a bias that comes originally from the Greek philosophers. The introduction of four elements, which, by the variety of their proportions, comprise all the bodies that we know, is pure conjecture, imagined long before we had the first notions of experimental physics and chemistry.

Figure 1. Cover of the 1789 edition of Antoine Laurent Lavoisier’s Traité élémentaire de chimie. Reproduced from reference (2). Subsequently, Lavoisier flew in the face of Aristotelian philosophy by giving supreme importance to observation, experimentation, and measurement in studying the properties and composition of material substances. He believed, and demonstrated, that elements survived in their compounds and could be recovered from them. This analytical approach, that emphasized concrete laboratory experimentation as opposed to theoretical speculation, was the groundwork of the chemical revolution (3). Finally, together with Louis-Bernard Guyton de Morveau (1737-1816), Antoine-François Fourcroy (1755-1809) and Claude-Louis Berthollet (1748-1822), he devised a new chemical nomenclature that was based on the principle that an element’s name should correspond to its composition (4). 2

Thus the lines were drawn regarding number, nature, and nomenclature of the elements. There now seemed to be no limit to the number of elements – it was a virtual “open season” for the eager discoverer. The nature of an element could now be determined using a growing battery of analytical tools. And the naming of an element, just as Adam named the animals that paraded before him in the Book of Genesis, belonged to the person who recognized it first.

Before 1789: Early Errors and Early Elements It is quite obvious that more substances of an elemental nature were recognized from ancient times than the traditional earth, air, fire, and water. Metals found in their free state such as gold, and sometimes sulfur and silver, were known from prehistory. Other metals such as iron and mercury could easily be extracted from their ores; other elements known from antiquity are carbon, copper, tin, and lead. In subsequent centuries, phosphorus, bismuth, arsenic, antimony, and zinc were discovered by alchemists. But the age of discovery of chemical elements began only in the second half of the eighteenth century, around 1750. Since 1789 marks the year of the chemical revolution, we can set it as the arbitrary dividing line between protochemistry and chemistry. During the forty year period between 1750 and 1789, about a dozen new elements were claimed to have been discovered. Very few chemists, if such can be named, were operating at this time, and none of them had access to more sophisticated tools than visual observation of physical properties of substances after treatment with chemical reagents such as acids and bases. At the same time, however, there was a growing interest in metallurgy because the prosperity of a nation was in direct proportion to the productivity of its mines. Mining led to the discovery of new minerals. New minerals had to be analyzed in order to exploit their utility. Chemical analysis revealed large numbers of new substances, many of which were recognized as elements, but many of these substances were mixtures mistaken for elements. So the field was ripe for errors of various kinds. Table 1 gives a sample of some of the erroneous discoveries made during this time period. An examination of Table 1 shows that many countries in Western Europe have a false discovery to their names. Most of the errors involved mistaking a mixture for an element, but there is one instance where a true element, tellurium, was only shown to be so ten years after the fact. Von Reichenstein really did have a problem on his hands, and he sought confirmation from Bergman, who was considered the greatest living chemist at the time. But Bergman died before he could render his opinion, and von Reichenstein simply let the matter drop. Ten years later, Klaproth volunteered to analyze the sample, and indeed, confirmed the presence of an as yet unknown element; he also volunteered the name “tellurium” when the true discoverer failed to come up with a name to his own liking (12, 13). In the next section, we will see how the European propensity for discoving erroneous elements spilled over into America. 3

Table 1. Erroneous Element Discoveries Before 1789 Date

“Element”

Discoverer

Country (5)

Presumed Actual Substance

1777

Terra nobilis (6)

Torbern Olof Bergman (1735-84)

Sweden

Impurity from diamond

1777-78

Hydrosiderum (7)

Johann Karl Friedrich Meyer (1733-1811)

Poland

Alloy of iron and phosphorus

1783

Metallum problematicum (8)

Ferenc Müller von Reichenstein (1740-1825)

Hungary

Tellurium*

1786

Saturnum (9)

Antoine Grimoald Monnet (17341817)

France

Mixture of copper and lead sulfides

1788

Terra adamantina (10, 11)

Martin Heinrich Klaproth (17431817)

Germany

Mixture of corundum and alluminite

1789 – 1869: From the Chemical Revolution to the Periodic Table The 80-year period leading up to the publication of the periodic table was filled with intense chemical activity. The concept of the chemical element was slowly growing in the consciousness of many chemists, although there were still individuals who believed in phlogiston! It was a period of ferment and change. Too great a reliance on technical skills, which were becoming more and more reliable in the isolation of new elements, often led unwary chemists astray. The period was characterized by gross errors in chemical analysis, multiple names for the same “element,” and rediscovery, decades later, of previously rejected elements. Over fifty false claims of element discovery were catalogued during this time. Many carried very fanciful names such as crodonium, pluranium, aurorium, and jargonium. Hardly any country in Europe was exempt from such errors; in addition to the countries noted in Table 1, England, Italy, Russia, and Lithuania, to name just a few, added their names to the list. Meanwhile, a certain degree of caution seemed to pervade the chemical community. Some researchers, hesitantly thinking that they had discovered a new element, declined to give the substance a name, and some even perpetrated jokes under the cloak of anonymity. For example, two gentlemen from Newark, Delaware, communicated the discovery of a supposed new element by the name of brillium to the Washington Post (14). Found in coal ashes, it reputedly gave off more heat than ordinary fuel. During this period, two chemists destined to one day become Presidents of the American Chemical Society, added their names to the list of discoverers of these indeterminate elements. 4

1852: The Element of Friedrich August Genth (1820-93) Friedrich August Ludwig Karl Wilhelm Genth was born in Wächterbach bei Hanau, Germany, on 16 May 1820. Following studies at the Universities of Heidelberg and Gießen (with Justus von Liebig (1803-73)), he earned his doctorate in chemistry in 1845 at the University of Marburg under Robert Bunsen’s (1811-99) supervision. He remained at Marburg as Bunsen’s assistant when he discovered the first nickel oxide crystal (15) (Figure 2). Then in 1848, he emigrated to the United States. In Philadelphia, Pennsylvania, Genth opened one of the first commercial analytical chemistry laboratories in the country. However, two years later, he closed it down in order to take up the position of superintendent of mines at Silver Hill, North Carolina, for a short time.

Figure 2. Book Jacket of a Volume Commemorating Friedrich Genth’s Discovery. Reproduced from reference (15), with permission. 5

While Genth was analyzing a platinum sample from California, he recovered two grains (about 100 mg) of a metal with an intense white color (16, 17). He found that it was malleable and melted immediately in the presence of charcoal and upon treatment with the oxyhydrogen blowpipe. It was attacked by hot hydrochloric acid, hot nitric acid, and with hydrogen sulfide, it yielded a brown precipitate. While Genth published his experimental results immediately, he never gave his presumed discovery a name. Successive studies showed that the substance that fooled him was a mixture of the oxalate and cyanide of platinum and the chlorides of palladium and iridium (18). A complex mixture indeed, but a mixture nonetheless. There is no record of a retraction by Genth, but since his element bore no name, it simply disappeared into the maw of history. Genth, meanwhile, returned to Philadelphia in 1850, reopened his laboratory, and devoted his energies to commercial chemistry, consulting, research, and instruction of a few students. In 1872, he was named professor of analytical and applied chemistry and mineralogy at the University of Pennsylvania, on condition that he be allowed to maintain his private practice. Certainly not resting on any laurels, he busied himself as well as chemist and mineralogist for the Second Geological Survey of Pennsylvania in 1874, and as the chemist for the Board of Agriculture of Pennsylvania from 1877 to 1888. In that year, he reopened his commercial analytical laboratory yet once again, and managed it until his death five years later on 2 February, 1893. In 1880, he served as fifth President of the American Chemical Society (Figure 3).

Figure 3. Frederick Augustus Genth, Fifth President of the American Chemical Society. Courtesy: American Chemical Society. 6

Adept as he was at multi-tasking, Genth was remarkable for the substantial contributions he also made to fundamental science, particularly in the area of complex cobalt-amine compounds, tellurides, phosphates, fertilizers, and rare minerals. He was passionate about the latter, having personally collected and catalogued about 5,000 specimens. Among these he identified and described about 20 new minerals, including calaverite, maconite, and penfieldite (19).

1862: The Element of Charles Frederick Chandler (1836-1925) Born on December 6, 1836 into a Massachusetts family of modest means, Charles Chandler distinguished himself early on as a hard worker who would take on extra duties and manufacture the time to do them by getting up early and going to bed late. His first interest in chemistry was awakened in high school by an inspiring and enthusiastic instructor who left a deep and lasting imprint on his eager student – chemistry soon became Chandler’s dominant lifelong interest. His work habits enabled him to earn enough money to outfit a small chemistry laboratory in the attic of the family home, and due to the many hours he spent there, he developed a proficiency unusual for a person so young. In 1853, he began to attend the Lawrence Scientific School of Harvard College, but soon found that a chemistry curriculum was virtually non-existent in the United States. Advised to study in Germany, he was helped along by a new acquaintance, Charles Arad Joy (1823-91), professor of chemistry at Union College, Schenectady, New York, who provided him with a letter of introduction to Friedrich Wöhler (1800-82), in Göttingen, who welcomed him immediately into his laboratory in 1854. Chandler, not only hard-working, but precocious, eventually earned the M.A. and Ph.D. degrees in 1856, when he was barely twenty years old! In the following year, he became assistant professor of chemistry at Union (although janitorial duties accompanied this appointment because that is where the budget line was), and succeeded to the chair of the department a few years later when Professor Joy moved on to Columbia College in the City of New York. At the time of his presumed discovery, Chandler was not yet 26 years old. In the preceding year, he had analyzed a mineral found in the Rogue River in Oregon, using the tried and true methodologies that he had perfected while doing the research for his doctoral dissertation under the guidance of Professor Heinrich Rose (1795-1864) in Berlin. Chandler was confident that he had come upon a new metal but felt that he needed additional sample in order to confirm it. Waiting in vain for a year, he finally decided to publish his results, done in triplicate according “to the ordinary routine of qualitative analysis,” with the same outcome. However, in the last paragraph of his paper (20), Chandler notes that a colleague apprised him of work done by Friedrich Genth ten years previously. The metal observed by Dr. Genth occurred among grains of platinum from California. It was malleable; it fused readily on charcoal before the blowpipe, becoming covered with a coating of black oxyd (sic); it dissolved in borax to a colorless bead, which became opalescent on cooling; it was dissolved by hot hydrochloric acid and by nitric acid; 7

and its solution gave a brown precipitate with hydrosulphuric acid. It seems quite probable, therefore, that the metal which I have observed in the Rogue River platinum, is identical with that observed by Dr. Genth.

Realizing that he had duplicated Genth’s work, Chandler wisely let the matter drop, but not before actually publishing it (21). Through a series of happy circumstances, not least of which was Chandler’s analytical expertise, he too received an invitation to Columbia, this time to instruct in the new School of Mines being formed at that institution. And there he remained for the following 46 years until his retirement in 1910; he acted as Dean of the school for 33 of those years (Figure 4).

Figure 4. Portrait of Charles Frederick Chandler. Michael De Santis, 1935, oil on canvas; Gift of Francis P. Garvan. Courtesy of the Union College Permanent Collection. 8

In the minutes of the Board of Trustees, following a recital of his many achievements, this entry can be found (22): Professor Chandler will carry with him into his retirement the affectionate regard and esteem of two generations of students as well as a host of colleagues on the teaching staff of the University. The Trustees record their grateful appreciation of this long and generous career of devoted service. Furthermore, Columbia’s President, Nicholas Murray Butler (1862-1947) refers, in his annual report of November 7, 1910, to Chandler’s wide-ranging activities during his tenure ((21), p. 186): To his teaching power as well as to his effective and conscientious service as administrator, the Department of Chemistry and the School of Mines, to which it primarily belonged, owed almost everything for many years. Professor Chandler has long been a point of contact between the University and the public, between science and industry and the public health. His career is unique of its kind, and we shall not soon look upon his like again. As Butler observed, Chandler’s interests were many and varied and his expertise and service ethic left a profound impact on his adopted city as well as the state of New York and the nation in general. Here are some examples:

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Chandler’s contacts with industry led to the accumulation of many chemical products that soon outgrew the space in his office; the initial collection served as the nucleus of Columbia’s famous Museum of Chemistry (23). Chandler opened his museum in order to show his students the things he talked about in his many lectures. A placard on display at the entrance to the chemistry department continues: “He began to collect material for the museum almost immediately on his arrival at Columbia as the Civil War was ending. For half a century after, he bought rare and interesting exhibits of chemicals and products of various chemical industries. Many times he bought items from exhibits out of his own pocket and materials were donated by the chemical industries...Professor Chandler took great delight in keeping up the museum.” Among the many exhibits are collections of dyes, rare earth elements (Figure 5), minerals, and laboratory equipment. When the New York College of Pharmacy was struggling as a “one-room schoolhouse,” Chandler offered his services and lectured and demonstrated (at his own expense) for three evenings per week year after year until the college could grow and expand, eventually becoming the College of Pharmacy of Columbia University In 1872, Chandler was appointed adjunct Professor of Chemistry and Medical Jurisprudence at the fledgling New York College of Physicians 9

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and Surgeons; for over 20 years, he lectured there every day between 5 and 6 PM, always raising his voice in favor of a more thorough scientific training for those entering the medical profession In 1870, together with his brother, William Henry Chandler (1841-1906), of Lehigh University, he founded the American Chemist and continued its publication until it was superseded by the Journal of the American Chemical Society in 1877 Chandler was a founding member of the American Chemical Society and twice served as its President in 1881 and 1889 Chandler initiated and oversaw the construction of Havemeyer Hall which opened its doors to students and faculty in 1898. A state-of-the-art chemistry department at the turn of the 20th century, this site was home to many famous scientists and notable scientific achievements. A century later, in 1998, it was designated a National Historic Chemical Landmark by the American Chemical Society (24) Chandler was recognized as the highest authority in his day in this country in industrial chemistry, serving as consultant in areas like sugar refining, petroleum refining, photomechanical processes, and calico printing; he invented and introduced the system of assay weights used to this day by assayers and metallurgists He was a pioneer in the area of water quality and sanitation, publishing landmark papers on these subjects Chandler devoted himself assiduously to all branches of hygiene and sanitary science, studying with the utmost care and thoroughness all factors bearing upon the health of a great city; for eleven years, he served as the President of the New York City Board of Health He served as one of the scientific directors of the New York Botanical Garden, and was for many years chemist for the Croton Aqueduct Commission For several years, Chandler was president of the New York State Charities Aid Association and took an active part in securing proper state care for the indigent insane His chemical expertise was called upon in the work of commissions under four U.S. Presidents: Chester A. Arthur (1829-86),Grover Cleveland (1837-1908), William McKinley (1843-1901), and Theodore Roosevelt (1858-1919). Chandler was the first scientist to receive an honorary D. Sc. degree from Oxford University (1900)

After a brief illness, on August 25, 1925, Charles Frederick Chandler died at his home in New York City at the age of 89. The young chemist who, in 1862, had the brashness to announce the false discovery of an element similar to platinum had become a national icon, celebrated both at home and abroad. He was mourned deeply by all who knew him or were familiar with his remarkable contributions to the well-being of humankind (25). 10

Chandler’s papers are now held in the Columbia University Archives (26) along with a finders’ guide. The materials are held off-site and require a 24-hour notice for access; it is a treasure trove for anyone planning a full-fledged biography of this extraordinary gentleman.

Figure 5. A Display of Rare Earth Element Compounds from the Chandler Museum Collection, Havemeyer Hall, Columbia University. Most of the Elements are Displayed as their Oxides. Photograph: Mary Virginia Orna.

1869 – The Advent of the Periodic Table 1869 seems like a very late date for Dmitri Mendeleev’s proposal of an orderly arrangement of elements to appear given the fact that phenomenological relationships among the elements had accumulated in the previous decades. From the 1820s on, scientists like Johann Wolfgang Döbereiner (1780-1849), Leopold Gmelin (1788-1853), Oliver Gibbs (1822-1908), and William Odling 11

(1829-1921) “played” with the idea of groups of elements with similar properties. Then in 1860, Stanislao Cannizzaro (1826-1910) provided what many historians of science consider the catalytic idea that precipitated the simultaneous proposals for what we now call the periodic table (27). In that year, many of the leading scientists of the day met at Karlsruhe, Germany, to discuss vexing questions such as the uncertainty of many chemical formulas of even the simplest of compounds. Cannizzaro, influenced by the theory of chemical bonding derived from Amedeo Avogadro’s (1776-1856) hypothesis, emphasized the importance of atomic weights. He distributed a paper based on a course he taught at the University of Genoa in which he stated: Once my students have become familiar with the importance of the numbers (i.e. atomic weights)…, it is easy to lead them to discover the law which results from their comparison. ‘Compare,’ I say to them, ‘the various quantities of the same element contained in the molecule of the free substance and in those of all its different compounds, and you will not be able to escape the following law: The different quantities of the same element contained in different molecules are all whole multiples of one and the same quantity, which, always being entire, has the right to be called an atom (28).’” Cannizzaro illustrated his statement with many examples taken from his own experimental data, and it is a known fact that many of the attendees of the congress were profoundly influenced by him, including Julius Lothar Meyer (1830-95) and Dmitri Ivanovitch Mendeleev (1834-1907). Nevertheless, the pathway from atomic weights to the periodic table was anything but direct, which is why it took almost ten years for both Mendeleev and Meyer to publish their ideas, which had percolated in their minds for a long time. I. S. Dmitriev, Director of the Mendeleev Institute at the University of Saint Petersburg, writes (29):

Mendeleev’s discovery of the Periodic Law did not follow a linear pathway, but rather one that was complicated, difficult, winding, one that utilized various criteria over a period of time.

It is beyond the scope of this chapter to give a detailed narrative of the periodic table’s development; many more comprehensive works have done a fine job in this regard (30, 31). Once chemists realized that not only could the periodic system bring order out of chaos, had predictive possibilities, was a guiding light in the search for new, and now “known to be missing” elements, but also that it served as a theoretical tool, as a map of the way in which electrons arrange themselves in atoms, and even the history of the development of life itself (32), it quickly took its rightful place as the “chemist’s Bible.” It has gone through many revisions since it was first visualized by Mendeleev and Meyer (Figure 6).

12

Figure 6. A monumental image of Dmitrii Mendeleev surrounded by “his” elements arranged like a sundial. Plaza facing the Faculty of Chemical and Food Technology Building of the Slovak University of Technology in Bratislava, Slovakia. Photograph: Mary Virginia Orna.

1877: J. Lawrence Smith (1818-1883) and His Discoveries Although J. Lawrence Smith served as the second President of the ACS in 1877, he was a Johnny-come-lately to the Mendeleev bandwagon: in the very year of his presidency, he had the misfortune of claiming the discovery of a new element that embroiled him in unlooked-for controversy. Although Mendeleev’s periodic table could predict missing elements in what we now call the main body of the table, the case of the rare earth elements was a different matter. They seemed to fit nowhere but had to fit somewhere in the table, but no one could fathom just how many there were, and if there was any limit to their number. Coupled with the fact that rare earth elements’ chemical similarities were so close that they were difficult to separate, chemists who worked with them often ended up with mixtures that they believed were pure substances. Smith, at this stage in his career, occupied himself with a serious study of the mineral samarskite, a rare earth-bearing mineral first extracted from the southern Ural mountains of Russia and named after the Russian mining official, Vasili Samarsky-Bykhovets (1803-70). Unknown to Smith, another chemist with formidable analytical skills and vast experience in the separation of rare earths, Marc Abraham Delafontaine (1837?-1911?), in Chicago, was devoting his time 13

to examining the same mineral. It is not surprising that virtually simultaneously (although Delafontaine had a slight edge), both announced the discovery of a new element. Delafontaine called his philippium in honor of his benefactor, M. Philippe Plantamour of Geneva; Smith called his mosandrium (or mosandrum) after the great rare earth chemist, Carl Gustav Mosander (1797-1858) (33). An immediate priority dispute erupted which was played out in the pages of Comptes rendus chimie (34). Delafontaine claimed that the presumed mosandrium was actually composed of about 75-80% terbium and a 20-25% mixture of yttrium, erbium, didymium (later resolved into its two constituent elements, praseodymium and neodymium, by Carl Auer von Welsbach (1858-1929) in 1903) (35) and “his” philippium. History has been kind to Delafontaine. The element discovered and isolated by Swedish chemist, Per Theodore Cleve (1840-1905), in 1878 and called holmium after the city of Stockholm, was later recognized to be identical to philippium, so today both Cleve and Delafontaine share credit for the discovery, although the latter lost the glory and privilege of conferring the name. Meanwhile, Smith’s claim fell into oblivion. Undaunted, Smith continued his assiduous study of samarskite. He announced the discovery of two new elements, columbium (36), presumably from its source in the mineral columbite and rogerium (in honor of his natural philosophy professor at the University of Virginia, William Barton Rogers (1804-82), thus being the first person to name an element after a person yet living. These claims were noted at a meeting of the National Academy of Sciences that took place on October 28-30, 1879 (37): Prof. J. Lawrence Smith gave an informal account of some recent researches for new elements. Some years ago, he found a field of research in the cerium and yttrium minerals, and was well satisfied that he had obtained a new substance, which he named mosandrum, in the cerium group. Since then he has been studying the compounds of samarskite and has found, he believes, two new elements, one of which he calls columbium, and the other he proposes to name in honour of his friend and the instructor of his youth, Prof. William B. Rogers. But having much other business requiring his attention, Prof. Smith has done little in that line of research, since then, except to purify some mosandrum. So, since neither of these discoveries seemed to be credible, no further work was done on samarskite except for a privately published piece in 1881 (38) that appeared two years later in the American Journal of Science (see reference (33)). Smith continued to publish almost up until the date of his death in 1883, but most of his interests had turned to the analysis of meteorites. In his Bakerian Lecture of 1883, William Crookes (1832-1919) remarks in a footnote (39): Dr. J. Lawrence Smith in a paper read before the United States National Academy of Sciences in 1879, announced the discovery in Samarskite of two new elements, which he named Columbium and Rogerium (Wyckoff, W. C. “U.S. National Academy.” Nature 1879, 21 (11 December), 14346). I have failed to find any further notice of these elements. This 14

Columbium should not be confounded with the well-known Columbium sometimes called Tantalum. With this remark, Crookes seems to have dealt the death blow to columbium and rogerium in the very year of Smith’s own death. Aside from Smith’s penchant for discovering erroneous elements, he was internationally well-recognized for his chemical expertise. A precocious child bordering on genius, he was proficient in mathematics at the age of four. Following two years of study at the University of Virginia, he took his medical degree at the University of South Carolina, submitting a thesis entitled “The Compound Nature of Nitrogen.” After receiving his medical degree, he continued his education in Paris, studying organic synthesis with the likes of Jean-Baptiste Dumas (1800-84) and medicinal chemistry with Mathieu J. B. Orfila (1787-1853). However, it was his chance visit to the laboratory of Justus von Liebig at Giessen in 1841 that turned the entire course of his life and landed him firmly in the field of chemistry. Thenceforward, dividing his time between Paris and Giessen, he did elegant chemical work that earned the respect of his mentor, Liebig. For example, in 1842, he published a landmark paper on the composition of spermaceti, possibly one of the earliest forays into organic chemistry by an American chemist (40) that was reproduced in several prestigious European journals. Ever peripatetic, Smith found employment in the Ottoman Empire for several years, conducting research on the natural resources of that area with a view to their commercial exploitation. Back in the United States, he spent several years lecturing in New Orleans, and later on in Virginia and in Washington, D.C. Then in 1854, he received an appointment as Professor of medical chemistry and toxicology at the University of Louisville (KY), in which post he remained for twelve years. Then in 1866, he resigned from formal teaching, finding it a restrictive burden on his many projects that required extensive travel. He established, instead, a home laboratory where he conducted private research in analytical and mineralogical chemistry, and in toxicology. He developed the potassium chromate test for barium and the J. Lawrence Smith method of analysis of silicates for alkali metals in 1853. Both of these tests remained the best analytical tools available for approximately a hundred years. In 1872, Smith was elected a member of the National Academy of Sciences, and in 1879, he became corresponding member of the Académie des Sciences, Institut de France, succeeding Sir Charles Lyell in this capacity; he was the first American to hold this post. He received many honors both at home and abroad. He was a Chevalier of the Légion d’Honneur, President of the American Association for the Advancement of Science in 1874, and the second President of the American Chemical Society in 1877. Early in his career, Smith began to collect meteorites, and a large body of his published work, especially during his later years, consists in their analyses. One of the final acts of his career was the transfer of this collection, by that time monumental, to Harvard University, by way of purchase (Figure 7). The last paper that Smith published was in June 1883 (41). In it, he promised further research, “my health” permitting. It did not. On October 12, 1883, J. Lawrence Smith succumbed to the ravages of chronic liver disease (42). 15

Figure 7. Meteorite Display at the Harvard University Museum of Natural History Earth and Planetary Sciences Gallery. © President & Fellows, Harvard College, courtesy Harvard Museum of Natural History. Photograph: Jess T. Dugan. In an adjacent display case, a caption reads, “The Harvard collection includes significant specimens from the personal collections of J. Lawrence Smith, a famous 19th century chemist.” The Harvard purchase of Smith’s meteorite collection was funded by subscription to which a number of prominent Bostonians contributed (43). Smith’s widow, Sarah Julia Smith, donated the funds in “deed of trust” (44) to the National Academy of Sciences to promote research on meteorites, and the monies are still used to this day for that purpose. In addition, Mrs. Smith also endowed the J. Lawrence Smith Medal (Figure 8) and Prize in memory of her husband. The National Academy has awarded it every three years since 1888 for recent, original, and meritorious investigations of meteoric bodies. 1901: Charles Baskerville and the “Discovery” of Carolinium At the dawn of the 20th century, one great organizing chemical principle was in place: the periodic behavior of the elements, having been enunciated some thirty years previously by Dmitri Mendeleev. It was a necessary, but unfortunately, not sufficient, breakthrough in the business of element discovery. In what we now call the main groups of the elements, it was clear that some elements still remained undiscovered in Mendeleev’s time, and the following decades witnessed their identification. At the same time, where to place the rare earths, presently called the lanthanides, and how many might actually exist, was still a mystery. How much more of a mystery, then, was the placement of what would later be called the actinides, a problem that Glenn Seaborg (1912-99) took up in 1944 while 16

he and his group were struggling with the placement of the transuranium elements in the periodic table (45).

Figure 8. The J. Lawrence Smith Medal. Courtesy: National Academy of Sciences. Image used with permission.

In 1901, a distinguished chemistry professor at the University of North Carolina, Charles Baskerville (1870-1922), was having a similar struggle with one of the members of this group, thorium. He did not have the advantage of two additional breakthroughs that would come later and clarify many things, i.e., the concept of atomic number and the concept of isotopes. Furthermore, the practice of fractional crystallization was the only means available at the time for substance purification, a fact that rendered so-called “very pure” samples almost moot. Charles Baskerville was an inspiring teacher and occupied an esteemed place in chemical education for almost thirty years. His lecture style was characterized by remarkable clarity and lucid expositions drawn from his highly organized store of notes and references. He was noted for his research on the rare earths and on the chemistry of anesthetics. On the industrial side, he studied pulp and paper recycling, and the refining and hydrogenation of vegetable oils (46). His early university education was at the University of Mississippi, after which he attended specialty courses at the University of Virginia and Vanderbilt University, finally 17

completing his B.S. degree in chemistry at the University of North Carolina in 1892. Following an intensive summer course in chemistry at the University of Berlin (1893), he returned to North Carolina to complete his doctoral work under F. P. Venable (1856-1934) in 1894. Remaining at the university following the completion of his degree work, he advanced rapidly through the professorial ranks, becoming chair of the chemistry department upon Venable’s accession to the presidency of the university. In 1903, Baskerville claimed through two papers published in the Journal of the American Chemical Society that he had discovered two new elements embedded in samples of thorium compounds extracted from monazite sands, carolinium (he had suggested the presence of carolinium as early as 1901) and berzelium. Though these discoveries were later refuted, Baskerville received enough acclaim from his professional colleagues to be invited to occupy a faculty position at the City College (later, City University) of New York, which he accepted in 1904. In reading Baskerville’s two papers claiming the discovery of carolinium, one cannot help but notice the “stream of consciousness” style that seems to have characterized his work. Far from dividing his papers into the sections customary today – experimental, results, discussion – we are treated to a running commentary full of details such as (47): …when the extraction had continued about six hours, it was discovered that an overlooked small defect of the cork had permitted the gradual introduction of about a drop of water, which came from the sweating of the condenser overhead. As this vitiated the experiment…the experiment must be discarded… Acknowledging that the chemistry of thorium was very complex, Baskerville attempted to prepare pure samples of compounds of thorium such as the oxide and the tetrachloride. In doing so, he took measurement of specific gravity as an index of the purity of the compound. In his determinations, he encountered the anomalous presence of an oxide with an unusually high specific gravity which, he claimed, “cannot be accounted for except by the presence of either a new oxide of a known element having greater density than the usual non-volatile residue after ignition, or an unknown element” ((46), p. 767). Later on in the same paper, Baskerville asserts that his preparations contained “a constant unknown impurity in practically all the materials used. This constituent must be an element of much higher atomic weight….between 260 and 280. On account of the extensive occurrence, in this state (North Carolina), of the monazite sands from which the original material was obtained, if the investigation give a successful issue, I should like to have the element known as Carolinium, with the symbol Cn.” ((46), p. 773) Then, Baskerville immediately says that this is a preliminary paper only….so, stay tuned for more information in a subsequent publication. However, Baskerville, in that same year, presented a framed set of “carolinium” samples to the Columbia University Chandler Museum, where it remains on display to this day (Figure 9). Since we know that carolinium does not exist, what the vials contain is also a mystery to this day. 18

Figure 9. Samples of Carolinium Compounds (Oxide, Sulphate, and Acetyl-actonate [sic]) on Display at the Chandler Museum Collection, Havemeyer Hall, Columbia University. Photograph: Mary Virginia Orna.

The Museum Caption reads: “In 1901, Dr. Charles Baskerville of the City College of New York announced ‘the Existence of a New Element Associated with Thorium.’ This potential new element, heavier and more radioactive than thorium, he named ‘Carolinium’ (Cn) for the North Carolina location of the monazite sand from which it was obtained.” The promised publication arrived in 1904 (48). In it, Baskerville reviews the history of thorium research and how it was previously misidentified as other new elements such as donarium and wasium, as well as the presumed thorium α and thorium β of Bohuslav Brauner (1855-1935) with atomic weights respectively of 220 and 260-280. The upshot of Baskerville’s conclusions is that thorium is not a primary radioactive body, that it is complex and not a chemical element, and that it can be resolved into at least three other distinct bodies sufficiently identified to deserve distinct names (including his own carolinium and berzelium), and that he awaits spark spectral confirmation of these conclusions ((47), p. 941). And a note appended to these observations by William Crookes, the doyen of state-of-the-art spectroscopy at that time, states that “Making allowance for the fact that my spectrum is from the metal, while that from your material is from the 19

chlorides solutions, all five spectra are practically identical, all the prominent lines being seen in each spectrum, while there are no lines in one which are not seen in the others…” ((47), p. 942) Crookes softens the blow by asserting that these are preliminary results and by no means prove that such bodies (carolinium and berzelium) do not exist. However, since we hear no more of these two presumed elements in the literature, we assume that Baskerville (Figure 10), reproduced from the Chemical Heritage Foundation Archives (49), provided no new evidence regarding their existence.

Figure 10. Charles Baskerville. Reproduced from reference (49), with permission. So, where did Baskerville take a wrong turn? His experimental work seemed to be very detailed, but closer examination reveals that it was incomplete and more than likely not reproducible. True, he determined the atomic weights of his presumed new elements (255.6 and 212.0 respectively). He determined that the properties of their oxides were not the same as those of the parent material, thorium, from which they were extracted. He confirmed the phosphorescence and radioactivity of the “new” substances, and when he realized that the arc and spark spectra were identical, he then supposed that the material examined was 20

not sufficiently pure and that the spectral data were not complete. However, it is very likely that he had bit off more than he could chew. Thorium was known to be difficult in every way. Its isolation and purification from other elements in the matrix were accomplished with difficulty only a decade after Baskerville had published his papers (50). Regarding the observed radioactivity of some of the fractions, one might presume that the presence of uranium in greater or smaller amounts as an impurity would account for the enhanced activity of the fraction attributed to carolinium. While Baskerville made note of the radioactivity, he did not attribute it to thorium since he even denied that thorium was radioactive. And it is certainly not possible that he had succeeded in isolating the 212Th isotope since he used only chemical means in his investigations, and one cannot separate isotopes of the same element that way since they exhibit identical chemical behavior. When all is said and done, it is quite possible that Baskerville failed simply because of the limitations inherent in his own research. Seventy-five years after Berzelius discovered thorium, Baskerville was using essentially the same laboratory techniques. Perhaps sensing that this line of research was a dead end, he abandoned his search for confirmation of the element that would have been the first to be named after a state of the United States for much more practical work in organic chemistry. During his relatively brief chemical career, Baskerville managed to author six books and over two hundred technical papers. He was an active member of several scientific societies, having served in official capacities in some of them and was even elected a Fellow of the American Association of the Advancement of Science. Hardly 52 years of age, Baskerville succumbed to a bout of pneumonia on January 28, 1922 (51).

1913 and 1915: Atomic Number and Isotopes As technology advanced, many elements were discovered that confirmed Mendeleev’s initial predictions. Some bumps along the road were how to accommodate the plethora of rare-earth elements, the unexpected discovery of the noble gases, and of numerous radioactive species that seemed to be individual new elements until the existence of isotopes came to be understood. Antonius Johannes van den Broek (1870-1926) was a Dutch amateur physicist notable for being the first who realized that the number of an element in the periodic table (now called atomic number) corresponds to the charge of its atomic nucleus. This hypothesis was published in two papers in 1911 (52, 53), just one month after Ernest Rutherford (1871-1937) published the results of his experiments that showed the existence of a small charged nucleus in an atom, and inspired the experimental work of Henry G. J. Moseley (1887-1915), who found good experimental evidence for it by 1913. Moseley foresaw that his X-ray method would “prove a powerful method of chemical analysis…It may even lead to the discovery of missing elements, as it will be possible to predict the position of their characteristic lines” (54). 21

Thus the atomic number not only placed a limit on the number of elements that could be discovered, but it also served as the key to atomic structure that identified each element – the defining property of an element could no longer be ascribed to its atomic weight. Following upon the results of this landmark paper, chemists realized that only seven of the naturally occurring elements remained to be discovered, thus cutting down drastically the number of reported false discoveries and setting in motion an element hunt full of controversial competing claims that lasted for decades (55). Barely two years following Moseley’s discovery of what came to be called atomic number, Frederick Soddy (1877-1956) and his research assistant, Ada Hitchins (1891-1972), working at Glasgow and later at Aberdeen, demonstrated that the density of elemental lead from various sources differed substantially, whereas the atomic volume remained constant. The only conclusion to draw from this observation was that the atomic weights of lead could vary, leading directly to the concept of isotopes, for which Soddy received the Nobel Prize in chemistry in 1921 (56).

1926: B Smith Hopkins (1873-1952) and Illinium In the decade that followed, chemists used these two new tools so cleverly that by 1925, only one rare earth element remained unidentified: the recalcitrant element 61. Its existence had been surmised by Bohuslav Brauner in Prague as early as 1902 (57): Apart from the 10 elements already listed…and more or less accurately studied by me, about seven to ten additional elements could be placed in this group…It is not impossible that one would be able to split neodymium, Nd = 143.8, into at least one element with a smaller atomic weight, and into another element with a higher atomic weight of about 145 and, similarly, some more gaps lying in the area between Ce and Ta could be filled.

When it was later shown that element 61 did indeed exist, Brauner claimed credit for the discovery in a letter to Nature (58):

I arrived at the conviction that the gap between the neodymium and samarium was abnormally large. In my paper…read in St. Petersburg in 1902, I came to the conclusion, not reached by any chemist before – that the following seven elements, possessing now the atomic numbers 43, 61, 72, 75, 85, 87, and 89, remained to be discovered. As regards element No. 61, the difference between atomic weights of Sm-Nd = 6.1, and it is greater than that between any other two neighboring elements. 22

Later work by Moseley showed definitively that an element should exist between neodymium and samarium (59). Researchers took up the challenge, and many were in hot pursuit, assuming that this element lay hidden in very small amounts among the other rare earths. Consequently, they continued to use the method that had led to success in so many other cases: fractional crystallization. In 1924, at the University of Florence, a newly-minted chemist, Lorenzo Fernandes (1902-77), urged on by his mentor and autocratic head of the chemistry department, Luigi Rolla (1882-1960), developed an expertise in the separation and purification of rare earths, as well as the X-ray check on the purity of the samples. The department had plenty of space, eager workers, and later on, enormous amounts of raw material to work with. Rolla realized that there was a gap between neodymium (Z = 60) and samarium (Z = 62), and who could better fill it than he and his department? Little did he realize that more than 56,000 fractional crystallizations later, he would still be uncertain about what he had (or did not have) in his hands. Thinking that he and Fernandes had observed the characteristic X-ray spectrum of 61, yet not wishing to make a premature announcement to claim priority, Rolla sent a sealed packet containing his sample and a note about his claim to the Accademia dei Lincei (the Italian equivalent of the National Academy of Sciences) – just in case someone else came along to make the claim while he was trying to obtain a larger and purer sample (60). Rolla had good reasons for his fears. Unbeknownst to him, a team headed up by B Smith Hopkins (Figure 11) at the University of Illinois in the United States had begun their search for this missing element in 1923. They were in “friendly competition” with Charles James at the University of New Hampshire who, prodded by William Ramsay (1852-1916) (61), took up the search more than a decade earlier in 1912. Hopkins, cognizant of the possibility that element 61 might be the rarest of the rare earths and virtually undetectable because of the difficulty of separating enough of it for X-ray analysis, nevertheless soldiered on. A caveat expressed early on indicated that he realized that since no new absorption bands had ever been observed in the intermediate fractions when the double magnesium nitrates of the rare earths had been subjected to fractional crystallization, element 61 might be concentrating with neodymium whose extensive absorption bands could succeed in masking any other bands present in the regions being examined (62). The team then followed the first paper with a companion piece (63) in which they claimed discovery of the new element, which they named illinium (after the university and state of Illinois (64)) based on the presence of 130 arc lines in the red and infrared spectrum and five lines toward the violet, corresponding closely to the theoretical positions for the Lα1 and Lβ1 for element 61 (65). Published virtually simultaneously was an article by James M. Cork (1894-1957), Charles James (1880-1928) and Heman C. Fogg (1895-1952) (66) which contained the L series spectroscopic values that came uncannily close to those published two decades later after the element had been isolated.

23

Figure 11. Professor B Smith Hopkins, 1937. Photo courtesy of the University of Illinois at Urbana-Champaign Archives, Faculty, Staff and Student Portraits, Record Series 39/2/26, Box 32, “Hopkins, B Smith.”

At this point, Luigi Rolla, throwing all caution to the winds, immediately revealed the contents of his sealed packet in the pages of the Gazzetta Chimica Italiana in three parts: (1) A two-page review of the results of his search for element 61 (67), (2) A much longer piece giving experimental details on the extraction and further procedures on the samples in hand (68), (3) A confirmation of observed spectra of element 61 by an independent laboratory (69). Predictably, the Illinois announcement coupled with Rolla’s speedy action precipitated a priority dispute between the Florence group and the chair of the Illinois chemistry department, the renowned W. A. Noyes (1857-1941), that played out in the pages of Nature (70). Noyes starts out by citing the history of the search for element 61 at the University of Illinois in 1919, in partnership with the U.S. Bureau of Standards, resulting in three publications over three successive years, 1921, 1922, and 1923. Noyes notes that the second publication appeared at just about the time that Rolla began his work, and two years before Rolla deposited his sealed packet with the Academia dei Lincei. Reviewing additional work at Illinois, Noyes concludes (71): 24

In the light of these facts it would seem that the honour for the discovery of No. 61 belongs primarily to Prof. Hopkins, and that the element should be called Illinium rather than Florentium. This does not detract from the credit which Prof. Rolla should receive for his independent discovery of the element. Both Prof. Rolla and Prof. Hopkins realise that a large amount of additional work must be done before the element can be fully accepted.

Rolla’s response involved reviewing the history of the reports on the spectra of the supposed element 61 by Carl Clarence Kiess (1887-1967) and Leonard Francis Yntema (1892-1976), quoting the latter (72): X-ray analysis of samples from different sources has so far given no evidence of the presence of this element. Rolla goes on to say that it is not sufficient evidence for assuming the existence of a new element discovered only from the fact of having seen new spectral lines. And then he delivers his punch line: We obtained…the first photographs of K-absorption spectra showing the characteristic band of element 61…the first certain data…[and] we believe that we should be credited with priority for the discovery. The spectroscopic analysis had been done by Prof. Rita Brunetti (1890-1942) of the University of Florence Physics Department in Arcetri using state-of-the-art instrumentation. Noyes’s immediate response gave short shrift to the Florentine claim (73): The fact that Prof. Rolla deposited a [packet] instead of publishing his paper, demonstrates that he was not, at that time, sure of his discovery…Harris, Hopkins, and Yntema…were sure of their results on the basis of four independent lines of evidence… Realizing that endless correspondence in the pages of the literature would not solve the impasse, Rolla, seeking confirmation of his discovery, actually traveled to Illinois to see with his own eyes what progress was being made in the isolation of element 61. Finding none, and with his mind at rest, he stopped off in Copenhagen on his way home and presented a sample of his material to Niels Bohr for confirmation. Subjecting Rolla’s enriched florentium sample to a much more scrupulous and accurate spectroscopic examination than had ever been done before, Bohr’s analysis came up empty: element 61 did not exist in Rolla’s material. It was only fourteen years later, in 1941, that Rolla sent a letter of retraction, partly written in Latin and couched in an extensive description of his work on neodymium and samarium, to the little-known Vatican journal of the Pontificia Academia Scientiarum (74). And while one might deem this no retraction at all, it 25

was far more than B Smith Hopkins ever did. To his dying day, Hopkins traveled extensively in the United States trying vainly to salvage illinium from the ashheap. His initial announcement of discovery was welcomed everywhere, and many textbooks had incorporated illinium, with the symbol Il, into their endflap periodic tables (75, 76). In Ref. (73) one can see the image of the periodic table published by Hopkins himself and labeled “The Nearly Completed Periodic Table of the Elements, 1926.” One can note the symbol, Il, for element 61, and the uncertain atomic weight, 146? As well as the blank spots for elements 85 and 87, and the symbol Ma for masurium, element 43, which was falsely claimed by Walter and Ida Tacke Noddack. At the University of Illinois, a large periodic table was painted onto the wall of the large lecture hall in the Chemistry Department; it was allowed to remain there for several years after Hopkins’s death in 1952. Illinium was allowed to remain in the periodic chart and the descriptive chemistry section of Hopkins, B. S. and Bailar, J. C., General Chemistry for Colleges, 5th Ed. (D. C. Heath & Co.: Boston, 1932) until 1956 in deference to the senior author (77). The genuine element 61 made its debut in 1947, two years after its actual discovery and isolation from uranium fission fragments by ion exchange chromatography at Oak Ridge National Laboratory in Tennessee (78). Originally called prometheum at the suggestion of Mrs. Grace Mary Coryell, wife of one of the discoverers, Charles D. Coryell (1912-71), its name was later changed by I.U.P.A.C. to reflect the suffixes of many of its companions in the periodic table: promethium. Prometheus was the Titan who stole fire from the gods and bestowed it upon humanity; in retaliation, Zeus condemned him to be chained to a rock where his liver was daily eaten by an eagle (presumably Zeus himself). Thus the name promethium is fraught with symbolism: the danger of having a powerful force in one’s hands and the frightful consequences thereof. Members of the American Chemical Society got their first look at compounds of promethium, yellow PmCl3 and pink Pm(NO3)3 at the June, 1948 ACS national meeting in Syracuse, New York. None of the would-be discoverers lived to find out that promethium actually did exist in nature. In 1956, Paul K. Kuroda (1917-2001) of the University of Arkansas hypothesized that the uranium ore coming from a pitchblende deposit in Oklo, Gabon (79), found to have a depleted level of U-235, could have undergone spontaneous nuclear fission. His hunch was later confirmed when the expected lighter elements, products of U-235 fission, were present in much higher quantities than normal. Kuroda subsequently organized a gigantic task force to extract naturally-occurring element 61 from pitchblende (80). The mass of the natural isotope is 147. The story of element 61’s discovery brings together over 100 years of chemical history, as noted by Clarence Murphy (b. 1934) (81): …no element has been “discovered” and named more times than 61…[it is intimately connected with the development of the understanding of atomic structure and of the Periodic Table…The story involves Roentgen’s discovery of X-rays and Moseley’s use of X-ray spectra to determine atomic numbers. It involves the more than one hundred-year effort to separate the rare earths and to find a place for them in the 26

Periodic Table. Finally it involves the development of ion-exchange chromatography and research on the atomic bomb during World War II… 1930: Professor Fred Allison (1882-1974), Virginium, and Alabamine When Bohuslav Brauner wrote his famous letter to Nature (58) regarding the atomic numbers of the elements he had divined that were yet to be discovered in 1902, he was witness to a frenzied hunt for the missing members of the periodic table that was already in progress, with two successes happening prior to his publication, i.e., in 1917 and 1923. Furthermore, Brauner was mistaken regarding his list of seven because one of these elements, actinium, had already been discovered in 1899 by André Debierne at the Institut du Radium, Paris. However, there were still seven elements if one substitutes element 91 for 89, as did Eric Scerri, correctly, in his book, A Tale of 7 Elements ((55), pp. 176-79). For clarity, please see Table 2. By 1930, only three elements remained on the Brauner-Scerri list (if we discount element 61, “discovered” in 1926 and destined to live on as illinium until at least 1956), and Fred Allison (1882-1974) was determined to reap the honors for discovering two of them, numbers 85 and 87. In 1927, Allison (Figure 12), then at the Alabama Polytechnic Institute, began to publish a series of papers on his use of the Faraday effect to identify the presence of dissolved metal ions in minute quantities. Essentially, by placing a salt solution in a glass cell through which a beam of polarized light can pass, the plane of polarization can be rotated by application of a magnetic field, and the rotation observed is accompanied by an increase in the brightness of the light. By adjusting the brightness to a minimum relative to a reference liquid such as carbon disulfide and changing the nature of the salt, it was possible to measure the time lag of the light’s appearance as a function of the identity of the metal ion in the solution. Allison found that each chemical substance, independent of the presence of other substances, produced its own characteristic minimum, or minima, of light intensity, persisting down to a concentration of about one part in 1011 (83). In a subsequent paper, Allison and his co-worker, Edgar J. Murphy, reported that their results showed that the positions of the minima observed were functions of the atomic equivalents of the metallic elements. Furthermore, they claimed that the number of minima was the same as the number of known isotopes of these metals. They claimed that the method had many advantages among which were rapidity, non-destructiveness, extraordinary sensitivity, and low cost. The one drawback noted was that operation of the apparatus required considerable skill and experience (84). In fact, one great advantage that Allison mentions almost in passing at the end of his article is that his method is so sensitive that it could identify the presence of the long-searched-for eka-cesium, element 87, in samples of pollucite and lepidolite ores. Not only that, he was able to observe six minima for this new element – now virginium, (in honor of the state of his birth) Va (later changed to Vi), in his lexicon – that he ascribed to VaCl, VaNO3, Va2SO4 and six stable isotopes of the new element (85). 27

Table 2. The Presumed Naturally Occurring Elements Remaining to Be Discovered in 1902 (82) Element No. Scerri

Element No. Brauner

Element Name

Date of Discovery

Discoverer(s)

Place of Discovery

28

43

43

Technetium (Tc)

1937

Carlo Perrier (1886-1948), Emilio Segrè (1905-89)

Palermo, Sicily

61

61

Prome-thium (Pm)

1945

Jacob A. Marinsky (1918-2005), Lawrence E. Glenden-in (1918-2008), Charles D. Coryell (1912-71)

Oak Ridge, TN, USA

72

72

Hafnium (Hf)

1923

Dirk Coster (1889-1950) and George de Hevesy (1885-1966)

Copenhagen

75

75

Rhenium

1925

Walter Noddack (1893-1960), Ida Tacke (1896-1978), Otto Berg (1874-1939)

Berlin

85

85

Astatine (At)

1940

Dale R. Corson (1914-2012), Kenneth R. MacKenzie (1912-2002), Emilio Segrè

Berkeley, CA, USA

87

87

Francium (Fr)

1939

Marguerite Perey (1909-75)

Paris

Element No. Brauner

Element No. Scerri

89 91

*

Element Name

Date of Discovery

Discoverer(s)

Place of Discovery

Actinium (Ac)

1899

André Debierne (1874-1949)

Paris

Protactinium (Pa)

1917

Lise Meitner (1878-1968), Otto Hahn (1879-1968), Kasimir Fajans (1887-1975)*

Berlin

Note on the discovery of protactinium. Kasimir Fajans and Ostwald H. Göhring (1889-1915?) discovered a short-lived isotope 234Pa (T1/2 = 6.69 h) in 1913. The longest-lived isotope, 231Pa (T1/2 = 32,500 y) was discovered in 1917 by three independently working teams: Lise Meitner and Otto Hahn working in Berlin; Kasimir Fajans working at Karlsruhe; Frederick Soddy, John A. Cranston (1891-1972) and Alexander Fleck, Baron Fleck of Saltcoats (1889-1968) working at Glasgow.

29

Figure 12. Fred Allison. Courtesy of the Auburn University Libraries Special Collections and Archives.

Figure 13 is a diagram of the magneto-optic apparatus (86). In a later publication with co-authors E. R. Bishop, A. L. Sommer, and J. H. Christensen, Allison opens the door to other possibilities for the minima observed, either to a cation with equivalent weight greater than that of Tl+, i.e., element 87, or to complex ions. After quickly eliminating the complex ions by experiment, the team concludes with an argument that seems like negative evidence (87): Since we have found nothing else which might give the minima attributed to compounds of 87 and since these minima persist after treatment with acids, bases, oxidizing and reducing agents, we conclude, as previously announced, that these minima are due to element 87. On a more positive note, they claim to have found 87 in sea water, California brine, Stassfurt kainite, crude cesium chloride, and monazite sand and samarskite from different sources. In all of these samples, the concentration of 87 was noted to be very low. 30

Figure 13. Magneto-Optic Diagram and Connections. Light from the spark source G passes through the optical train from L to E to measure differential time lags of the Faraday effect in liquids. Reproduced from reference (85). Copyright 1932. American Chemical Society.

In the same month that the Alabama team claimed the discovery of element 87, they also announced the detection of element 85 by the same magneto-optic method, and quickly named it alabamine (named after the state and the Polytechnic Institute where Allison worked) with the symbol Am. After treating a large amount (100 pounds) of monazite mineral, they found the appropriate minima that showed that alabamine went into solution as peralabamic acid, HAmO4. After preparing alabamine as lithium alabamide, the research team found that alabamides are readily oxidized to hypoalabamite, alabamite, alabamate, and peralabamate salts and their corresponding acids. They also found that the peralabamates were the most stable form of the element, and that its estimated atomic weight was 221. However, no experimental evidence aside from the well-described magneto-optic method, was given (88). Departing from his usual appeal to the physics and chemistry research community, in 1933 Allison published an article in the Journal of Chemical Education (89) in which he also laid claim to having discovered element 85 and 87, among many other things, and even a heavy isotope of hydrogen a year before deuterium’s discovery by Harold Urey (1893-1981) (90). But already toward the end of 1931, Allison’s reputation was beginning to fray. The 1931 Annual Survey of American Chemistry, published in 1932, ventured the following remark (91): 31

[Allison] brought forward additional evidence that the number of minima in the time lag of the Faraday effect in electrolytic solution is equal to the number of isotopes of the positive ion. With some exceptions, this relationship seems to hold and the concentrations at which the minima first appear are reported to be approximately inversely proportional to the relative abundances of the isotopes…thus two isotopes of hydrogen, two of thallium and four of barium were predicted and these numbers subsequently proven to exist. But also, one chromium, two strontium and two ruthenium isotopes were required by these observations and subsequently four chromium, three strontium and six or seven ruthenium isotopes were reported. If the method is as sensitive as is claimed, these isotopes should not have been overlooked. Questions about the validity of Allison’s method and claims began to multiply, principally because of his insistence that only those adept and experienced in using the method could achieve the correct results. Furthermore, attempts to reproduce Allison’s published results met with failure. By 1934, largely at the instigation of Nobel Laureate Irving Langmuir (1881-1957), the American Chemical Society began to ban the submission of any articles experimentally based on the magneto-optic effect. Perhaps the literature death-blow to the method came in the December, 1934 issue of the Physical Review where Herbert G. MacPherson (1911-93), at Berkeley, published the results of his year-long, exhaustive study of the magneto-optic effect, ending with the words (92): “…it is clear that this objective test does not reveal the existence of any real minima on the apparatus used. The author believes…that such minima as [Allison] saw from time to time can be explained as due to physiological or psychological factors.” Even so, the false American elements illinium, alabamine, and virginium did not go away very soon. They were still being listed as “rare metals” in the 1940 Materials Handbook (93) and as late as 1956 in the eighth edition of this same publication (94)! And, as we have noted previously, the end-flap periodic tables in major textbooks and wall charts published by scientific supply houses continued to retain these elements as late as 1956 (Figure 14). Far from being discredited, even by a very sarcastic and damning lecture on the subject, called “pathological science” by Irving Langmuir in 1953 (95), Allison seemed to move from strength to strength. He was chair of the physics department of Auburn University from 1922 to 1953, from whence he departed to become chair of the physics and mathematics at the University of Texas at Austin for a fouryear stint. He moved on to Huntingdon College to chair that institution’s physics and mathematics department for another five years, retiring in 1961, although he continued to teach there throughout the decade of the 1960s. And everywhere he went, he was greatly esteemed and admired by faculty colleagues and students alike, so much so that the buildings housing physics at these institutions are named after him (96). 32

Figure 14. A Periodic Table compiled by H. D. Hubbard and revised in 1941 for the W. M. Welch Manufacturing Co. of Chicago, Illinois. Note that, in addition to elements 85 and 87 (Ab and Vi) being present in this 1941 version, Il still occupies the space between Nd and Sm, and masurium (Ma) still claims to be element 43. Reproduced from reference (76). Copyright 1975. American Chemical Society.

Much later, Allison’s work was incorporated into a set of topics on “bad science” taught in the general chemistry and analytical chemistry at a liberal arts college. A number of cases, including Allison’s, were presented in a continuum of what the instructor termed bad, pseudo-, pathological, or deviant science. The author cites Allison’s magneto-optic method as something that started out as legitimate science and degenerated into bad science (97). In the 1993 Concise Columbia Encyclopaedia, Allison is still listed as the person who “announced” the existence of astatine (as alabamine), while the Berkeley research group of D. R. Corson, K. R. MacKenzie and E. Segrè are listed as its “producers (98).” However, in that same work, there is no question about who discovered element 87. Marguerite Perey is definitely given the credit she is due for discovering francium, element 87. However, the last sentence in the entry on p. 997 reads, “In the U.S., it was at one time called virginium.” So, despite their evanescent nature, both alabamine and virginium died hard.

Conclusion The three great advances that revealed the relationships and structures of the chemical elements were made over the space of time in which the “All-American” errors were made. None of them prevented the would-be discoverers from falling into error. Table 3 summarizes these erroneous claims. 33

Table 3. Summary of the “All-American” Elemental Errors “Chemical Element” (prior to the concept of Atomic Number)

“Physical Element” (Atomic Number)

Discoverer

Year

Symbol

Unnamed

--

H. A. Genth

1853

None

Unnamed

--

C. F. Chandler

1863

None

Mosandrium

--

J. L. Smith

1877

None

Rogerium

--

J. L. Smith

1879

None

Columbium

--

J. L. Smith

1879

Cb

Carolinium

--

C. Baskerville

1901

Cn

C. Baskerville

1904

Bz

Berzelium

Atomic No.

--

Illinium

B. S. Hopkins et al.

1926

Il

61

--

Alabamine

F. Allison

1930

Ab

85

--

Virginium

F. Allison

1930

Va, Vi, Vm

87

If we look a little closer at the fundamental reasons for these errors, we can discern quite a few: •

•

• •

•

claims based upon chemical tests and examination of chemical and physical properties of mixtures, not of pure samples (elements of Genth and Chandler); the difficulty of separating rare earth elements from one another due to their chemical similarity, plus lack of means to follow up on claims asserted (Smith); misconceptions about the nature of thorium earths, impure samples, incomplete spectral data, and poor experimental technique (Baskerville); excessive reliance on spectral lines of impure samples thought to be pure, and actually thinking to have seen something that was not actually there – wishful thinking – (Hopkins); use of an instrumental method that was entirely incapable of detecting less than nanogram quantities of material unless through delusion or fraud (Allison).

None of these would-be discoverers was ever actually ejected from the scientific community; in fact, some of them, like Fred Allison, went on to occupy high academic positions and retirement with honor – although he was the only one to receive censure from some sectors. 34

Examining all the cases in this paper, you will find a curious common denominator: not one of the authors of the alleged discoveries has ever issued a public retraction, thereby acknowledging and explaining the reasons and the causes of their (human) errors. The admission of one’s own error, when cultural and scientific conditions would require a corrected publication, would be a form of individual moral growth, but also a breakthrough for the scientific community, which at first was erroneously informed and therefore misguided. Unfortunately, in today’s overly competitive society, such an intellectual exercise would be counterproductive for a career. This may explain why some of these characters have ignored or forgotten their previous work. This could be the case, for example, for Baskerville who, after his dismal failures, changed his research interests drastically. Another exception can be made for Allison and Smith-Hopkins: they were acting in perfectly good faith; they believed blindly in their work and never wanted to consider the idea of having beaten the air instead of making a lasting contribution to scientific knowledge (29). A third consideration can be advanced as a justification of J. Lawrence Smith. He made his discoveries between 1877 and 1879, when his health was already shaky. Maybe he did not have time to write a proper retraction given the fact that he died not long afterwards. In any event, all of us can learn a great lesson from these errors: the history of science is never written in unwaveringly straight lines – there are always dead ends and backtracking along the way.

Acknowledgments The authors would like to acknowledge the contribution of Janan M. Hayes to the first portion of this paper on the role of American Chemical Society Presidents in reporting spurious elements. We refer to her paper, “Even ACS Presidents Announced the Discovery of New Elements, and They Were Wrong,” read at the ACS National Meeting, San Francisco, CA, August 11, 2014. The invaluable help of Vera Mainz (University of Illinois), Tammy Hartwell (Auburn University), and Eric Slater of the ACS Copyright Office in obtaining image permissions is deeply appreciated.

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64. There is some question about the naming attribution – Hopkins was a great fan of the Illinois soccer team, the “Illini,” and they perhaps should indeed receive some credit. 65. Siegbahn, M.; Friman, E. Vacuum spectrograph for high-frequency spectra investigations. Applications to the study of the rare earths. Physik. Z. 1916, 17, 176–78. 66. Cork, J. M.; James, C.; Fogg, H. C. The Concentration and identification of the element of atomic number 61. Proc. Natl. Acad. Sci. U.S.A. 1926, 12, 696–99. 67. Rolla, L.; Fernandes, L. Ricerche sopra l’elemento a numero atomico 61. Nota I. Gazz. Chim. It. 1926, 56, 435–36. 68. Rolla, L.; Fernandes, L. Ricerche sopra l’elemento a numero atomico 61. Nota II. Gazz. Chim. It. 1926, 56, 688–94. 69. Rolla, L.; Fernandes, L. Ricerche sopra l’elemento a numero atomico 61. Nota III. Gazz. Chim. It. 1926, 56, 862–64. 70. Noyes, W. A. Florentium or Illinium? Nature 1927, 120, 14. 71. Noyes, W. A. Illinium. Nature. 1927, 119, 319. 72. Rolla, L.; Fernandes, L. Florentium or Illinium? Nature. 1927, 119, 637–38. 73. Noyes, W. A. Florentium or Illinium? Nature 1927, 120, 14. 74. Gatterer, A.; Junkes, J.; Rolla, L.; Piccardi, G. Sugli spettri d’arco delle miscele neodimio-samarifere. Pontif. Acad. Sci., Comment. 1942, 6, 387–413. 75. The Internet Database of Periodic Tables. http://www.meta-synthesis.com/ webbook/35_pt/pt_database.php?PT_id=550 (accessed Sept 15, 2017). 76. Hopkins, B. S. Building blocks of the universe. Sci. Am. 1927, 136 (2), 87–89. 77. Trimble, R. F. What happened to alabamine, virginium and illinium? J. Chem. Educ. 1975, 52, 585. 78. Marinsky, J. A.; Glendenin, L. E.; Coryell, C. D. The chemical identification of isotopes of neodymium and of element 61. J. Am. Chem. Soc. 1947, 69, 2781–85. 79. Meshik, A. P. The workings of an ancient nuclear reactor. Sci. Am. [Online] 2005, 293, 83−91. http://www.scientificamerican.com/article/ancientnuclear-reactor/ (accessed Sept 15, 2017) 80. Attrep, M.; Kuroda, P. K. Promethium in pitchblende. J. Inorg. Nucl. Chem. 1968, 30, 699–703. 81. Murphy, C.; Charles James, B. Smith Hopkins and the tangled web of element 61. Bull. Hist. Chem. 2006, 31 (1), 9–18. 82. Data gleaned from Emsley, J. The Elements, 3rd ed.; Clarendon Press: Oxford, U.K., 1998. 83. Allison, F. The effect of wave-length on the differences in the lags of the Faraday effect behind the magnetic field for various liquids. Phys. Rev. 1927, 30, 66–70. 84. Allison, F.; Murphy, E. J. A magneto-optic method of chemical analysis. J. Am. Chem. Soc. 1930, 52, 3796–3806. 85. Allison, F.; Murphy, E. J. Evidence of the presence of element 87 in samples of pollucite and lepidolite ores. Phys. Rev. 1930, 35, 285letter. 39

86. Allison, F. Magneto-optic method of analysis as a new research tool. Ind. Eng. Chem., Anal. Ed. 1932, 4 (1), 9–12. 87. Allison, F.; Bishop, E. R.; Sommer, A. L.; Christensen, J. H. Further research on element 87. J. Am. Chem. Soc. 1932, 54, 613–15. 88. Allison, F.; Bishop, E. R.; Sommer, A. L. Concentration, acids and lithium salts of element 85. J. Am. Chem. Soc. 1932, 54, 616–20. 89. Allison, F. The magneto-optic method of analysis, with particular reference to the detection of elements 85 (alabamine) and 87 (virginium). J. Chem. Educ. 1933, 10, 71–78. 90. Urey, H.; Brickwedde, F.; Murphy, G. A hydrogen isotope of mass 2. Phys. Rev. 1932, 39, 164–66. 91. Annual survey of american chemistry, 1931; West, C. J., Ed.; National Research Council; Chemical Catalog Co.: New York, 1932; Vol. VI, p 40. 92. MacPherson, H. G. An investigation of the magneto-optic method of chemical analysis. Phys. Rev. 1935, 47, 310–17. 93. Brady, G. S. Materials handbook: An encyclopedia for purchasing agents, engineers, executives, and foremen, 4th ed.; McGraw-Hill Book Co.: New York, 1940; p 397. 94. Brady, G. S. Materials Handbook: An Encyclopedia for purchasing agents, engineers, executives, and foremen, 8th ed.; McGraw-Hill Book Co.: New York, 1956; p 666. 95. Langmuir, I. Pathological science: the science of things that aren’t so. Reprinted in Scientific Work and Creativity. Advice from the Masters; A presentation of the citizen scientists league. Smith, R. D., Ed.; Citizen Scientists League: Clearwater, FL, 2012; Vol. 1, pp 363−91. 96. Kauffman, G. B.; Adloff, J.-P. Fred Allison’s magneto-optic search for elements 85 and 87. The Chemical Educator. 2008, 13, 358–64. 97. Epstein, M. Using bad science to teach good chemistry. J. Chem. Educ. 1998, 75, 1399–1404. 98. The concise Columbia encyclopaedia, 5th ed.; Chernow, B., Vallaski, G., Eds.; Columbia University Press: New York, 1993; p 168.

40

Chapter 2

The Periodic Table of the Elements: A Review of the Future Paul J. Karol* Chemistry Department, Carnegie Mellon University, Pittsburgh, Pennsylvania 15213, United States *E-mail: [email protected].

Early in 2017, the remaining four elements of the Periodic Table were officially added, completing the seventh row with eka-radon, now oganesson. A perfectly natural vision of what the future holds in store is the subject of this appraisal. How will the Periodic Table continue to evolve? How much further will it go and how fast will new elements be found? Will any new elements be found? Where? Are there any in nature? How will new elements be synthesized? How will the synthesized products be measured? And for our particular interests, what chemistries are to be anticipated? The answers to these questions will constitute a review of our expectations – the future – based on various studies. Emphasis is on qualitative predictions rather than on the theoretical details that address both the nuclear and electronic aspects of structure and reactions.

The Periodic Table We are obligated to begin with a definition of the Periodic Table. In chemistry, the most respected source and official compendium of definitions is found in the “Gold Book” of the International Union of Pure and Applied Chemistry (IUPAC), i.e. the 2014 “Compendium of Chemical Terminology”. Ironically, there is no definition of the Periodic Table. Nevertheless, we proceed under the reasonable assumption that the readership is comfortable with their understanding of the Periodic Table, perhaps the quintessence of chemistry.

© 2017 American Chemical Society

Our focus is on the upper limit, “superheavy elements”, if any, on the number of elements that might be discovered or synthesized beyond those currently named (1–6). Physicist John Wheeler used the term “very heavy nuclei” in 1955 and subsequently used “superheavy nuclei” that same year for nuclei with a higher atomic number Z than 100, the highest known at the time. Many nuclear chemists define superheavy elements as being transactinides, Z >104 but there are arguments as to why they should refer to a mass number A >280 instead. A very good monograph on our subject is the 1990 publication “The Elements Beyond Uraniuim” by Glenn Seaborg and Walter Loveland. As an aside, note that both theorists and experimentalists in the entire heavy and superheavy element community shun the unwarranted IUPAC Greco-roman systematic naming system that extends to Z=999 (7). It is interesting to look back and review some predictions on the anticipated upper limit to the Table. The first fifty years of speculation on the upper limit are tabulated below in Table 1 based partially on Kragh (1). Note the Z=137 limit which will be discussed more thoroughly later.

Table 1. History of Predictions about the Upper Limit of the Periodic Table Year

Proposed by

Z limit

J. Newlands

1878

A → 480

E. Mills

1884

92

V. Meyer

1889

100

C. Baskerville

1901

93

S. Losanitsch

1906

>92

W. Tilden

1910

92

N. Bohr

1922

hundreds or thousands

S. Rosseland

1923

92

S. Rosseland and N. Bohr

1923

137

A. Sommerfeld

1924

137

The most recent fifty years shown in Table 2 continue the general lack of agreement as to where the end of the Periodic Table will emerge: soon or later, but likely well short of Z = 200. As of this writing, the number of known elements is 118 (8–10). The rate of discovery, disregarding the first dozen or so ancient elements, has been nearly constant over the past two and a half centuries at an average of one element every two and a half years as shown in Figure 1. Whether or not this rate of discovery will continue is implicitly discussed in the following sections. Assuming the discovery rate will continue as displayed, two centuries will elapse before a hypothetical 200th element (not necessarily Z=200) will be discovered, an unlikely achievement. 42

Table 2. History of Predictions about the Upper Limit of the Periodic Table Year

Proposed by

Z limit

G. Seaborg

1969

126

B. Fricke and W. Greiner

1971

172

A. Migdal

1974

> (137)3/2 ~ 1600

J. Berger et al.

2001

≈300

V. Nefedov

2006

164

A. Khazan

2007

155

J. Emsley

2011

128

W. Brodziński and J. Skalski

2013

126

Y. K. Gambhir et al.

2015

146

Figure 1. The cumulative number of elements as a function of time.

43

Transuranium Elements in Nature Since one of the co-sponsors of the Symposium on the Periodic Table is the Environmental Division, an appropriate exploration is to review the search for the existence in nature of the heaviest elements: the transuraniums. A summary is shown in Table 3 In 1948, G. Seaborg and I. Perlman extracted chemically purified neptunium and plutonium from a Canadian pitchblende ore (11). The measured alpha activity corresponded to 10-14 g 239Pu/g, most likely from neutron capture on natural uranium over the years. W. Grimm reports that in 1969 S. G. Thompson et al. searched for element 110, eka-Pt, in platinum ores but obtained negative results with a variety of detection methods (12), Also, that year (13), G. N. Flerov et al. found positive evidence for element 114, eka-lead in some lead-bearing samples by the observation of strongly ionizing events identifying fission in proportional counters and of fission tracks in plastics and glasses. Another attempt to discover superheavy elements involved minerals showing giant radiation damage halos hypothetically produced by high-energy alpha particles from transuraniums (14). Alternative explanations for the long-range tracks repudiated those interpretations. Characteristic X-rays were subsequently measured with PIXE (proton induced x-ray emission analysis) (15). Several of the observed x-rays were interpreted as originating from superheavy atoms with Z ranging from 116 to 126. However, this interpretation was refuted by subsequent independent experiments. In 1971, D. C. Hoffman et al. reported on 244Pu in nature (16). The best estimate for terrestrial abundance was 10-18 g 244Pu /g. This was consistent with both nucleosynthetic production 4.7 billion years ago and cosmic-ray influx. M. Jaschek and E. Brandi in 1972 observed spectral lines for Pu, Am and Cm in “peculiar A stars” (17). (A stars are young, main sequence stars.) Evidence for the possible existence of long-lived superheavy nucleus ekathorium with atomic mass number A=292 and atomic number Z ≈ 122 in natural Th was reported by A. Marinov et al. in 2007 (18). An abundance of about 1 x 10-11 relative to 232Th was reported in this work based on mass spectrometric measurements. R. Barber and J. De Laeter disputed the validity of the technique’s asserted sensitivity (19). Existence of long-lived isotopes of superheavy elements in natural Au was claimed by A. Marinov et al. (2011) (20). In essence they claimed to have observed 261Rg and 265Rg at an abundance level of (1-10) x 10-10 in gold. Rg is eka-Au. J. Lachner in 2012 instituting another search for 244Pu found an upper limit 15% of that of the positive finding by Hoffman’s 1971 study (21). Chemically etched radiation damage tracks of heavy nuclei in olivine from pallasite meteorites was used by Bagulya et al. in 2013 to assign approximately 6000 nuclear charges greater than 55 from galactic cosmic rays (22). Three superheavy nuclei were detected whose charge was within the range 105 < Z < 130. 44

Extensive accelerator mass spectrometry searches for 42 superheavy nuclides (A= 288-310) around the much discussed “island of stability” (Z=114, N=184) in natural Pt, Au, Pb, Bi materials were reviewed in 2015 by Korschinek (23). No positive evidence for the existence of long-lived superheavies (t1/2 > 108 yr) with abundance limits of 10-12 to 10-16 was found. Data on the composition of galactic cosmic-rays with Z > 70 were obtained in experiments by Ter-Akopian (2015) with the aid of Lexan-polycarbonate sheets used to detect the radiation damage tracks of incident heavy cosmic-ray nuclei (24). The high probability obtained for the existence, even at extreme trace levels, of Pu (and possibly Cm) nuclei in galactic cosmic-rays indicates a component of “freshly synthesized” (< 108 years old) nucleosynthetic transuranium matter. Ter-Akopian also discusses experiments carried out in Dubna, Russia found that flerovium is similar to noble gases in its chemical behavior (24). A 140-gram sample of xenon gas extracted from the atmosphere to look for rare events of spontaneous fission provided either by the flerovium nuclei or by their daughters was measured. The attainable limit of the flerovium concentration on the Earth amounts to 10-20 g/g assuming a billion year flerovium lifetime. Belli (2015) noted that searches by others for superheavy elements in Os, Pt and PbF2 established limit 3 were observed corresponding to a hassium limit concentration of < 10-14 g Hs/g Os with the standard assumption that the lifetime of eka-Os is ≈109 yr. Another search based on the similarity of the chemical properties of seaborgium to those of tungsten follows W in chemical purification processes including the growth of a ZnWO4 scintillation detector crystal allowing near 100% radiation detection efficiency. Signals induced in ZnWO4 by alpha particles are different from those caused by beta particles or gamma rays. The limit for a billion year seaborgium life time was 137 because E acquires an imaginary value. In Figure 15, the exact total E’s for the first few atomic orbitals are shown as a function of the nucleus atomic number with the cutoffs implied by the end of each energy curve. One-electron total energies (including rest mass, mc2) for a point nucleus in the figure correspond to total energy relative to mc2.

Figure 15. Total orbital energy (including inertial rest mass) of one-electron systems as a function of nuclear charge for a point nucleus. Below zero, the total E becomes imaginary. The 1s at Z=1 corresponds to E=-13.6 eV when the rest mass energy is ignored, for example. 59

In 1969 it was recognized that the Z=137 cutoff could be circumvented by introducing a finite-sized nucleus rather than a point nucleus (30). The one-electron total energies now followed the pattern shown in Figure 16 and place the 1s electron in the negative energy continuum in the presence of a nuclear charge >172.

Figure 16. One-electron total energy (including rest mass, mc2) for a finite-size nucleus. At E below –mc2, the negative energy continuum is breached and positron-electron pair production is predicted to occur.

The fate of the “vacuum” surrounding a completely ionized atom in overcritical Z cases evinces spontaneous pair production (which requires an energy of two electron rest masses = 1022 keV). The “vacuum” becomes negatively charged. Theorists find no stable atomic structures conceivable above the critical nuclear charge of 172. The previous illustration was for one-electron systems. A germane question is how the scheme might change when more than one electron is present. In Figure 17 below, the finite-size nucleus electron total energy (including rest mass, mc2) results are shown for atomic orbitals containing the 18 electrons of an argon-like system (31). Zcrit is essentially unchanged from the single electron result. 60

Figure 17. Total orbital energies for systems containing eighteen electrons as a function of non-zero-sized nuclear charge, Z. The 1s level dives into the negative energy continuum (-2mec2 = -511 keV) at Zcrit. The dashed extension curves are for Z > Zcrit.

Figure 18. Comparison of (theoretical) orbital radial density distributions with (solid) and without (dashed) relativistic inclusion in copernicium (Z=112). 61

Radial density distributions are shown in Figure 18 for the non-relativistic (solid curves) and relativistic (dashed curves) 7s, 7p, and 8s atomic orbitals calculated for 112Cn (31). These demonstrate the effects of relativity that include contraction of the s and p orbital sizes. Those reductions are accompanied by reciprocal expansions in the d and f and higher angular momentum orbitals. Figures 16 and 17 also reflect the increased spin-orbit splitting of the p, d, (f…) atomic orbitals. Theoretical approaches continue to improve but all still lead to electronic instabilities above Zcrit ≈ 172.

The Future A number of promising attempts to synthesize superheavy elements further out on the Periodic Table, eka-francium and eka-radium, have been made and were not successful. These are indicated below in Table 4.

Table 4. Unsuccessful Routes to New Elements 1985

254Es

+ 48Ca → 119X183-x + xn

2007

244Pu

+ 58Fe → 120Y182-x + xn

2009

238U

2011

248Cm

2011

249Cf

+ 50Ti → 120Y179-x + xn

2011

249Bk

+ 50Ti → 119X180-x + xn

+ 64Ni → 120Y182-x + xn + 54Cr → 120Y182-x + xn

So, what’s the problem? Neutron population. Fission and alpha-decay of the intermediate compound nuclei sidetrack the reaction pathway. But what if, in the future, laboratories are successful at overcoming technical challenges such as increased beam intensities (as seems on the horizon) and decay detection systems (which are already quite efficient). Suppose higher mass number and atomic number to Z=164 nuclei and maybe slightly beyond could be synthesized or their existence from astrophysical sources confirmed? What about the positioning of further superheavy elements onto the Periodic Table, the subject of this report? Three major attemptsat calculating electronic structure incorporating relativistic effects have appeared (32–34). These are reviewed in the depictions below in Figures 19, 20 and 21. 62

Figure 19. The Mendeleev-Seaborg extended Periodic Table prediction through Z=168 displayed in the spdf (shell partitioned display format) (35) following Seaborg (32).

Figure 20. The extended Periodic Table of Fricke et al. through Z=170 (33).

Figure 21. The extended Periodic Table through Z=172 following Pyykko (34). 63

In Figure 19, Z = 164 would be a p-block element; in Figure 20, Z = 164 would be an s-block element; in Figure 21, Z=164 would be a d-block element. Moreover, in all three variations of the Periodic Table, the element with atomic number Z = 125 has the valence electron configuration of 8s25g5. But to illustrate the emerging issue of assigning valence electrons, Nefedov calculated that the valence electrons occupy a mixed configuration describable as 81% 8s25g6f28p2 + 17% 8s25g6f7d28p + 2% 8s26f37d8p (36). If location on the Periodic Table is determined by valence electron configuration, where would mixed configuration assignments go? That question is currently unanswered. The author has a suggestion: Continue with the spdf (shell partitioned display format) Table structure, Figure 19, and allow for exceptions in electron configurations as is already explicit in the current Periodic Table. Despite the implied complexity on the horizon, peeking into the immediate future allows us to conclude with Figure 22, asserting some reasonable degree of confidence in its realization and correctness. Particularly likely are the alkali and alkaline earth chemistries of eka-francium and eka-radium. Beyond that expectations remain highly provisional (34, 37).

Figure 22. The extended Periodic Table in the near future through eka-radium (Z=120).

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22. Bagulya, A.; Kashkarov, L.; Konovalova, N.; Okateva, N.; Polukhina, N.; Starkov, N.; Gorbunov, S. Search for Superheavy Elements in Galactic Cosmic Rays. JETP Lett. 2013, 97, 708–719. 23. Korschinek, G.; Kutschera, W. Mass spectrometric searches for superheavy elements in terrestrial matter. Nucl. Phys. A 2015, 944, 190–203. 24. Ter-Akopian, G.; Dmitriev, S. Searches for superheavy elements in nature: Cosmic-ray nuclei; spontaneous fission. Nucl. Phys. A 2015, 944, 177–189. 25. Belli, P.; Bernabei, R.; Cappella, F.; Cerulli, R.; Danevich, F.; Denisov, V.; d’Angelo, A.; Incicchitti, A.; Kobychev, V.; Poda, D.; Polischuk, O.; Tretyak, V. Search for long-lived superheavy eka tungsten with radiopure ZnWO4 crystal scintillator. Phys. Scr. 2015, 90, 085301. 26. Wallner, A.; Faestermann, T.; Feige, J.; Feldstein, C.; Knie, K.; Korschinek, G.; Kutschera, W.; Ofan, A.; Paul, M.; Quinto, F.; Rugel, G.; Steier, P. Abundance of live 244Pu in deep-sea reservoirs on Earth points to rarity of actinide nucleosynthesis. Nat. Comm. 2015, 6, 1–9. 27. Rutherford, E. Collisions of α Particles with Light Atoms. IV An Anomalous Effect in Nitrogen. Philos. Mag. 1919, 37, 581–587. 28. Göppert-Mayer, M.; Jensen, J. Elementary Theory of Nuclear Shell Structure; Wiley: New York, NY, 1955. 29. Kowalski, M.; Kuchowicz, Br. On the Formation of Superheavy Nuclei (with Z=110) in Nature. Phys. Lett. B. 1969, 30, 79–80. 30. Pieper, W.; Greiner, W. Interior Shells in Superheavy Nuclei. Z. Physik. 1969, 218, 327–340. 31. Schwerdtfeger, P.; Pašteka, L. F.; Punnett, A.; Bowman, P. O. Relativistic and quantum electrodynamic effects in superheavy elements. Nucl. Phys. A. 2015, 944, 551–577. 32. Seaborg, G. Prospects for Further Considerable Extension of the Periodic Table. J. Chem. Ed. 1969, 10, 626–634. 33. Fricke, B.; Greiner, W.; Waber, J. T. The Continuation of the Periodic Table up to Z = 172. The Chemistry of Superheavy Elements. Theor. Chim. Acta. 1971, 21, 235–260. 34. Pyykkö, P. A suggested periodic table up to Z ≤ 172, based on Dirac-Fock calculations on atoms and ions. Phys. Chem. Chem. Phys. 2011, 13, 161–168. 35. Karol, P. The Mendeleev–Seaborg Periodic Table: Through Z = 1138 and Beyond. J. Chem. Ed. 2002, 79, 60–63. 36. Nefedov, V. I.; Trzhaskovskaya, M. B.; Yarzhemskii, V. G. Electronic Configurations and the Periodic Table for Superheavy Elements. Dokl. Phys. Chem. 2006, 408, 149–157. 37. Pershina, V. Relativistic electronic structure studies on the heaviest elements. Radiochim. Acta. 2011, 99, 459–476.

66

Chapter 3

How the Periodic Table of Available Elements Shaped Natural History Benjamin J. McFarland* Department of Chemistry and Biochemistry, Seattle Pacific University, Seattle, Washington 98119, United States *E-mail: [email protected].

The narrative of natural history, normally told in terms of geology and biology, can be better understood by accounting for environmental and chemical factors represented in the periodic table. The inorganic chemist R.J.P. Williams wrote extensively about how the elements were constrained in form, function, and availability throughout deep time. Geological and biological observations from natural history can be explained with chemistry. For example, the distinction between “old” elements used by simpler, earlier microbes and “new” elements used by more complex organisms can be explained using solubilities and redox potentials.

Introduction The title “Elements Old and New” provokes two questions: “How old?” and “How new?” Ancient humans recognized several “old” elements thousands of years ago that were isolated and purified by geological processes themselves billions of years old. Those same processes actively sorted elements and made them available in different environments throughout natural history. Some elements predominated the ecosphere four billion years ago and can be considered truly old, while others came to predominate a mere billion years ago and can be considered new (or at least newer). The billion-year scale of natural history alters the common relationship between chemistry and biology. In broad terms, biology studies living organisms on the meter-to-micrometer scale, biochemistry studies macromolecules on the micrometer-to-nanometer scale, chemistry studies molecules and atoms on the © 2017 American Chemical Society

nanometer-to-picometer scale, and particle physics studies subatomic particles at even smaller scales. However, large scales in time and space include geochemical changes of large amounts of atoms and molecules, which can be explained by appealing to chemical priciples of reactivity such as periodic trends. This is chemistry on the megameter scale, describing reactions that influenced the entire planet. At this scale, chemical changes defined the geological arena in which biological competition, mutation, and evolution could take place. Life filled the available space and flowed through diverse environments on the planet — yet it was all built from the same periodic table, on which, for example, oxygen holds a unique place in its distinctive abundance and reactivity. Constant chemical laws, expressed through elemental availability, shaped the contingencies and uncertainties of individual organisms and ecosystems. In this essay, the title “Elements Old and New” is interpreted ecocentrically rather than anthropocentrically. Elemental isolation, discovery, and refinement are usually considered to be human activities. However, if natural processes may be described as active, then a natural history can be told in which elements were isolated, discovered, and before humans evolved:

1. 2. 3.

Isolation. The oceans isolated different elements in the liquid phase according to their solubility and abundance. Discovery. Living cells discovered the chemical utility of particular elements for novel reactions through natural selection. Refinement. Through pumps and proteins, microbes, animals, and plants refined certain elements into particular extracellular and subcellular compartments.

As a result, the earth over time underwent a redox-dependent sequence of elemental availability, dominated by certain elements long ago and different elements more recently, as the periodic table of elements available to life expanded and contracted. I have described this natural history at book length for a general audience in A World from Dust: How the Periodic Table Shaped Life (1). This essay provides new data supporting this chemical narrative by emphasizing how the Earth’s changing environment shaped chemical availability, which in turn shaped the chemical capabilities and functions available to life. In this way, the periodic table continues to explain the grand arc of the narrative of natural history.

Chemical Logic Explains Elemental Abundance in Different Environments The periodic table is by definition incomplete, because larger elements can always be formed, but it has become more complete over time as more elements were made. The first periodic table would contain two elements representing the mix of hydrogen and helium immediately after the Big Bang. Stellar nucleosynthesis fused lighter elements into heavier elements (2), and the periodic 68

table grew over time into the diverse set of 90 naturally occurring elements that now constitute the environment. Even 13 billion years of this process has not been sufficient to alter the overall dominance of the two lightest elements; for example, in a diffusion cartogram of the periodic table weighted by current abundance in the universe, only hydrogen and helium can be clearly seen (3). A graph of universal abundance organized by atomic number is also dominated by these two elements, even when using a logarithmic scale (Figure 1). The universe differentiated into diverse environments, with different proportions of elements sorted into each by physical processes. Within stars, gravity pulled heavier elements toward the core. Gravity also pulled heavier elements together outside of stars, so that rocky planets like the Earth are enriched in heavier elements like magnesium, calcium, iron, aluminum, silicon, and oxygen. Another diffusion cartogram of the periodic table, this one weighted by elemental abundance in the Earth’s crust, lists these six elements most prominently (3). The same six elements comprise Robert Hazen’s “Big Six” classification of abundant geological elements (4). Gravity collected these rocky elements together in the proto-Earth as the solar wind blew hydrogen and helium into interplanetary space. From the earliest days of the Earth, the periodic table of elements abundant on the planet was arguably more complex than that of most of the universe. As the Earth cooled, a solid crust formed on its liquid mantle, on which liquid water condensed to form oceans beneath a nascent, gaseous atmosphere. Three of these four environments were fluid, allowing elements to mix, react, and exchange. Even the solid environment of the crust was mixed by subduction into the mantle, moved through the processes of plate tectonics. Because of this contact and mixing, each of these environments gathered a different combination of available elements through a combination of chemical properties such as melting points and solubility products. Because these dynamic processes form stable, persistent cycles such as the water cycle, together they resemble a chemical reactor, so that environmental chemists are fond of describing the earth as a “giant chemical reactor.” This exact phrase can be found in diverse places in the atmospheric chemistry and geochemistry literature (5–7), because it describes well how chemical laws mixed and sequestered different forms of the elements in different environments. The limited availability of elements and distinct patterns of chemical reactivity combine to imply that all possible combinations of elements can not be realized in all environments. The partitioning of elements into different environments can be understood with chemical logic, by comparing chemical properties and thermodynamic parameters. One example of this explains why free ozygen was absent from the ancient atmosphere. Both oxygen and sulfur would have been abundant on the early earth, and each could have combined with many elements because of the physical processes mixing the planet’s environments. If for each available element, we compare the heats of formation of oxides to heats of formation of sulfides, a thermodynamic advantage is found for oxygen reacting with aluminum, silicon, carbon, and magnesium because of their larger standard enthalpies of formation relative to their sulfides (Figure 2). 69

Figure 1. Relative abundance of all elements in the universe, arranged by atomic number. Reproduced with permission from ref. (1). Copyright 2016 Oxford University Press.

These four elements would have been so abundant on the surface of the early earth that most of the oxygen would have reacted with them, leaving little left to react with the other elements and still less to reside in the atmosphere as O2 gas. Therefore, the low levels of oxygen on the early Earth despite oxygen’s elemental abundance, as observed by multiple lines of geological evidence (4), can be assigned a chemical explanation on the basis of thermodynamic reference tables. When the earth was first formed, oxygen was incorporated into the Earth’s crust as oxides and into the oceans as water, and its free, diatomic form was essentially absent from the atmosphere. After that, over billions of years, oxygen levels increased to today’s levels, as shown by several lines of geological evidence, including the presence of iron oxide in the global geological oxidized precipitate known as the Banded Iron Formations (Figure 3). 70

Figure 2. Elements sorted by preference for oxygen over sulfur, calculated from standard heats of formation. Reproduced with permission from ref. (1). Copyright 2016 Oxford University Press.

Life Requires Elements That Are Both Available and Functional In the anoxic Archaean environment, microbial life harnessed and dissipated energy gradients in order to persist and do work, which required a persistent chemical environment inside the cell relative to the fluctuations in the external environment. Biochemical processes used the liquid phase in order to respond to fluctuations in the gradients, and so the material used for life’s reactions had to be soluble in water. Many have hypothesized that thermal and geochemical gradients at underwater vents would have provided the matter and energy needed for life (8), although others suggest that geothermal pools on the surface could have served a similar function (9). Wherever it happened, life emerged and evolved multiple chemical mechanisms to persist, drawing material from the environment in order to build and maintain dynamic, complex structures of proteins, sugars, lipids, and DNA. Biological possibilities were constrained and restricted by the chemical form, reactivity, and availability of elements, as illustrated by the “arsenic life” controversy. In late 2010, a team of scientists presented evidence from Mono Lake in California, a high-arsenic, low-phosphorus environment, that bacteria not only tolerated high concentrations of arsenic but used arsenic instead of phosphorus (10), possibly incorporated into the DNA backbone as arsenate. This was contested because arsenate esters have much faster hydrolysis half-lives in water 71

than phosphate esters. Arsenate is therefore much less suitable for maintaining DNA structure in solution, despite its chemical similarity to phosphate (11). Later experiments supported the conclusion that the bacteria from Mono Lake do not incorporate arsenate and remain dependent on phosphate (12, 13), like all known life forms. Phosphate-binding proteins from this particular bacterials strain are able to discriminate subtle structural differences between arsenate and phosphate by precise placement of hydrogen bonds (14), allowing arsenate to be rejected even in high-arsenate environments. The availability of arsenate in Mono Lake did not overcome its instability and therefore its inability to build structures persistent enough for aqueous life.

Figure 3. Oxygen levels over time, expressed as percentage of modern atmospheric levels and aligned with significant geological and biological events. Reproduced with permission from ref. (1). Copyright 2016 Oxford University Press.

The interior of a living cell is separated from the environment by a semipermeable membrane, perforated with channels and pumps that maintain the internal environment. Bacteria persist in a high-arsenate environment by evolving these proteins to maintain internal consistency despite external challenges. The prominent bioinorganic chemist R.J.P. Williams (1926-2015) proposed that the elements kept inside each cell are selected based on availability and function, and he outlined three broad functional categories for the different biochemical uses of the elements (15), which can be re-stated as the following: 72

1.

2.

3.

Balance. Highly soluble elements must be pumped across the membrane to equalize osmotic pressure. Large amounts of these elements are required. Building. Elements that can form covalent bonds are used to make stable structures inside and immediately outside the cell. Moderate amounts of these elements are required. (Calcium, sulfur, and phosphate can be used for both balance and building and therefore are considered members of both categories.) Biochemistry. Elements with useful chemical properties, such as transition metals, can be used to catalyze particular reactions. Trace amounts of these elements are required.

Williams argued that cells selected phosphate rather than sulfate or arsenate for the backbone of DNA because of phosphate’s availability and chemistry (e.g., its ability to form stable phosphate esters, its flexibility, its negative charge, etc.), which are well suited for the information storage and transfer functions of DNA (15). The selection of phosphate led to the selection of magnesium to balance phosphate’s negative charge, because magnesium binds well to phosphate (e.g., the ionic radius of magnesium fits well into the spacing of oxygens in a diphosphate ester, as seen in the association of magnesium with ATP as Mg-ATP) (16). Williams theorized that the presence of charged phosphate and magnesium inside the cell required the rejection of sodium and the retention of potassium for osmotic balance in the chemical environment of the ocean (17). With this style of reasoning, Williams gave a chemical rationale for the elemental concentrations inside living cells. Broad chemical features of the environment derived from the periodic table are predictable in a way that specific biological features like gene sequences are not. The first discovery that Williams published (from research accomplished as an undergraduate) was the Irving-Williams series describing the stability of transition-metal complexes with various ligands (18). Williams later described how this series predicted the free transition-metal ion concentrations inside cells, because the inorganic chemistry of transition metal ion binding is universal despite extensive biological variability (17). Cellular pumps and channels must maintain levels of metal ions at the concentration where there is enough of the ion to bind and use, but not so much that it cross-links and congeals the fluid and dynamic structures necessary for life. Williams’s insight that large-scale chemical predictability co-exists with and shapes biological unpredictability can be seen in cases in which the environment shapes life, such as in recent large-scale “-omics” studies. Multiomic sequence information of microorganisms in the Saanich Inlet were explained with a biogeochemical model in which a chemical feature of the environment (elemental fluxes across a redox gradient) determines microbial community structure and gene expression patterns (19). In a global study of the ocean microbiome, the geochemical category of environmental conditions, including the elemental forms and concentrations that shape biological niches, was found to be more strongly determinative and predictive of the patterns of microbial life than the biological category of taxonomy (20). On a different scale, a study of microbial communities 73

sampled along an intercontinental cycling trip concluded that local environmental factors select from a pool of globally distributed microbial species, so that “most of the [biological] variation can be explained by environmental and not spatial patterns (21).” This may apply to how cancer responds to its environment as well. When metastatic cancer cells encounter a new environment in a different part of the human body, tumors originating from different primary sites will predictably converge to similar phenotypes, being shaped by the adaptive landscape of the local environment (22). Biological species and genes, and geographical spatial patterns, are contigent upon various degrees of unoredictability, but the chemistry of the environment plays a major role in shaping the characteristics of life in a manner consistent with the chemistry of the elements. This implies that when the chemistry of the planetary environment changed fundamentally over billions of years and as the array of elements available from the periodic table changed, that these environmental-chemical changes could have fundamentally reshaped life.

Figure 4. Concentrations of elements in the modern ocean, with elements categorized by biological functions of balance (dark gray), building (medium gray), and biochemistry (light gray). Reproduced with permission from ref. (1). Copyright 2016 Oxford University Press. 74

Environmental Availability Restricted the Biochemical Periodic Table Differently Billions of Years Ago In both ancient and modern oceans, the more soluble forms of the elements would have been more available for biological functions. This solubility depends on the presence of counterions that may form less soluble precipitates and is dependent on the concentrations of different elemental forms. Williams noted that his three broad categories of elemental function correlate with aqueous solubility in the ocean today (Figure 4) (15): 1. 2. 3. 4.

Elements available at millimolar solubilities are used for balance. Elements available at micromolar concentrations are used for building. Elements available at nanomolar concentrations are used for biochemistry and catalysis. Elements available at concentrations lower than nanomolar are not used.

Concentrations in the ocean are directly related to elemental abundance in the universe, which is why most of the elements used for biochemical processes are those with smaller atomic weights, which have been formed in more abundance from hydrogen and helium by stellar nucleosynthesis (Figure 1). The three elements present in two categories in Figure 4 each have chemical rationalizations. Calcium’s availability in the low-millimolar/high-micromolar range of concentrations and its high charge density explain its ability to provide both balancing and building functions. Sulfur is used for balance as sulfate, while it is used for building as sulfide (e.g., in proteins as disulfide bonds). Phosphate’s unique structural and energetic utility for building DNA means that it is enriched in the cell and must be accounted for in electrostatic balance as well as in biomolecular building. Several elements are redox-sensitive, and as the global redox state changed, their availability would have changed, and so would their use in life. The vertical arrows on Figure 4 show how Williams predicted that redox-sensitive elements would have responded to increasing levels of free oxygen in the atmosphere and ocean. Over billions of years of oxidation, iron, nickel, and cobalt decreased, while molybdenum and copper increased. Also, reduced forms of carbon, sulfur, and nitrogen would have been more prominent in the ancient ocean compared to the more oxidized forms present today. If these changes in concentrations were significant enough to change the availability of these elements in the ocean, their patterns of usage in life would have changed as well. Abundance and solubility influence availability and function, restricting the biochemical periodic table today to about two dozen elements. The anoxic environment three billion years ago would have resulted in a different emphasis on the table, especially in the transition metal block, with the elements on the left of the block used more and the elements on the right used less. As predicted by Williams based on solubility constants and redox potentials, older elements included iron, cobalt, nickel, and reduced forms of carbon, sulfur, and nitrogen (especially methane, sulfide, and ammonia) while newer elements included copper, molybdenum, and oxidized forms of carbon, sulfur, and nitrogen 75

(especially sulfate and nitrate) (23). This sequence is illustrated in Figure 5 by an arrow that places the elements in approximate order of oxidation from a reduced form to an oxidized form. The x-axis in this figure is both redox potential and time.

Figure 5. Redox potential (bottom) correlates with old and new forms of the elements used by life (top). Reproduced with permission from ref. (1). Copyright 2016 Oxford University Press. These predictions were made by Williams based on chemical properties and biochemical observations of how the metals used by archaea and ancient bacteria differ from those used by animals today. In the first decade of the twenty-first century, enough genomic data had accumulated that this chemical sequence could be more rigorously tested by multiple authors (24–26). These analyses upheld the general sequence shown in Figure 5, with more reduced forms of the elements used early in natural history and more oxidized forms used later, and with the overall patterns of metalloprotein usage in line with the previous publications by Williams.

Recent Evidence for How Oxygen Affected the Ancient Environment The chemical sequence of elements old and new throughout natural history is a continuing area of research. For example, recent studies using selenium isotopes support the general sequence of oxygen levels shown in Figure 3, including transient “whiffs” of oxygen before the Great Oxidation Event accompanied by low ocean oxygen levels even during the “overshoot” in atmospheric oxygen levels that followed that event (27–29). A more contested point is more whether increased atmospheric oxygen was a cause or effect of the widespread diversification of complex animal fossils termed the Cambrian explosion, as 76

described in Chapter 9 of A World from Dust (1). An interesting new technique measuring oxygen levels in halite inclusions supports high oxygen levels before the Great Oxidation Event (30), providing additional evidence that oxygen preceded biological complexity and could have been its cause. Most investigations to date have focused on oxygen levels, but recent investigations have started to determine the levels and forms of other elements. A pattern in nitrogen isotopes coincident with the Great Oxidation Event was analyzed and interpreted as showing that the rise in oxygen led to an increase in nitrate, allowing evolution of new aerobic metabolic cycles based on this element (31). This geochemical evidence supports Williams’s chemical sequence, and provides initial evidence for how oxygen shaped both the geochemistry and biochemistry of another element. One story recounted in A World from Dust that was challenged by recent evidence is from Chapter 9, in which molecular clock genetic evidence was used to suggest that newly-evolved lignin-degrading peroxidases caused a decrease of carbon burial and the end of the Carboniferous era (32). However, a later analysis of the genetic evidence concluded that, as the title of the study put it, “Delayed fungal evolution did not cause the Paleozoic peak in coal production (33).” Another study highlighted the complexity of lignin degradation genes, emphasizing that there are many genes with poorly understood functions and implying that our understanding is not complete enough to encapsulate in a simple cause-and-effect story (34). Our understanding of the interplay between chemistry and biology continues to evolve in cases like this.

The Unique Perspective of Chemistry on Natural History Throughout the narrative of A World from Dust, Williams’s ideas are contrasted with those of the naturalist Stephen Jay Gould. Gould explained the narrative of evolution in his book Wonderful Life with the metaphor of rewinding and replaying the “tape of life” to repeat natural history (35). Gould’s ultimate conclusion emphasized biological contingency: “But I suspect, from the rarity of Pikaia in the Burgess and the absence of chordates in other Lower Paleozoic Lagerstiitten, that our phylum did not rank among the great Cambrian success stories, and that chordates faced a tenuous future in Burgess times. … And so, if you wish to ask the question of the ages -- why do humans exist? -- a major part of the answer, touching those aspects of the issue that science can treat at all, must be: because Pikaia survived the Burgess decimation. … The survival of Pikaia was a contingency of ‘just history.’ I do not think that any ‘higher’ answer can be given, and I cannot imagine that any resolution could be more fascinating.” (p. 322-323) Although my perspective as a chemist causes me to de-emphasize the role of biological contingency, I agree that Gould’s “tape of life” metaphor expresses deep truths about our interconnected and complex universe. This is not a debate about 77

the facts but about the prioritization and interpretation of the facts. Yet Gould’s fundamental conclusion in this case is built upon one type of observation: the relative scarcity of the protochordate Pikaia in the fossil record at certain times. Simon Conway Morris, cited by Gould as the scientist who first “reached a firm conclusion” about the nature of Pikaia (35), interprets the same data differently and emphasizes the predictable importance of evolutionary convergence. Conway Morris argues that some Ediacaran body plans were lost during the Cambrian era because they were less efficient and more poorly suited to the increased competition after the Cambrian explosion (36). This is a biochemical judgment built partly on chemistry, including the ability of biological structures to utilize carbon and oxygen. If chemical inefficiency provides a reason why these body plans were removed from the array of possible biological shapes, then the same chemical inefficiency would remove those same shapes upon repetition of the tape of life. There may be similar chemical and/or environmental factors that led to the selection of Pikaia despite its unassuming profile in the fossil record. If so, then the proliferation of chordates would be more predictable than Gould concluded because of the ‘higher’ reason of chemical necessity. Evidence for chemical constraints shaping biology can be found in the records of natural history, as predicted by chemists like Williams. In the case of the Cambrian Explosion, current evidence suggests that a major geochemical event involving atmospheric oxygen increase immediately preceded a period of rapid biological innovation. Biological complexity can be connected to both the energy and chemical structures provided by oxygen, and to the chemistry of oxidized elements and elements released by oxidative weathering. If geological evidence continues to place oxygen increase before biological complexity, then the unique chemical characteristics of oxygen can connect the two in a causal chain, turning what is now correlation into something that may be considered chemical causation. Such a chain of reasoning would apply across the universe and would predict a Cambrian explosion event wherever photosynthetic water-based life can persist for billions of years, producing eventual oxygen increase. One concern with an emphasis on the predictability of natural history is that it may smuggle unwarranted teleology into scientific explanations, supervenient on the non-teleological processes of evolution. However, some scientists have proposed ways that evolutionary direction could emerge from large-scale physical constraints and laws. R.J.P. Williams is one example; another is Terrence Deacon, who theorized in Incomplete Nature about how goal-oriented behavior can develop from stochastic processes, in the context of functional and adaptive organisms (37): “… natural selection is indeed a thoroughly non-teleological process. Yet the specific organic processes which this account ignores, and on which it depends, are inextricably bound up with teleological concepts, such as adaptation, function, information, and so forth.” (p. 137) “Thus, although the evolutionary process is itself non-normative (i.e., is not intrinsically directed toward a goal), it produces organisms which are capable of making normative assessments of the information they receive.” (p. 413) 78

Deacon’s conclusions apply to how the random proliferation of particular genes and body plans may be constained by the chemistry of the environment to give functions, goals, and an overall direction to evolution. Chemical laws impose constraints on the available elements and the types of reactions those elements can undertake, eliminating possibilities and restricting evolution. Deacon calls these processes “absential” because they are primarily shaped by what is absent from the system rather than what is present. In living organisms, predictability can emerge from randomness through constraint in a pattern consistent with Deacon’s arguments. For example, mammalian cells transcribe genes in a messy and unpredictable process, but transport through the nuclear membrane dampens the random fluctuations and results in transcript levels in the cytoplasm that are tightly adapted to the microenvironment (38), and therefore predictable from detailed knowledge of the chemical and biological features of the microenvironment. Here, nuclear transport constrains the randomness of transcription, keeping the random fluctuations in the nucleus absent from the cytoplasm, which produces predictable results in the cytoplasm. (This is similar to how the pumps and metal-binding proteins in a cell maintain free metal ion concentrations consistent with the Irving-Williams series.) Deacon applies these concepts to explain how biological function and goal-directed behavior can emerge from evolution’s random processes, such as those involved in gene mutation and processing. Deacon’s concepts of functional emergence complicate simple identification of cause and effect. Instead of causality, he discusses goal-oriented behavior in terms of work, which is physically defined as directed movement: “Work is a more complex concept than mere cause, because it is a function of relational features. And relations, both actual and potential, are precisely where the action is. More important, work is only possible because of limitation. To abuse an old metaphor: the fabric of mind is not merely the thread that composes it.” (p.141). If directed movement and work can emerge from random processes that are limited by external constraints, then direction in evolution may emerge as well from global chemical constraints. In Deacon’s ideas, the keys to explaining how life shaped the planet and the planet shaped life can be found in relational. absential features such as feedback loops and the need for life to operate within the constraints imposed by its environment. Some have suggested that life has survived for so many billions of years on this planet because global feedback loops have allowed the planet to traverse a “Gaian bottleneck” as the sun’s luminosity changed over time (39). This interpretation sees the Earth’s surface oceans and atmosphere as an effect rather than a cause of life. Chemistry constrains all of these processes, including the Gaian bottleneck, because they take place within the chemical constraints of phase diagrams: the temperatures and pressures needed to retain surface oceans and a thick atmosphere given the energy output of the sun. To use Deacon’s terminology, this kind of planetary feedback loop is absential and directed by the requirements of dynamic persistence within a set of physical-chemical constraints. A chemical perspective that takes into account patterns of availability, solubility, and reactivity connects to both geological and biological perspectives on natural history, providing explanatory insights in a conversation with geology, 79

biology, and other fields of knowledge. The rise of oxygen and subsequent global chemical transformation distinguishes old elements from new over the course of natural history. The geological effects of this change are well documented, while the biological causes and effects are still being disentangled. The importance of chemistry in understanding also implies that this perspective is useful in teaching, which is why I have used these stories in physical chemistry, biochemistry, and science writing courses to help student learning. When natural history is discussed, the discussion should include not only rocks and animals, but the periodic table of elements as well.

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Chapter 4

An Elemental History of the Early Universe E. Prasad Venugopal* Department of Chemistry and Biochemistry, University of Detroit Mercy, 4001 W. McNichols Road, Detroit, Michigan 48221-3038, United States *E-mail: [email protected].

Elemental properties play a crucial role in constraining theories of the early universe in cosmology. Attempts to explain the creation and abundances of the elements combined aspects of cosmochemistry, nuclear physics and astronomy to produce the Big Bang theory of the evolution of the universe. This chapter provides a brief overview of historical and contemporary research on elemental chemistry in cosmology.

Introduction In 1927, Georges Lemaitre, a Belgian astrophysicist and Catholic priest, published a new set of analytical solutions to Einstein’s equations of General Relativity. Lemaitre’s solutions (1), describing an expanding universe filled with matter and radiation seemed to suggest that the universe was neither static nor infinitely old. As such, Lemaitre’s work, and that of Soviet mathematician and physicist Alexander Friedmann who independently derived a similar set of solutions in 1922 (2), represented a dramatic departure from earlier cosmological models (3, 4). In solving the equations of General Relativity, Einstein was led, based on existing empirical evidence, to a universe that was spatially limited but temporally infinite and filled with a finite matter density. As is well-known, the introduction of an “ad hoc” cosmological constant into the field equations provided a repulsive force to balance the gravitational attraction of matter that otherwise predicted the collapse of the universe, a contribution that Einstein found to be “detrimental to the formal beauty of the theory” (5). The expanding universe models of Lemaitre and Friedmann did not receive much attention due to a variety of reasons, including the lack of observational evidence (6). However, that situation changed in 1929 when American astronomer, Edwin Hubble, published a study showing a linear relationship © 2017 American Chemical Society

between galactic distances and the redshift of spectral lines from those galaxies (7). Through the Doppler effect, Hubble calculated the recessional velocities of galaxies from their redshifts and related them to galactic distances through Hubble’s law:

where V is the recessional velocity in the radial direction in km/s, D is the distance of the galaxy in Megaparsecs (1 Mpc = 3.26 million light-years) and H0 is called the Hubble constant. Initially Hubble calculated the value of H0 to be about 500 km/s/ Mpc. While Hubble’s observations could not identify the source of the redshift, it was soon recognized that Hubble’s Law was possible evidence for the expansion of space, as predicted by General Relativity (8). The cosmological timescale during which the expansion occurred was then easily calculated as the inverse of the Hubble constant, and provided an approximate (model-dependent) value for the age of the universe of 1-2 billion years. The scientific debates and controversies about a universe that seemed to evolve from a putative beginning began in earnest in the 1930s (6). Cosmologists were confronted with the challenge of building a finite-age universe populated by chemical elements whose relative abundances were known from terrestrial and cosmic sources. Elemental properties played a crucial role in constraining theories of the early universe that attempted to answer two fundamental questions: how were the known elements created in the universe; and, what mechanism(s) could explain the relative abundances of these elements? While the implications of these issues were far from clear in the early decades of twentieth-century cosmology, attempts to explain the creation and abundances of the elements continue to play a determinant role in viable theories of the evolution of the universe. This chapter provides a brief overview of historical and contemporary research on elemental chemistry in cosmology.

The Role of Radioactivity in Early Cosmology Henri Becquerel’s discovery of radioactivity in 1896 (9) undermined the notion of stability inherent in the Periodic Table of elements. A flurry of experimental activity followed, resulting in the discovery by Pierre Curie that radium salts emitted a constant stream of heat energy (10). Ernest Rutherford and his collaborators demonstrated that radioactive materials transmutated to other elements with the emission of radiation (11). These experiments opened new frontiers in solar-system physics and cosmology, resolving a long-standing controversy in the former while creating a new one for evolutionary models of the universe. Nineteenth-century science was witness to a sharp disagreement over the age of the Earth (12). The physicists, led by Lord Kelvin, estimated the age of the sun to be about a hundred million years based on a premise that the source of the 84

sun’s heat was solely due to gravitational contraction. Kelvin followed up with a calculation that placed a similar upper-limit on the age of the earth by treating it as a sphere that began in a primordial heated state and gradually cooled to its present state. This timescale was far shorter than that required by Charles Lyell’s theories of geology as well as the theory of biological evolution championed by Darwin and his followers. The dispute was resolved by the discovery of radioactivity in terrestrial elements, such as uranium, contradicting Kelvin’s premise that no other sources of heat were present in the earth’s interior. In subsequent years, radioactive dating of the earth’s crust provided more consistent estimates of the age of the earth, consistently placing it at a few billion years (13). Presenting his view that the paradox had been resolved, Rutherford concluded: “The discovery of the radio-active elements, which in their disintegration liberate enormous amounts of energy, thus increases the possible limit of the duration of life on this planet, and allows the time claimed by the geologist and biologist for the process of evolution” (quoted in (12)). The recalculated age of the earth however, posed a serious problem for cosmologists, who realized that even small changes in the density of matter or radiation in the equations of general relativity predicted a universe that was not only in conflict with the Hubble time, but was younger than the earth and sun. The independent determinations of the Hubble value and the half-lives of radioactive elements meant that this was not a problem that was easily resolved. Writing in 1952, cosmologist Hermann Bondi wrote: “[F]or more than fifteen years, all work in cosmology was affected and indeed oppressed by the short value of [the age of the universe] T (1.8 x 109 years) so confidently claimed to have been established observationally. The time-scale difficulty, as the discrepancy between T and the ages of the Earth and stars was called, of cosmological theories, and the effects of the removal of this influence, have not yet been worked out fully (14).” Nevertheless, the phenomenon of radioactivity caught the attention of cosmologists, and some prominent chemists, as a mechanism for creating the known elements in the universe. Though Rutherford was referring to the ages of the earth and sun, his 1904 suggestion that “it is natural to ask what part radio-active substances play in cosmical physics” (15) foresaw the emergence of the field of cosmochemistry. Lemaitre’s model of 1931 suggested that the universe began as a “unique” and “unstable” atom that subsequently divided into “smaller and smaller atoms by a kind of super-radioactive process” (6). Others, including Walther Nernst, Gilbert Lewis and James Jeans, suggested similar mechanisms for the production of elements in stars and the early universe based on the disintegration of heavy elements into lighter one, but in the absence of a viable theory of the nucleus, these ideas remained highly speculative. The situation changed in the late twenties when multiple physicists, including the nuclear theorist George Gamow, provided a theoretical explanation of 85

radioactive alpha-decay through the quantum-tunneling effect in which an alpha-particle was emitted by an unstable nucleus despite being forbidden by classical energy conservation principles. Soon, it was recognized that a reverse nuclear process could be responsible for synthesizing the heavier elements from lighter ones (16), a notion that was solidified by the discoveries of the neutron and deuterium in 1932. The advent of the field of nuclear astrophysics was instrumental in the post-war period in constructing theories of primordial and stellar nucleosynthesis.

Primordial Nucleosynthesis and Stellar Thermodynamics Questions about elemental constituents and energy production in stars had long interested both cosmologists and chemists, a project that began in the early nineteenth century when Joseph von Fraunhofer produced a detailed solar absorption spectrum. The advent of improved equipment and experimental techniques in the early twentieth century significantly aided the study of stellar spectra. In 1925, Cecilia Payne completed a detailed study of the spectral characteristics of the sun as part of her dissertation (17) for a Ph.D. in Astronomy. Payne demonstrated that hydrogen and helium were the dominant components in the solar atmosphere, and by extension stellar interiors and the universe. Further observations of stellar spectra confirmed her results. In 1937, geochemist Victor Goldschmidt (18) published the results of a decade-long study in which he presented detailed tables of data on terrestrial and cosmic abundances of most elements then known to exist. Goldschmidt’s study was a major contribution to the growing field of cosmochemistry (19, 20). Attempts to reproduce the relative abundances of matter and radiation in the universe were initially based on the idea of modeling the universe as a system in thermodynamic equilibrium. Among the earliest theorists to apply this concept were the chemists Harold Urey, Charles Bradley and Richard Tolman. In a 1931 paper (21), Urey and Bradley calculated the abundances of terrestrial isotopes on a hypothesis of “equilibrium by a number of transmutation reactions” and found little agreement with data. Tolman, a mathematical physicist and physical chemist at Caltech, combined special relativity and thermodynamics to calculate “the possible formation of helium out of hydrogen in accordance with a quasi-chemical reaction” and the “possible transformation of matter into radiation.” Disappointingly, Tolman’s calculations led to either a helium-filled universe or one virtually devoid of matter (22). Nuclear physics had flourished in the nineteen thirties into a mature field based on the pioneering work of theorists such as Hans Bethe and experimentalists such as Enrico Fermi who bombarded various elements with neutrons to induce radioactivity by neutron capture. The impetus to apply the newly developed quantum theory of nuclear reactions to explain cosmic elemental abundances received a boost with the publication of a paper by Cornell University nuclear physicist, Hans Bethe, in 1939. Using previous estimates of stellar temperatures, Bethe identified two sources of energy production in stars (23), believed to be hot, dense gases in equilibrium. These processes, the proton-proton chain for 86

stars with lower core temperatures, and the carbon-nitrogen chain for hotter stars, were responsible for producing He4 (identified as α-particles in his paper) from hydrogen. He found the agreement with observational data from stellar spectroscopy to be “excellent”. However, Bethe made a categorical claim in his paper that “no elements heavier than He4 can be built up in ordinary stars” because both processes that he considered led to the production of He4. The non-existence of stable nuclei at atomic weights of five and eight precluded the building up of heavier elements through proton capture by helium. Bethe considered other possible reactions to bypass this elemental roadblock. In particular, he tested the possibility that the C12 isotope could be “formed directly in a collision between 3 α-particles”. Production of the C12 isotope could have led to heavier nuclei of nitrogen and oxygen through collisions with protons. Bethe also found that reactions involving the carbon-nitrogen-oxygen group were able to successfully account for the energy production in stars. However, his calculations showed that the 3α   C12 process was highly temperature-dependent and required temperatures of about a billion degrees to “make it as probable” as the proton-proton collisions. Ultimately, Bethe concluded that “there is no way in which nuclei heavier than helium can be produced permanently in the interior of stars under present conditions.” The advent of World War II produced a relative lull in the application of nuclear chemistry to cosmology and stellar astrophysics. But, Bethe’s work had revived interest in the notion of a superdense, superhot early universe that was capable of synthesizing heavier elements. Attempts were made to combine known nuclear reaction rates with the principle of thermal equilibrium (the equivalent of the law of mass action applied to nuclei) to accurately calculate elemental abundances. Most attempts failed to reproduce Goldschmidt’s data and led to various disagreements about the hypothesis of a hot, dense universe. But, theorists such as George Gamow, who had studied under Friedmann in the Soviet Union, continued to apply their knowledge of nuclear physics to explaining the origin of the chemical elements. In 1948, Gamow published a paper with Ralph Alpher and Hans Bethe that was generally considered to be the first significant exposition of a hot Big Bang cosmology (24). Their paper, titled “The Origin of Chemical Elements” was predicated on the two key principles of thermodynamic equilibrium and neutron capture applied to a hot, dense, expanding primordial “neutron gas.”. However, they quickly recognized that: “various nuclear species must have originated not as the result of an equilibrium corresponding to a certain temperature and density, but rather as a consequence of a continuous building-up process arrested by a rapid expansion and cooling of the primordial matter (24).” They postulated that protons formed from neutron decay from the reaction n → p + e- would first form deuterium by capturing a neutron, following which “subsequent neutron captures resulted in the building up of heavier and heavier nuclei.” With assumptions about the time-dependence of matter density, Alpher, who completed his dissertation on this topic, was able to demonstrate 87

agreement with Goldschmidt’s data on the logarithmically decreasing abundances of the elements. Subsequent papers also recognized that the early universe was radiation-dominated with a blackbody spectrum that they calculated to have a current temperature of about 5 K. The incipient big bang model of the universe was born. The αβγ-paper, as it was popularly called, was successful in explaining the hydrogen and helium abundances through primordial processes. But, its failure to account for the mass gaps at atomic weights five and eight remained a serious flaw in the theory. In addition, the theory effectively posited that elemental abundances would remain static over time, an assumption that contradicted Bethe’s calculations showing the chemical evolution of hydrogen in stellar nucleosynthesis. Since the latter was well-founded on known nuclear reactions and spectroscopic data, it undermined the claims of big bang protagonists. These uncertainties and contradictions essentially remained until the appearance of a seminal paper on nucleosynthesis in the year 1957. Fred Hoyle, the eminent British physicist and astronomer, was no fan of the big bang theory. An opponent of the idea that the universe had a beginning, Hoyle, along with his collaborators, proposed an alternative model of the universe (25, 26) in the same year as the αβγ-paper. The steady-state theory posited that the universe was infinitely old and large, static and unchanging. To accommodate the effects of expansion, matter was continually created, primarily in the form of hydrogen. The production of heavier elements was posited to occur solely in stars. After more than a decade of work on stellar nucleosynthesis, Hoyle coauthored a landmark paper with fellow astrophysicists Margaret Burbidge, Geoffrey Burbidge, and William Fowler, that substantially resolved the vexing question of the origin of the elements. The 1957 paper (27), titled Synthesis of the elements in stars (but known more as the B2FH paper after the initials of the authors), began by challenging existing theories on the primordial formation of elements. Including the big bang model, they noted that none of the existing theories could explain the element abundance curves originally published by Goldschmidt (18) and later improved by Hans Seuss and Harold Urey (28). On the contrary, they proposed that “stars are the seat of origin of the elements” related to “the known fact that nuclear transformations are currently taking place inside stars.” The B2FH paper provided a detailed and comprehensive description of stellar processes that produced elements heavier than hydrogen. They found that the temperatures and densities in stellar cores were sufficiently high to produce “progressive conversion of light nuclei into heavier ones as the temperature rises.” Remarkably, they found that the 3α-process that Bethe had earlier discounted was a significant reaction in the buildup of heavier elements from He4. In 1953, Hoyle had predicted the existence of an new, and subsequently discovered, resonant state of the C12-nucleus as a necessity for carbon synthesis in stars. The B2FH authors found that the cross-section for C12-production was dramatically increased due to the presence of this resonance, circumventing the roadblocks at atomic weights five and eight. The paper also gave a plausible explanation of the uneven cosmic distribution of these elements through the ejection of material from stars, including in supernovae explosions. 88

The publication of the B2FH-paper in 1957 and the discovery of the cosmic microwave background radiation in 1965 (29) established Big Bang cosmology as the dominant theory of the creation and evolution of the universe. In the intervening decades, the main contours of element production and processes in the theory had been mapped out. Increasingly precise elemental observations have also produced stringent tests of the predictions of Big Bang cosmology, which the theory has passed with flying colors. A detailed and eminently readable account of the Big Bang can be found in Joseph Silk’s book, The Big Bang (30).

Elemental Abundances and Modern Cosmology A few seconds after the big bang, the universe was dominated by radiation accompanied by a much smaller fraction of ordinary matter, including protons, neutrons, and electrons. The ratio of baryons (protons and neutrons) to photons that was established in the early universe, believed to remain constant over time, is the only free parameter in the standard big bang model. Historically, observed primordial abundances were used to place limits on this value. However, in recent years, it has been independently measured with increasing precision from detailed maps of the cosmic microwave background radiation (CMBR), resulting in a currently accepted value of η = (6.1 ± 0.2) × 10−10 (31). At the onset of nucleosynthesis, the heavier and slow-moving neutrons and protons were in thermal equilibrium, constantly creating and destroying each other through decay reactions with electrons and positrons. However, there was roughly one neutron for every six protons, a consequence of the proton’s lower mass and the temperature of the universe at that time. As the universe cooled and expanded, this interconversion ceased and the ratio “froze out” at 1:7 due to the decay of a small fraction of the remaining neutrons into protons, electrons and anti-neutrinos, with a half-life of about 10 minutes. This ratio, well-established from the standard model in particle physics, is an additional input into theoretical calculations of light element abundances in the standard big bang model. The primordial synthesis of light nuclei was the result of a competition between temperature-dependent nuclear reaction rates and a temperature-lowering expansion of the universe (32). The first element to be created was deuterium (“heavy water”), the result of the capture of a neutron by a proton. Much of the produced deuterium was destroyed, either through disintegration by high-energy radiation or through conversion into tritium, helium-3 and finally helium-4 by neutron absorption. A small fraction, roughly one for every hundred thousand hydrogen nuclei, that survived can be detected today. Since deuterium would not survive production in denser and hotter stellar interiors, any detected abundance must have been produced in the early universe. The calculation is very sensitive to the value of η, with larger values dictating a smaller deuterium abundance due to an increased cross-section for neutron absorption. Thus, deuterium poses a significant test and constraint on big bang theory. The series of neutron capture reactions that began with the deuterium nucleus resulted in the production of helium-4 with its high binding energy. Consequently, ordinary helium is the second most abundant element in the universe, with about 89

one helium nucleus produced in the big bang for every twelve hydrogen nuclei. Over 99%, by weight, of ordinary matter in the universe is made up of hydrogen (about 75%) and helium (about 24%). Though erroneously attributed to stellar interiors, Hans Bethe’s admonition in 1939 that ““no elements heavier than He4 can be built up” became a defining feature of the standard big bang theory. The absence of stable nuclides at atomic mass numbers of 5 and 8 effectively ended primordial nucleosynthesis within minutes after the big bang. There was no way that a helium nucleus could capture a proton or another helium to produce a heavier element. The temperatures and densities in the early universe also did not allow for the alternative nuclear processes suggested by stellar nucleosynthesis. Calculations do predict a small fraction (~10-5) of helium-3 that avoided neutron capture, as well as trace amounts of lithium-7 (fraction~10-5) from the capture of a helium-4 nucleus by tritium. In the past fifty years, one of the significant challenges facing modern astronomy has been spectroscopic observations and determinations of light-element abundances in the universe to test the predictions of the big bang and competing models. Primordial abundance observations are complicated by the ongoing chemical evolution of elements in the universe, particularly in stars, galaxies and the interstellar medium. Spectroscopic signatures of individual atoms, or the molecules they constitute, are typically obtained from many cosmic sources considered to either be good indicators of original abundances, or for which the processes of chemical evolution are well understood (33). In the latter case, the abundances of elements such as oxygen and nitrogen are good indicators of the extent of chemical evolution that has occurred in these sources. Astronomers typically refer to the “metallicity” of gas clouds, stars and galaxies to indicate the fraction of their mass present in elements beyond hydrogen and helium. Consequently, sources with low or zero metallicity are better indicators of primordial element abundances. The presence of deuterium is inferred from the absorption lines of some of the oldest objects in the universe, quasars. Quasars are believed to be the nuclei of active galaxies formed more than ten billion years ago. The absorption of quasar light by highly redshifted hydrogen clouds produces a distinct signature of any deuterium present. Current data suggests that the deuterium to hydrogen ratio is about 27 × 10−8 (33). He4 abundances are typically measured from the emission spectra of highly ionized interstellar gas clouds with low metallicity and show an abundance ratio of 0.24 ± 0.006 relative to hydrogen. The abundance of He3 is more elusive to measure. The spectral signature from He3 mimics that of He4 to a large extent making it difficult to distinguish between the two. Despite this, measurements from ionized gas in our galaxy infer a ratio of about (1.1 ± 0.2) × 10−5 or about one He3 for every 100,000 hydrogen nuclei (32). Data on Li7 are obtained from old, low-mass, low-metallicity stars found in the halos of galaxies. The presumption is that older stars were formed earlier in the history of the universe and contain primordial abundances of elements like Li7 in their outer non-reactive layers. However, theoretical assumptions built into these measurements have produced large uncertainties in the Li7-H ratio of (1.81 − 2.65) ×10−10 (33). 90

Figure 1 shows the comparison between the predictions of big bang nucleosynthesis and observations. As mentioned earlier, the baryon-photon ratio (η) plotted along the horizontal logarithmic-scale axis is an independent parameter into theoretical calculations. The vertical strip indicates the currently accepted value of η = (6.1 ± 0.2) × 10−10 (31). The relative abundances of the elements are plotted along the vertical axis, also a logarithmic scale. The horizontal lines indicate observational values. Overall, it is clear the agreement between data and theory is rather impressive for all expect Li7, which, as noted earlier, has higher uncertainties embedded in the observed value. Extensions or alternatives to the standard big bang model are subject to the strong constraints on elemental abundances reflected in the figure.

Figure 1. Comparison of primordial element abundances predicted by the big bang against observations. Reproduced with permission from reference (32). Copyright 2006 Albert Einstein Institute. 91

Conclusion While our solar neighborhood is dominated by radiation and ordinary matter, the elements of the periodic table constitute less than 5% of the matter density in the known universe. The remaining 95% is postulated to exist in the form of dark matter and dark energy (30), the composition, distribution and impact of which remain among the many open questions in cosmology today. Yet, even as cosmology grapples with new ideas into the nature of the universe, viable theories will always be those that can successfully incorporate its elemental history.

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27. Burbidge, E. M.; Burbidge, G. R.; Fowler, W. A.; Hoyle, F. Rev. Mod. Phys. 1957, 29, 547–650. 28. Suess, H. E.; Urey, H. C. Rev. Mod. Phys. 1956, 28, 53–74. 29. Penzias, A. A.; Wilson, R. W. Astrophys. J. 1965, 142, 419–421. 30. Silk, J. The Big Bang; W. H. Freeman and Company: New York, 2001. 31. NASA/GSFC. Cosmological Parameters Table. https://lambda.gsfc.nasa. gov/product/map/dr5/parameters.cfm (accessed May 25, 2017). 32. Weiss, A. Elements of the past: Big Bang Nucleosynthesis and observation. http://www.einstein-online.info/spotlights/BBN_obs (accessed March 8, 2017). 33. Signore, M.; Puy, D. Eur. Phys. J. C 2009, 59, 117–172.

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Chapter 5

Building Classes of Similar Chemical Elements from Binary Compounds and Their Stoichiometries Guillermo Restrepo* Bioinformatics Group, Department of Computer Science, Leipzig University, Härtelstrasse 16-18, D-04107 Leipzig, Germany Laboratorio de Química Teórica, Facultad de Ciencias Básicas, Universidad de Pamplona, km 1 vía Bucaramanga, 543050 Pamplona, Colombia *E-mail: [email protected]; [email protected].

Similarity is one of the key concepts of the periodic table, which was historically addressed by assessing the resemblance of chemical elements through that of their compounds. A contemporary approach to the similarity among elements is through quantum chemistry, based on the resemblance of the electronic properties of the atoms involved. In spite of having two approaches, the historical one has been almost abandoned and the quantum chemical oversimplified to free atoms, which are of little interest for chemistry. Here we show that a mathematical and computational historical approach yields well-known chemical similarities of chemical elements when studied through binary compounds and their stoichiometries; these similarities are also in agreement with quantum chemistry results for bound atoms. The results come from the analysis of 4,700 binary compounds of 94 chemical elements through the definition of neighbourhoods for every element that were contrasted producing similarity classes. The method detected classes of elements with different patterns on the periodic table, e.g. vertical similarities as in the alkali metals, horizontal ones as in the 4th-row platinum metals and mixed similarities as in the actinoids with some transition metals. We anticipate the methodology here presented to be a starting point for more temporal and even more detailed studies of the periodic table.

© 2017 American Chemical Society

Introduction Chemistry is about substances and especially about compounds. Historically, chemical analysis has led to chemical elements, which are characterised by different properties, being of especial historical importance the atomic weight (1, 2). Through the study of how substances react with each other, it has been possible to find similarities among substances and to relate those similarities classes, which constitute the core of chemistry (3). A particular result of studying compound similarities is that it leads to classifying chemical elements by resemblance (4). The periodic table (PT) by Mendeleev resulted from combining atomic weights and similarity of chemical elements such that atomic weights were taken as the ordering principle and similarities as the source of classes (groups of the PT). Hence, ordering and similarity of an element define its position in the PT. Later on, the atomic number was the accepted ordering criterion, which brought an atomistic ontology for the concept of element: a collection of atoms with the same atomic number (number of protons) (5–7). The exponential growth of chemical substances made difficult assessing similarities through the historical approach (8). This, in combination with the emerging atomistic ontology and the advent of quantum mechanics, particularly its application to chemical elements to understand the structure of the PT, led to model similarities among chemical elements through similarities on the energetic distributions of valence shell electrons (9, 10). This is the cause of the typical contemporary overemphasised textbook introduction to the periodic table through electronic configurations, which normally correspond to those of free atoms (11). However, these configurations are rather dissimilar to those of the bound atoms present in substances, all in all the relevant species for chemistry (11–14). Jørgensen put it plainly as “There is not the slightest doubt, however, that no simple relation exists between the electron configuration of the ground state of the neutral atom and the chemistry of the element under consideration (11).” Besides the textbook oversimplification of the PT to the electronic structure of free atoms, for example in the form of the Madelung rule, authors of these books also overlook a fundamental piece of information of the historical approach, namely that Mendeleev’s studies were mainly based on compounds, particularly on oxides, hydroxides, hydrides and halides and that by studying their similarities he came up with resemblances for the chemical elements (15, 16). Mendeleev highlighted the need to rely on compounds and their proportions of combination rather than on properties of chemical elements; this is evident in his statement that “if CO2 and SO2 are two gases which closely resemble each other both in their physical and chemical properties, the reason of this must be looked for not in an analogy of sulphur and carbon, but in that identity of the type of combination, RX4, which both oxides assume (17, 18).” He adds, “the elements, which are most chemically analogous, are characterized by the fact of their giving compounds of similar form RXn.” Some few contemporary authors have stressed the importance of compounds and have elaborated on the related problem of quantifying qualitative parts of chemistry in similarity studies (19). Following Mendeleev’s ideas, a key concept 96

to understand similarities of chemical elements is that of valency, which is obtained by stoichiometric decomposition of compounds in chemical analysis (20). Hence, relying on chemical compounds and their stoichiometries, we discuss in the current chapter results of a mathematical and computational study of similarity of chemical elements (4).

Materials and Methods As a first approach to chemical classification of elements through compounds, binary substances were analysed (4). As usual, the binary compound was defined as any substance containing two different elements, independently of their proportions. Hence, H2O, KCl and CH4 entered in the study, but H2SO4 or C60 did not. In total, 4,700 binary compounds were analysed, which accounted for 94 chemical elements with at least one binary compound reported in the literature. As the binary compound exists, its decomposition reaction to chemical elements also exists (Figure 1A). The set of decomposition reactions gives place to a network, a hypergraph, as exemplified in Figure 1B, which, for our purposes, can be reduced to a subnetwork (Figure 1C) whose vertices are the elements produced by decomposition (21). The edges of this subnetwork, or lines between two vertices, are pairs of elements produced by a decomposition reaction. We call this network the product-network.

Figure 1. A) Two decomposition reactions ρ and ρ′. B) Modelling of ρ and ρ′ as a hypergraph. C) Product-network of B.

As the interest is analysing the similarity among chemical elements, the structural similarity of the chemical elements in the product-network is to be explored in such a way that elements linked to common elements are similar. Thus, for example, Na and K are similar for they form fluorides, chlorides and bromides, to name but a few of their common binary compounds. This is shown in Figure 2A and 2B, where Figure 2B shows that Na and K are linked to F, Cl and Br, therefore Na and K are similar; F, Cl and Br are also similar, for they are linked to Na and K. To quantify such a similarity, the neighbourhood of each element was determined, such that it contains the element in question and the elements that are connected to it in the product-network. For the three elements of Figure 1, the neighbourhoods are Nx= {x,y}, Ny= {x,y,z} and Nz= {y,z}. The corresponding neighbourhoods for the elements of Figure 2 are shown in Figure 2C. 97

The more similar the neighbourhoods, the more similar the elements are; or the more different the neighbourhoods, the more different the elements. These differences can be calculated by counting the number of elements in the symmetric difference between pairs of neighbourhoods. The symmetric difference of two sets results from their union and the removal of their intersection. The number of elements of the corresponding symmetric differences of Figure 2 are shown in Figure 2D, where, for example, the two elements of the symmetric difference between NK and NNa are Na and K. Figure 2 shows that Na and K are similar and F, Cl and Br also form another similarity class.

Figure 2. A) Hypergraph of reactions ρ1 to ρ6, its product-network (B) and the neighbourhoods of each element (C). D) Number of elements of the symmetric difference between pairs of neighbourhoods.

Now, let us suppose that we have the hypergraph shown in Figure 3A, whose product-network is shown in Figure 3B along with the respective neighbourhoods of the elements (Figure 3C). This network indicates that H is evenly similar to F and B. But F only forms a binary compound with H while B forms by far more with H. A method to take into account this diversity of combinations, following Mendeleev’s ideas, is to consider the stoichiometry of the combinations as follows: for a compound xayb, the neighbourhood of x is given by {xa/b,yb/a}, in this way the stoichiometry of the binary combination is attached to the respective element. 98

Hence, the neighbourhoods for the elements in Figure 3 are extended to NF= {F1/1,H1/1}, NH= {F1/1,H1/1,H6/2,H9/5,H14/10,…,B2/6,B5/9,B10/14,…}, NB= { H6/2,H9/5,H14/10,…,B2/6,B5/9,B10/14,…}, from which it can be concluded that there are more resemblance between H and B than between F and H, for there are by far more commonalities between NH and NB than the only two (F1/1,H1/1) between NF and NH.

Figure 3. A) Hypergraph of some decomposition reactions of binary compounds with H and either F or B. B) Product-network of A. C) Neighbourhoods of each element in B.

These stoichiometric neighbourhoods were determined and their differences quantified. The elements were clustered according to their neighbourhood resemblance using hierarchical cluster analysis (HCA) with the average union as grouping methodology. An illustrative example of the HCA algorithm is shown in Figure 4, which starts with the differences for the neighbourhoods shown in Figure 4A. It is found that A and B are the most similar elements; then, after grouping A with B new differences are recalculated (Figure 4B). As the grouping methodology is the average union, the difference between {A,B} and D, for example, is given by the average of the differences between A and D and that of B and D, i.e. 4 (22). With this new table of differences (difference matrix), the algorithm keeps running (Figures 4C, 4D) until all elements are grouped. A depiction of these groupings is a dendrogram (Figure 4E), which shows that A and B are the most similar elements and that they, in turn, are similar to element C. This group of three elements is similar to D and, finally, the most dissimilar element to the group of four elements is E. Another depiction of the similarities takes into account the different levels of similarity of the merging process, which are also seen in the hierarchical structure of a dendrogram. Figure 4F shows such a similarity landscape. 99

Figure 4. A) Number of different elements between Ni and Nj. B) New difference matrix after merging A and B, C) {A,B} with C and D) {A,B,C} with D. E) Dendrogram and F) its similarity landscape.

Results and Discussion The 94 elements explored through the 4,700 binary compounds are shown in Figure 5 making use of the conventional medium-long form of the PT (9).

Figure 5. 94 elements explored through binary compounds. The similarity landscape for the chemical elements is shown in Figure 6 and the dendrogram is found in reference (4). 100

Figure 6. Similarity landscape of 94 chemical elements.

Figure 6 shows that the most different element is H, i.e. there is no other element whose presence and proportion in binary compounds is similar to that of H. Other similar cases are found for B, C, N and O, which are evidences of the singularity principle, i.e. the chemistry of the second period elements is often different to the latter members of their respective groups (23, 24). In fact, elements on the red regions of Figure 6 are not only different from the elements of their groups but entirely different from all the other elements. This indicates that no other element forms binary compounds with the elements they combine and with the stoichiometries they have. Still on the red regions, halogens show up, right between the singularities of H, C and O and those of S, B, P and N. This indicates that halogens, as a group of the PT, have no other element behaving as they do with the elements they form binary compounds. However, halogens are very similar among themselves, as they belong to a blue region. The strongest similarity within halogens occurs between Cl and Br; a bit different is I and, as an evidence of the singularity principle, F is the least similar halogen. By increasing the similarity, Se and Te form a very similar couple and not far from them As and Sb result also similar. In fact, of the chalcogens, only Se and Te are similar, for all other members of this group show up as single classes. Likewise pnictogens behave, where only As and Sb are similar. Right in the middle of {Se,Te} and {As,Sb} Si is found; which shows that some elements which it forms binary compounds with are elements which Se and Te, and As and Sb also form with and with similar stoichiometries. 101

At a higher level of similarity, alkali metals show up constituting a similarity “bridge” between the already discussed elements, especially {As,Sb}, Si, {Se,Te}, the single pnictogen Bi and the transition metals; especially Cr and Mn. Within the very strong similarities of alkali metals, the couples {Li,Na} and {Rb,Cs} are the most similar classes. K is most similar to Rb and Cs than to Li and Na. Bi, Cr and Mn can also be regarded as elements connecting the red-orange regions of markedly differences with the green-blue regions, where most of the similar elements are gathered. At the periphery of this regions belong the remaining transition metals, forming several clusters of very strong similarities. It is found that the V-group is one of these clusters, as well as the Zn-group, this latter interestingly sharing similarities with alkaline earth metals and with {Ge,Sn,Pb}. These resemblances of the Zn-group elements indicate that they form binary compounds with several of the elements which the alkaline earth metals do with and also that the Zn-group elements combine in a similar fashion as Ge, Sn and Pb do. Hence, the Zn-group can be regarded as an intermediate group between alkaline earth metals and Ge, Sn and Pb. Another interesting transition metal cluster is {Fe,Co,Ni,Pd}, whose elements are part of the so called “platinum metals,” group VIII in the old IUPAC group numbering or VIIIB in the CAS numbering. This shows that they indeed have commonalities regarding the elements which they combine with. This reminds us that the similarity of these “platinum metals” was what led Mendeleev to group them together as noted in his claim that “Only among these metals are compounds of the type RO4 or R2O8 formed (which is why they are designated as the eighth group) (18, 25).” Although these similarities, Ru, Os, Rh, Ir and Pt are not so similar to Fe, Co, Ni and Pd. Rh and Ir, in turn, are more similar to lanthanoids and actinioids. All elements of Ti-group are similar, especially Zr and Hf, but they are also similar to two actinoids: Th and U. The similarity of Th and {Zr,Hf} was already in 1945 highlighted by Seaborg (26). Schwarz has also recently discussed the resemblance of early actinoids with some 6th-row transition metals (27). In particular, the resemblance of Zr and Hf was recognised by Goldschmidt as the effect of the lanthanoid contraction, which in modern terms is understood as a spatial shrinking of lanthanoids species as the result of the filling of 4f shells that brings a contraction of the 5p and 6s shells of the species. Such a contraction makes that, e.g. Zr(IV) and Hf(IV) have almost equal ionic radii when six-coordinated (28–30). Although the Zr-Hf similarity is well documented in the literature, it is striking to find that it is considered an exception, “due to an anomalous cancellation of relativistic effects,” of the differences between 5th- and 6th-row elements of the same group (30, 31). In the current study we found that out of the 17 possible pairs of 5th- and 6th-row elements belonging to a group, there are other five pairs sharing similarities: {Nb,Ta}, {Mo,W}, {Tc,Re}, {Ru,Os} and {Rh,Ir}. The first two were also discussed by Huheey and Huheey, who found as a cause of it the almost equal radii for 5th- and 6th-row species (29). Fricke et al. also discussed the generality of this resemblance and the similar oxidation states for 5th- and 6th-row elements (32). 102

Another cluster, which involves main group elements, is {Al,Ga,In,Tl,Cu,Ag,Au}, where the Al-group is merged with the coinage metals. By increasing the similarity level is found that the Al-group is a subcluster and {Cu,Ag} another one. It is observed that the similarity of the Al-group with the Cu-group is given by the resemblance between Tl and Au, which may be caused by relativistic effects. The largest cluster gathering similar elements is the one of lanthanoids and actinoids. Lanthanoids are more similar among them than actinoids among them. Jørgensen states that such a strong resemblance among lanthanoids, characterised by a constant dominant oxidation state III, was initially noticed by Rydberg in 1914 when mentioning the “existence of fifteen consecutive elements having almost the same chemical properties (11, 27).” Fricke et al. gave an explanation by relating ionisation energy with valency: “The ionization energy […] is a valuable quantity for theoretical predictions of the valency of the elements. […] the ionization energy becomes larger if the wave function of a new additional electron is not screened by the other electrons with the same wave function. This altogether means that the ionization energy curve for the lanthanides is very flat and, therefore, the valency is always nearly the same (32).” And they explained the differences among actinoids “because the 5f electrons are not so buried in the atom as the 4f electrons in the lanthanides and thus the valency at the beginning is larger whereas it becomes smaller at the end.” Within lanthanoids, Ce is the most different element and there is a strong resemblance of lanthanoids with Sc and Y, a similarity already found by Goldschmidt, especially for Y and Dy and Ho (28). In fact, Sc, Y and the lanthanoids constitute the so-called rare earths, which are grouped together given their chemical similarities that have found support on quantum chemical grounds (33). Scerri has discussed about the element at the beginning of the third row of the transition elements, which in some tables is La and in others Lu (9). Schwarz and Rich have stated that Lu cannot be considered a lanthanoid, for it does not fill f orbitals as they are already filled; and have suggested that Lu should be regarded as a transition metal (33). According to our results, La appears in between two clusters, one of 11 lanthanoids and another of transition metals, namely {Y,Sc}. Lu is part of the clusters of 11 lanthanoids and the smallest cluster containing it is {Ho,Er,Lu}, which shows that Lu is more similar to lanthanoids than to transition metals, while La share similarities with lanthanoids and with transition metals. Therefore La must be the element located at the beginning of the third row of transition metals if chemical resemblances is what it is to be emphasized. Besides the similarity of lanthanoids with Sc and Y, lanthanoids are also similar to some actinoids, as found in the presence of Tb and Pr in a large cluster of actinoids and noble gases. A resemblance that has been discussed by Schwarz on quantum chemical grounds (27). However, the similarities are more notorious among members of the same row for lanthanoids and actinoids, as also reported by Diwu et al. and Cary et al. (34, 35) The relaxed similarities among actinoids can be seen for example in their dominant oxidation states, which vary from II in No to VI in U; variations related to 103

the actinoid contraction, which is more irregular and not as large as the lanthanoid one (27, 32, 36). If lanthanoids have similarities with Sc and Y, actinoids have them with more transition metals, e.g. with Zr, Hf, Tc and Re. It is found, for example, that U is similar to the Ti-group of transition metals and to Th. Similarities of actinoids with transition metals have been reported by Rayner-Canham and studied by Schwarz and Rich (23, 33). The largest set of similar actinoids is Am, Cm, Bk, Cf, Es. An interesting place for Pu is found in the similarity landscape, which results alike to some actinoids {Cm,Bk,Es,Am,Cf,Ac} and to the lanthanoids Tb and Pr. Chemical singularities of Pu have been recently studied by Schwarz in a systematic study of tricarbonato-actinyl anions (27). Cary et al. argue that the special behaviour of Pu is given by its unique electronic properties, which come from the changing roles of the 5f orbitals, which among other features, allow it to equilibrate four oxidation states in solution, something not reported for any other chemical element (35). Scerri has also made the point about the element at the beginning of the fourth row of the transition elements, being Ac in some tables and Lr in others (9). Unfortunately no relevant data was available for Lr at the time of the study, therefore it is not possible to discuss its resemblance to other elements (4, 37). Ac is found in a single class, which is also similar to a class of elements with very few number of binary compounds, i.e. Ra, Kr, Xe and At (38). In general, Figure 6 shows that the most different elements are: most of the second row elements, H, halogens and alkali metals; whose trends of combination are unique and which are very different from those of transition metals, lanthanoids and actinoids. These three later sets are, in general, alike regarding the elements which they form binary compounds with and also alike in the proportion they do it.

Conclusions and Outlook Although there have been several studies on the similarity of chemical elements through different chemical, physical and physicochemical properties of the elements; the study reported in Leal et al., here further explained, is the only one where the recovery of trends in the PT has been attained by uniquely using chemical information and where more elements have been regarded (4, 39–42). This study follows Mendeleev’s ideas of devising similarities for the elements based on their compounds through a mathematical and computational approach. It is found that the method is able to detect several well-known classes of similar elements with different patterns on the PT, e.g. vertical similarities as in alkali metals, halogens, Al-group and Cu-group; horizontal ones as in 4th-row platinum metals, lanthanoids, actinoids; and some other mixed patterns as in lanthanoids and Sc and Y (rare earths); and actinoids with some transition metals. It is found that for most of these similarities, there is a quantum chemical possible explanation based on non-simplistic rules of electronic configuration fillings but on studies that take into account, e.g. relativistic effects, for heavy atoms and also the bound character of atoms in compounds. 104

The approach followed in the current chapter constitutes also an alternative way to introducing the PT to students, with a more chemical “flavour” than the traditional oversimplification relying on electronic configurations of free atoms. The ingredients of the approach here presented are fundamental concepts of chemistry such as compound, reaction and stoichiometry, which make part of the bulk level by Nelson for describing chemistry (43). However, in teaching, the concepts must be presented as simple as possible and the current approach is not that simple. A work to do to reduce the complexity of this presentation is to look for those particular regions of the neighbourhoods defined for the elements, which make that the results do not vary too much. This would also shed some light on the most relevant neighbours for keeping the similarity structure of the PT. The similarity classes found match to a big extent the chemical way of presenting the elements in some few specialised chemical books like Chemistry of the Elements, which is to be expected from books rooted in chemical reactions (44). The study here discussed was based upon binary compounds, but the current amount of chemical information, stored in electronic libraries, and the current computing capacity, allow thinking in running the study with ternary, quaternary and even with the whole set of chemical substances. An interesting example of exploration of large networks of chemicals are the studies of Grzybowski and his team (45–47). The study can also be run at different time periods to explore how the patterns on the periodic table have changed in time. At first glance, it looks like the approach to chemical similarity here discussed cannot stand the test of time, for it relies on compounds, which are especially scarce for the heavy elements. Moreover, for these elements the few compounds that are obtained are synthesised in a one-atom-at-a-time fashion, which is rather different to the bulk one, of the traditional chemistry (48, 49). This brings not only a clash of chemical traditions, but also the mixture of two different ontological levels for compounds, i.e. the bulk and the one of atomic aggregates. But the method overcomes these problems, for it is actually based, more than upon compounds, on their mathematically generality, i.e. their composition and stoichiometry. Both can be extracted from either bulk or atomic aggregate compounds; wet-lab synthesized or in a one-atom-at-a-time way or even estimated through quantum chemical approaches (50). From a methodological viewpoint there were two important steps for finding similarity classes, namely the construction of neighbourhoods and their clustering. In Bernal et al. it has been shown that those neighbourhoods can be further explored through an interesting connection with Formal Concept Analysis, a mathematical technique based on the finding of closed sets of objects (51). There, objects are characterised by their attributes and, for the particular case of binary compounds, objects are chemical elements and their attributes are those chemical elements belonging to their neighbourhoods. The use of neighbourhoods for chemical elements brings also possibilities to explore in more detail the landscape of similarities through the inclusion of basic elements of topology (41, 42, 52). Works on this direction have started with Restrepo et al. and have also found a connection with the results obtained by Formal Concept Analysis (41, 51). The richness of this subject is still to be 105

explored and constitutes an interesting field of study for mathematical chemistry (53). Regarding the clustering of elements, the algorithm used was the hierarchical cluster analysis, one of the most used classificatory techniques, particularly in chemistry (54). However, it has limitations, e.g. the stability of clusters is reduced given equidistances in the difference matrices during the merging processes. This is a problem known as ties in proximity, which was brought to the attention of the chemical community by MacCuish and which has been further explored by Leal et al. (54, 55) Leal and his team developed measures of cluster’s frequency taking into account ties; therefore it is worth exploring how frequent and stable are clusters of chemical elements given ties using these techniques. A related issue of hierarchical cluster analysis is that clusters may vary if the grouping methodology changes. The work to do is to explore if clusters obtained in this work remain stable after using several other grouping methodologies. However, previous results on chemical elements, although not using relational properties as in the current work, have shown that clusters are in general stable, especially for those around noble gases, e.g. alkali metals and halogens (41, 42). Finally, from a philosophical perspective, Schädel has stated the current, in our opinion counterhistorical, understanding of the PT, whose premise is that the atomic number and the electronic configuration of an element determine its position in the PT (49). This brings the consequence that from the electronic configuration of the position, the chemical properties arise, which leads to claim that chemical properties can be linked to trends of the electronic configurations along groups or periods. In our opinion, the premise is that similarity in chemical properties of compounds, along with atomic number lead to the position of the elements in the PT (3). This brings the consequence that from similarities in chemical properties, electronic configurations are obtained. Fricke et al. clearly stated it for the case of superheavy elements: “Theoretical predictions of the chemistry of superheavy elements can be done in two ways (32). First the behaviour of the well-known elements as a function of their chemical group and period can be extrapolated into the unknown regions. Secondly eigenvalues, wavefunctions, most stable configurations etc. can be calculated theoretically so that on this basis the predictions of the chemistry can be done in a more accurate way.” Given the stability of our approach, Schädel’s question whether the PT is still valid regarding chemical properties of the superheavy elements, from Rf onwards, has a positive answer (49).

Acknowledgments G. R. thanks the Universidad de Pamplona and the Alexander von Humboldt Foundation/Stiftung for funding this research. Eugen Schwarz is specially thanked for his valuable comments on early versions of this document and for pointing out important aspects of the quantum chemistry of heavy elements.

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13. In reference 11 Jørgensen discusses how “the very high ionization energies of atomic spectroscopy are to a large extent compensated by the electrostatic attraction of the surrounding anions” for crystals. 14. Wang, S-G.; Schwarz, W. H. E. Icon of chemistry: The periodic system of chemical elements in the new century. Angew. Chem., Int. Ed. 2009, 48, 3404–3415. 15. Schwarz, W. H. E.Which brings chemically relevant results for about 20% of all elements. Personal Communication. See reference 14. 16. Schwarz, W. H. E.; Wang, S-G. Some solved problems of the periodic system of chemical elements. Int. J. Quantum Chem. 2010, 110, 1455–1465. 17. Mendeleev, D. Principles of Chemistry; Longmans, Green & Co: London, 1905; Chapter 15. 18. Jensen, W. B. Mendeleev on the Periodic Law, Selected Writings, 1869-1905; Dover: Mineola, NY, 2005. 19. Schwarz, W. H. E. Recommended questions on the road towards a scientific explanation of the periodic system of chemical elements with the help of the concepts of quantum physics. Found. Chem. 2007, 9, 139–188. 20. Schwarz, W. H. E. Towards a physical explanation of the periodic table (PT) of chemical elements, achievements of the previous generations. In Fundamental world of quantum chemistry; Brändas, E. J., Kryachko, E. S., Eds.; Springer: Dordrecht, The Netherlands, 2004; Vol. III, pp 645−669. 21. Klamt, S.; Haus, U-U.; Theis, F. Hypergraphs and cellular networks. PLoS Comput. Biol. 2009, 5, e1000385. 22. MacCuish, J.; MacCuish, N. E. Clustering in Bioinformatics and Drug Discovery; Mathematical and Computational Biology Series; CRC Press: Boca Ratón, FL, 2011; pp 1−244. 23. Rayner–Canham, G. Periodic patterns. J. Chem. Educ. 2000, 77, 1053–1056. 24. An atomic interpretation of this effect comes from the small atomic radius of 2nd-row species regarding the bigger size of atoms in higher rows. 25. Mendeleev, D. The periodic law of the chemical elements. J. Chem. Soc. 1889, 55, 634–656. 26. Seaborg, G. T. The chemical and radioactive properties of the heavy elements. Chem. Eng. News. 1945, 23, 2190–2193. 27. Liu, J-B.; Chen, G. P.; Huang, W.; Clark, D. L.; Schwarz, W. H. E.; Li, J. Bonding trends across the series of tricarbonato-actinyl anions [(AnO2)(CO3)3]4- (An = U−Cm): The plutonium turn. Dalton Trans. 2017, 46, 2542. 28. Wedepohl, K. H. The importance of the pioneering work by V. M. Goldschmidt for modern geochemistry. Naturwissenschaften 1996, 83, 165–171. 29. Huheey, J. E.; Huheey, C. L. Anomalous properties of elements that follow “long periods” of elements. J. Chem. Educ. 1972, 49, 227–230. 30. Pyykkö, P.; Desclaux, J-P. Relativity and the periodic system of elements. Acc. Chem. Res. 1979, 12, 276–281. 31. Pyykkö, P. The physics behind chemistry and the periodic table. Chem. Rev. 2012, 112, 371–384. 108

32. Fricke, B.; Greiner, W.; Waber, J. T. The continuation of the periodic table up to Z = 172. The chemistry of superheavy elements. Theoret. Chim. Acta. 1971, 21, 235–260. 33. Schwarz, W. H. E.; Rich, R. L. Theoretical basis and correct explanation of the periodic system: Review and update. J. Chem. Educ. 2010, 87, 435–443. 34. Diwu, J.; Grant, D. J.; Wang, S.; Gagliardi, L.; Albrecht-Schmitt, T. E. Periodic trends in lanthanide and actinide phosphonates: Discontinuity between plutonium and americium. Inorg. Chem. 2012, 51, 6906–6915. 35. Cary, S. K.; Vasiliu, M.; Baumbach, R. E.; Stritzinger, J. T.; Green, T. D.; Diefenbach, K.; Cross, J. N.; Knappenberger, K. L.; Liu, G.; Silver, M. A.; DePrince, A. E.; Polinski, M. J.; Van Cleve, S. M.; House, J. H.; Kikugawa, N.; Gallagher, A.; Arico, A. A.; Dixon, D. A.; Albrecht-Schmitt, T. E. Emergence of californium as the second transitional element in the actinide series. Nat. Commun. 2015, 6, 6827. 36. Given these variations in actinoids’s similarity, Jørgensen11 suggested not referring to them as “actinides”, or actinoids, in more contemporary terms, but to 5f group elements; for, as pointed out by Schwarz and Rich,33 actinide/ actinoid means “like actinium.” 37. A recent revision of the literature shows that there are at least three binary compounds for including it in a new study. 38. For all these elements the current literature has more binary compounds, allowing running a finer study for them. 39. Zhou, X.-Z.; Wei, K.-H.; Chen, G.-Q.; Fan, Z.-X.; Zhan, J.-J. Fuzzy cluster analysis of chemical elements. Jisuanji Yu Yingyong Huaxue. 2000, 17, 167–168. 40. Sneath, P. H. A. Numerical classification of the chemical elements and its relation to the periodic system. Found. Chem. 2000, 2, 237–263. 41. Restrepo, G.; Mesa, H.; Llanos, E. J.; Villaveces, J. L. Topological study of the periodic system. J. Chem. Inf. Comput. Sci. 2004, 44, 68–75. 42. Restrepo, G.; Llanos, E. J.; Mesa, H. Topological space of the chemical elements and its properties. J. Math. Chem. 2006, 39, 401–416. 43. Nelson, P. G. Teaching chemistry progressively: From substances, to atoms and molecules, to electrons and nuclei. Chem. Educ. Res. Pract. 2002, 3, 215–228. 44. Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Elsevier: Oxford, UK, 2005; pp v−xviii. 45. Fialkowski, M.; Bishop, K. J. M.; Chubukov, V. A.; Campbell, C. J.; Grzybowski, B. A. Architecture and evolution of organic chemistry. Angew. Chem., Int. Ed. 2005, 44, 7263–7269. 46. Bishop, K. J. M.; Klajn, R.; Grzybowski, B. A. The core and most useful molecules in organic chemistry. Angew. Chem., Int. Ed. 2006, 45, 5348–5354. 47. Grzybowski, B. A.; Bishop, K. J. M.; Kowalczyk, B.; Wilmer, C. E. The ‘wired’ universe of organic chemistry. Nat. Chem. 2009, 1, 31–36. 48. This difference is not only a matter of scale and of amount of substance, but also of the synthetic machinery and of its attached theory. According to Schädel:49 “As a single atom cannot exist in different chemical forms 109

49. 50.

51.

52. 53.

54.

55.

taking part in the chemical equilibrium at the same time, the classical law of mass action—well established for macroscopic quantities and characterizing a dynamic, reversible process in which reactants and products are continuously transformed into each other—is no longer valid. For single atoms, the concept of chemical equilibrium needs to be substituted by an equivalent expression in which concentrations, activities, or partial pressures are replaced by probabilities of finding the atom in one state or the other”. Besides formation of compounds, other properties studied at this level are volatilities, formation of complexes in aqueous solutions and their interaction with other phases.49 Interestingly, properties such as ionic radius and the stability of oxidation states are indirectly obtained by using the PT, for they result from the comparison with the known properties of lighter members of the group the element belongs to.49 Schädel, M. Chemistry of superheavy elements. Angew. Chem., Int. Ed. 2006, 45, 368–401. Pyykkö, P. A suggested periodic table up to Z ≤ 172, based on Dirac-Fock calculations on atoms and ions. Phys. Chem. Chem. Phys. 2011, 13, 161–168. Bernal, A.; Llanos, E. J.; Leal, W.; Restrepo, G. Similarity in chemical reaction networks: Categories, concepts and closures. In Advances in mathematical chemistry and applications; Basak, S. C., Restrepo, G., Villaveces, J. L., Eds.; Bentham: Sharjah, United Arab Emirates, 2015; Chapter 2, pp 24−54. Restrepo, G.; Mesa, H. Chemotopology: beyond neighbourhoods. Curr. Comput-Aided Drug Des. 2011, 7, 90–97. Restrepo, G. Mathematical chemistry, a new discipline. In Essays in the Philosophy of Chemistry; Scerri, E., Fisher, G., Eds.; Oxford University Press: New York, 2016; Chapter 15, pp 332−351. Leal, W.; Llanos, E. J.; Restrepo, G.; Suárez, C. F.; Patarroyo, M. E. How frequently do clusters occur in hierarchical cluster analysis?: A graph theoretical approach to studying ties in proximity. J. Cheminf. 2016, 8, 4. MacCuish, J.; Nicolaou, C.; MacCuish, N. E. Ties in proximity and clustering compounds. J. Chem. Inf. Comput. Sci. 2001, 41, 134–146.

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Chapter 6

It’s All in the Sludge: Elements That Are Always By-Products Justin Pothoof, Grace Nguyen, and Mark A. Benvenuto* Department of Chemistry and Biochemistry, University of Detroit Mercy, 4001 W. McNichols Road, Detroit, Michigan 48221-3038, United States *E-mail: [email protected].

The periodic table has been taught predominantly as an exercise in memorization for decades, with connections to reactivity, but seldom to abundance of elements. As society changes and technological developments are brought to the market and general public, the need for formerly unused or under-used elements may increase drastically, and in some cases already has. An understanding of the origins and sources of the elements thus becomes an important component in educating students – the next generation of the adult work force – about the periodic table.

Introduction Is every element on the Periodic Table of the Elements of equal importance, and present in or on the Earth in equal abundance? Are the elements all extracted from their natural sources in the same manner, or are some simply part of the “sludge” of the production of other elements? Unfortunately, the table we use and depend on so heavily gives no indication. To any student new to the study of chemistry, the idea of the Periodic Table of the Elements is that this is a large document (often a wall chart) that simply lists all the elements known, in an apparently random order, using one or two apparently random letters, and with the implied idea that since all of them are given equal billing, all of them are of roughly equal importance. Students are eventually taught the correlation between such ideas as atomic number and number of protons and electrons, but at first glance, the table seems to have no order. Element names certainly have no system to them, and even the shape of the table does not appear © 2017 American Chemical Society

to have any obvious reason. “Learning” the Periodic Table is essentially a large exercise in memorization. Our current image of the Periodic Table of the Elements is roughly 70 years old, and its shape is connected with the work of Professor Glenn Seaborg and others, who were interested in completing the transuranic elements and the seventh row of the Table through the f-block electrons. There are certainly similarities to this now established work and the much more recent work in which elements up to and including Element 118 – Oganneson – have been synthesized and verified, thus completing the seventh row of the Table through the p-block elements (1). And while such work extends our knowledge of the periodic table, it does not imbue any deeper understanding of either the prevalence of any elements, or their usefulness in society.

Abundance, Prevalence, and Uses of the Elements A thought exercise that can be conducted even in a freshmen-level chemistry class is a series of questions concerning the abundance of all matter. Table 1 lists the questions, as well as general answers, that focus the experiment on the abundance of a few, heavier elements.

Table 1. The Distribution of Matter Question

General Answer

Of what is almost all of the universe composed?

>>>99% is empty space, or nothing.

Of what is almost all of the remaining part of the universe, the matter of the universe, made?

>>>99% are the stars.

What are stars made of?

Hydrogen and helium

Where is almost all matter that is not stellar?

The planets.

Of the planets, where is most of the matter, and most of the light elements?

The gas giant planets.

How many planets do we know of that are composed of elements from hydrogen to uranium?

One, Earth.

What heavier, metallic elements are thus actually present in relatively small amounts, but are heavily used by humanity?

Iron, copper, aluminum (there are others as well).

The exercise makes it apparent how little of the matter in the known universe is actually any element other than hydrogen or helium. While this is admittedly an exercise whose scope is greater than what is usually undertaken in most chemistry 112

classes, it does give a person some feel for how vast the scale is with which we can consider all matter, from hydrogen to the heaviest of elements. These questions can be stated in a number of different ways, and could go even further than what is presented here. But the large idea behind this is to focus student thinking on how elemental abundances exist, and how much or little of any of them exist in forms that are available to us. A series of questions such as those in Table 1 provide some focus and understanding of the Periodic Table of the Elements beyond mere proton counts. But a look at the chemical industry and what elements are used in large amounts represents another view of the table. For example, Table 2 shows several elements that are used on a very large scale, to produce relatively simple chemical compounds that in turn enable the quality of life enjoyed by much of the world’s people today.

Table 2. Products Made from the Elements (2) Element

Processing

Product

Example uses

Sulfur

Addition to O2 and water

Sulfuric acid

Fertilizer production

Nitrogen

Addition of elemental hydrogen

Ammonia

Fertilizer

Nitrogen

Addition to oxygen

Nitric acid

Oxidizers

Copper

Refined from ores

Elemental copper

Wiring (among many)

Iron

Refined from ores

Elemental iron

Iron and steel objects

Aluminum

Refined from bauxite

Elemental aluminum

Lightweight alloys

Chlorine

Electrolytic, of salt water

Sodium hydroxide, elemental chlorine, elemental hydrogen

Chlorine disinfectant, polyvinylchloride plastics

These six elements are produced and isolated at a level of millions of tons per year, and Table 2’s listing of example uses is always a very small fraction of a much large whole. Some, such as iron and copper, have been used for millennia, and have seen incredible improvements in production over that time. Others, such as aluminum, have a much shorter history from discovery to the present day, and thus have seen relatively rapid improvements in their large scale production. But approximately twenty elements are never mined, extracted, or in some way isolated as the main product of an operation, or are produced largely as secondary products to some other refining process. This was not necessarily a problem in the past, but with the changing needs of society, and the expanded uses of novel materials, their recovery, extraction and isolation becomes very important today. 113

As well, increased needs for certain elements means that even if they were a major product in some past method of production, they are now valuable enough that they are pursued even as secondary products. Silver, gold, and the platinum group metals (PGM) are all examples of this. Table 3 gives a listing of these elements.

Table 3. Elements Co-Produced with Other Materials (2, 3) Element

Source

Co-produced with:

Example use

Bromine

Brines in Arkansas, and the Dead Sea

Common salt, magnesium

Flame retardants, pesticides

Cerium

Monazite and bastnäsite

Lanthanides

Polishing compounds

Gadolinium

Monazite ores

Other rare earth elements

High resistance alloys

Gallium

Bauxite

Aluminum

Semi-conductors

Germanium

Copper, lead, zinc ores

Copper, lead, or zinc

Fibre-optics

Gold

Copper ores

Copper, as part of anode muds

Electrical connections, store of wealth

Lithium

Brine pools

Salt

Ceramics, batteries

Neodymium

Monazite and bastnäsite

Other rare earth elements

Magnets

Platinum group metals

Anode mud or anode slime

Copper

Catalysts

Scandium

Apatite

Phosphorus

Lightweight alloys

Selenium

Copper, nickel, lead ores

Copper, nickel, or lead

Magnesium production

Silver

Lead ores

Lead, via the Parkes Process

Electronics, coins

Thorium

Monazite ores

Rare earth elements

Mag-thor alloys

It can be seen that several of the elements in Table 3 are ultimately extracted from monazite or bastnäsite ores. They are routinely co-located because of the similarities in chemical reactivity of many of the lanthanides and of thorium. But prior to the Second World War, there were few needs for these elements, and thus the isolation and refining of them – a tedious process requiring that steps be repeated many times – was not considered important, or worth improving. The current profile of needs and applications for elements such as cerium, neodymium and gadolinium are such that some greater understanding of these materials is necessary; logically this can be first taught to new, undergraduate students. 114

In similar vein, the production of gallium is always a subordinate operation in the production of aluminum from bauxite. Yet the rise of a large, now mature semiconductor industry means that there is a steady need for this by-product element. As a third broad example, importantly, the electro-refining of copper on a large scale has meant the co-production of what are called anode muds or anode slimes. These waste materials, never refined out of copper before the advent of electrochemical methods of purification, are valuable enough co-products that they can sometimes become the driving force for the larger production of copper. If, for instance, copper sells on the world’s markets at $2.50 per pound, a 300-lb refined copper ingot has a base price of $750. If one ounce of gold is an anode mud co-product of a particular copper batch, and sells at $1,200 per troy ounce (a price at which is hovered for much of the year 2016), the co-product now generates more income than the main product. This becomes even more apparent if some amount of silver or the platinum group metals are also present in the anode mud.

Mendeleev’s “Eka” Elements and Their Uses Professor Dimitri Mendeleev is often credited as being the father of the Periodic Table of the Elements because of his ability to predict elements that had not yet been found, inclusive of what some of their chemical and physical properties should be, as well as the formulas of those compounds. Four undiscovered elements that he labeled with the term “eka” include eka-boron (scandium), eka-aluminum (gallium), eka-manganese (technetium), and eka silicon (germanium). At the time of Mendeleev’s publication of “The Principles of Chemistry,” technetium had still not been discovered and isolated, and in his table he notes that manganese forms compounds that include a formula of MnO3Z, and that an element he calls an “unknown metal” designated as “Em=99” should combine to form compounds of the formula EmO3Z (4). While there are applications for all four of these eka elements today, and thus easy connections to them to be made when teaching, technetium is arguably the most applicable now, because of the use of metastable technetium-99 (99mTc) in medical isotopes. The connection is that numerous undergraduate students who declare a major course of study in chemistry or biochemistry do so because they wish to go into the health professions in one career or another. The dependence upon this isotope of a rare element that the modern medical community has developed has been a cause of concern to some medical doctors, and a discussion of this has even reached the popular press (5). Clearly, an understanding of where elements are sourced from, and how abundant each is, becomes relevant for a variety of reasons.

Mining and Mining By-Products While several chemicals are obtained primarily through mining and drilling, such as coal, oil, and common salt, currently three metallic elements are produced on an extremely large scale through the mining of a variety of their ores: iron, copper, and aluminum. For example, 29 million metric tons of pig iron, 1.3 million 115

metric tons of copper, and 1.7 million metric tons of aluminum were produced from primary sources – mining – in 2015 alone (2, 3). These numbers do not include recycled metal, which means the amounts actually used are even higher than the just-mentioned numbers. These figures equate to roughly two thirds of aluminum and copper used in the United States being produced domestically, while over 80% of iron in the United States is produced from domestic sources (2). In Table 2 we have already seen a short list of elements refined and used on massive scales, several of which are mined, and in Table 3 we have seen several elements that are not produced as the primary product in some mining operation. Yet the field of metallurgy, with its constantly expanding number of alloys and materials, continues to see an increase in the number of alloys that are finding specific uses, even if they are niche uses. Even those that have been used for decades are being refined to higher purities, or being re-examined through the addition of trace elements (6). Perhaps the most ubiquitous example of a relatively new alloy, at least in terms of the needs of end users, are the magnets in cellular phones (7). Had no neodymium-iron-boron alloy with strong magnetic properties been discovered by General Motors and Sumitomo Special Metals, and subsequently utilized, it is conceivable that cell phones would still be as large as those first marketed in the 1980’s! It is not a great leap of imagination to believe that the future will see still more new, useful alloys and materials, some of which will utilize elements currently not produced on a large scale.

Conclusions First, what can be called an understanding of the origins, the industrial aspects and the end-user aspects of elemental and chemical production and refining can easily be incorporated into any discussion and teaching of the Periodic Table of the Elements. A greater understanding of the disparities of the presence of elements on the Periodic Table can be incorporated into chemistry classes as early as the freshmen year, when students are still engaged in memorizing element names, and can perhaps obviously be reintroduced in upper-level classes (8, 9). Such an approach brings relevance to our teaching, and deepens student understanding, making the table more than just a large group of symbols that students must memorize. Second and importantly, how the uses of various elements and the societal requirements for them change with time can become a means by which the table comes alive for students. The rise of aluminum from an elemental curiosity in the nineteenth century to a common, work horse metal in the early twentieth is one obvious example. The continuing rise in the need for rare earth elements in various applications, such as miniaturized electronics, is another such example. The rise and decline of bromine as a component of flame retardants, and of arsenic in pesticides, serve as still further examples of how the need for a specific element changes over the course of time. Third, the title of this paper points to the fact that some elements exist in widely dispersed deposits, in relatively small amounts, and that they will always be by-products or secondary products in any large chemical extraction 116

and refining process. But these “secondary” elements become vitally important as the needs of society, and the products that enable the modern quality of life, change as we advance. Bringing such facts and this level of understanding to the teaching of secondary school and college-level chemistry classes adds relevance to the discussion, and brings to the fore a realization that the periodic Table of the Elements is a living, changing document in the classification of matter.

References 1. 2. 3. 4. 5.

6. 7. 8. 9.

What it takes to make a new element. Chemistry World 2017, 14 (1), 22–32. USGS. Mineral Commodity Summaries 2015. https://minerals.usgs.gov/ minerals/pubs/mcs/2015/mcs2015.pdf (accessed Jul 7, 2017). American Iron & Steel Institute. http://www.steel.org/ (accessed Feb 26, 2017). Mendeleéff, D. The Principles of Chemistry; P.F. Collier & Son: New York, 1902. Visualizing the Medical Isotope Crisis. Scientific American. https:// blogs.scientificamerican.com/sa-visual/visualizing-the-medical-isotopecrisis/ (accessed Feb 26, 2017). Read, W. T. Industrial Chemistry; John Wiley & Sons, Inc.: New York, 1943. Rare Earth Magnetics. http://www.rareearth.org/ magnets_patents_history.htm (accessed Feb 13, 2017). Benvenuto, M. A. Metals and Alloys: Industrial Applications; Walter DeGruyter GmbH: Berlin, Germany, 2016. Benvenuto, M. A. Industrial Chemistry; Walter de DeGruyter GmbH: Berlin, Germany, 2014.

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Chapter 7

Chemistry’s Decision Point: Isotopes 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’.

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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

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

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

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

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.

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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

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

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

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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117. Coplen, T. B.; Shrestha, Y. Isotope-abundance variations and atomic weights of selected elements: 2016. Pure App. Chem. 2017, 88, 1203–1224. 118. Meija, J.; Coplen, T. B.; Berglund, M.; Brand, W. A.; De Bièvre, P.; Gröning, M.; Holden, N. E.; Irrgeher, J.; Loss, R. D.; Walczyk, T. Atomic weights of the elements 2013 (IUPAC Technical Report). Pure App. Chem. 2016, 88, 265–291. 119. Wieser, M. E.; Holden, N.; Coplen, T. B.; Böhlke, J. K.; Berglund, M.; Brand, W. A.; De Bièvre, P.; Gröning, M.; Loss, R. D.; Meija, J. Atomic weights of the elements 2011 (IUPAC Technical Report). Pure App. Chem. 2013, 85, 1047–1078. 120. Coplen, T. B.; Holden, N. Atomic weights no longer constants of Nature. Chem. Int. 2011, 33, 10–15. 121. Clarke, F. W. The Constants of Nature, Part 5. Recalculation of the Atomic Weights. Smithsonian Misc. Publ. 1882, 441, 1–259. 122. Reedijk, J. On the Naming of Recently Discovered Chemical Elements—the 2016 Experience. Chem. Int. 2017, 39, 30–32.

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Chapter 8

Analytical Methodologies for Arsenic, Selenium, and Mercury: A Historical Perspective Larry Kolopajlo* Chemistry Department, Eastern Michigan University, Ypsilanti, Michigan 48197, United States *E-mail: [email protected].

Historical interrelationships between mercury, arsenic, and selenium are examined and assessed in terms of their environmental impacts, cold vapor/hydride analytical method development, and emergence of the discipline of chemistry. The narrative is accentuated by highlighting important contributions made by environmental chemistry professionals, and is appropriate for an audience of environmental chemical analysts, chemical educators, students, and teachers.

Introduction The first question one might ask about is why arsenic, selenium and mercury are included together in a paper that takes a historical perspective, and the answer is of course that that all three elements have unleashed upon the environment a veritable Pandora’s Box of appalling impacts while tarnishing the image of chemistry in the process. Studying the history of these elements then presents an opportunity for teachers, students, and analysts to learn about the consequences of pollution, and better appreciate the value of environmental initiatives carried out by agencies such as the U.S. Environmental Protection Agency (1), and the World Health Organization (2). What’s more, for decades in the field of analytical chemistry, each of these three elements has been analyzed by cold vapor and hydride methods. It can also be argued that mercury and arsenic are important to the history of chemistry because they have been known and investigated since antiquity (3). Setting itself apart from all other elements by having a privileged place in alchemy, © 2017 American Chemical Society

mercury played a central role in the subsequent shift from alchemy to chemistry. Understanding mercury in today’s context would allow one to trace how mercury has influenced chemical history. Moreover, the history of chemistry should be important to chemical educators, because the discipline helps students and teachers alike, step back and view the big picture of how science developed, and evolved into separate disciplines. Only then can students and teachers alike appreciate the contributions of diverse geniuses whose drudgery led to stunning scientific accomplishments. Unfortunately, however, studies have shown that there is a history knowledge gap in current U.S. secondary teaching and chemistry programs whose over-crowded curricula focus on content, methods, and clinical experience at the expense of a historical perspective. Kaufman (4) has argued that a historical deficit can lead students to formulate a distorted view of chemistry. Furthermore, while Stock (5) demonstrated that few history of chemistry courses exist across the curriculum, Suay-Matallana (6) noted that that the discipline of chemical history has few independent associations and societies dedicated to it. To bridge that gap in secondary teacher preparation programs, studying the history of these elements can imbue chemical educators with an appreciation of the nature of science (7), including the evolution of pure and applied chemistry, the scientific method, the process of scientific discovery, and the human impact of pollution. There also appears to be little historical information available that addresses the history of environmental analytical chemistry, so important in an unsustainable world. Finally, by studying the history of mercury, arsenic, and selenium, one can also understand how chemical history fits within Kuhn’s (8) theory of scientific revolutions, and how changes in one area cause ripple effects in others. The purpose of this chapter is to frame mercury, arsenic, and selenium within their historical context, to understand their environmental impacts and ensuing analytical method development, to demonstrate the interrelationships between these elements, and to communicate how they have played an important role in the transition from alchemy to chemistry. This will be done by integrating the understanding of mercury, arsenic, and selenium from past to present times in such a way as to be of educational value to chemical educators. By reading this chapter, chemical educators can acquire relevant information to enhance their analytical history curriculum, and understand that the history of mercury, arsenic, and selenium are much more than a linear collocation of successive discoveries, and creation of government agencies.

Alchemy to Chemistry Mercury in Arabic Alchemy In this section will be examined how an understanding of mercury, and arsenic to a lesser extent, influenced the development and understanding of Middle Age alchemy (400 – 1400 A.D.), and how the study of these elements, especially mercury, helped catalyze the transition from alchemy to modern chemistry (9). The term alchemy conflates the study of nature (art, chemistry, medicine, and physics) and metaphysics (astrology, mysticism, and spiritualism) 142

into a distinct discipline. Many scholars (10) have supported the idea that alchemy originated in the Middle East between 640 and 1100. There Arabic and Islamic scholars cultivated a strong scientific tradition that had a lasting influence on the arcane, embryonic field of alchemy. For example, today’s commonplace chemical vocabulary contains several terms originating from Arabic alchemy, including for example, alkali (al-qaliy), alcohol (al-kohl), amalgamate, antimony, benzene, borax, elixir (al-iksir), soda, and talcum (11). Furthermore, Arab scholars clearly forged the term alchemy from the definite article al, and kimiye (12). The rich tradition of Arabic alchemy created numerous scholarly manuscripts and lab artifacts, many of which being forerunners of today’s modern instruments, are currently exhibited at international museum collections (13, 14). What’s more, it was Arabic alchemists who devised the theory of the mythical Philosopher’s Stone (15) that could not only transform base metals such as mercury into gold, but when used as an elixir (16), possessed the power to confer immortality. For centuries, alchemy’s consummate but unattainable quest was to synthesize this entity, which today is analogous to the modern theory of the transmutation of metals through radioactive decay. In addition, Arabic alchemy triggered the rise of European alchemy around 1100 (17). Hence the role of alchemy in the Middle Ages relative to other known disciplines is analogous to the role of chemistry as a central science today. Before Arabian alchemists conjured up theories of matter, Greek dogma (18) patently defined matter as being a mixture of four natural elements: earth, water, fire, and air that possessed the four properties of coldness, hotness, wetness and dryness. According to Aristotle’s theory, solar heat caused the earth to emit two types of exhalations, one being cold and moist, and the other hot and dry. When moist exhalations met dry rock, the mixture condensed and converted into metals. In the time of Arabic alchemy, metals were still the most recognized form of matter because they were both abundant, and used for a variety of applications such as jewelry, dining, and weaponry (19). In general terms, it was known that metals could be obtained from ores, and further, that alloys and amalgams could be obtained by mixing melted metals. Moreover, only seven metals were known, and most existed naturally as readily available sulfide ores (20): mercury in cinnabar (HgS), silver in acanthite (Ag2S), copper in chalcocite (Cu2S), lead in galena (PbS), and iron in pyrite (FeS2). Obviously, gold could be found in its elemental state, while tin on the other hand, was found in cassiterite (SnO2). Since tin has a relatively low melting point (232 °C), it could be easily prepared in the pure state by smelting cassiterite (21) in the presence of carbon at low temperatures:

To ancient alchemists, these seven elements must have held a very important place in the universe, and hence they took on a metaphysical status as fundamental principles, not just material building blocks. Of these seven metals, the element mercury, having fascinated scholars since antiquity because of its salient properties, namely the fluidity of water, and the luster and color of silver, hence became known as quicksilver. Thus, alchemists 143

imbued mercury with possessing metaphysical properties that led to it playing a crucial role in shaping both the concept and theory of matter. Mercury compounds were even employed by Egyptians as pigments and for gold mining (22). The most common mercury containing mineral was cinnabar (23), composed of mercuric sulfide, HgS, which is a bright red in color. Important cinnabar deposits have been found across the Middle East (24), for example, in Turkey, the Balkans, Spain (Almadén), Egypt (Giza), and in Italy (Tuscany). Mercury could be prepared by heating cinnabar:

Since ancient times, cinnabar was used in red-colored pigments, red lacquers, and until recently, in cosmetics when its use was halted due to recognition of mercury’s toxicity (25). Of all the Islamic alchemists, the scholar Jabir ibn Hayyan (26), who lived from 721- 813 in present day Persia, is unequivocally recognized as the leading figure. Jabir worked as both a pure and applied alchemist, and in addition, introduced experimenting as a laboratory method. Among the mercury compounds that he investigated, Jabir’s manuscripts list mercury(II) oxide and mercury(II) chloride (27). One of Jabir’s most celebrated accomplishments was his theory of metals (28). In general, some of the speculative ideas that led to this theory are as follows. Since during Jabir’s time: (a) most known metals were mainly obtained from sulfide ores, (b) both mercury and sulfur were easily obtained from cinnabar, (c) mercury was a liquid in its natural state, and other metals could be melted and amalgamated with it, and d) because mercury exhibited special properties, then it would have been reasonable for him to infer that each of the seven known metals were composed of just two substances, mercury and sulfur. Jabir thus unveiled his famous mercury-sulfur theory of metals, each containing a different ratio of mercury to sulfur (29). However, in contrast to the above reasoning, perhaps the Jabirian theory of metals was rationalized in analogy to what was understood about mixing colors in ancient times. For example, it was known that orange pigment could be prepared by mixing red and yellow pigments, the resulting hue depending on the ratio of red to yellow. Furthermore, in Jabir’s time, the minerals cinnabar and orpiment were employed as red and yellow pigments respectively. Since mercury was obtained from cinnabar, and sulfur from orpiment, then it would have been reasonable to assume that all metals had a dual nature deriving from mercury and sulfur. Hence Jabir could have assumed that metals possessed a weighted average of the properties of mercury and sulfur, just like in color mixing. However, this seemingly simple theory was very complicated and represented not only a refinement of the Aristotelean scheme, but a comprehensive extension. It was published in his 7th-century treatise titled the Secret of Creation (30) and later in The Book of Explanation (31). Up to Jabir’s time, little progress had been made toward an understanding of matter, but Jabir’s theory established a new 144

paradigm whose legacy would not only dominate alchemy for nearly a millennium, but would provide a research path for future alchemists to follow. Jabir’s mercury-sulfur theory is so important to the history of chemistry that it merits elaboration. Most significantly, in Jabir’s mercury-sulfur theory, Aristotle’s moist and cold exhalations were superseded by mercury while his hot, dry exhalations were replaced by sulfur. According to Jabir, metals formed as growths inside the earth when sulfur rose as fumes, and combined and fused with mercury under the influence of heat. Different metals formed in different types of soils containing different amounts of sulfur. The silver fluidity of mercury was thought to impart the properties of fluidity and metallic luster to metals, whereas sulfur gave rise to hardness and combustibility. Moreover, both gold and mercury were known to be dense substances, and the fluidity of quicksilver was viewed to be analogous to the property of malleability in gold. Therefore gold, exhibiting a bright luster and being malleable, was the most perfect metal since it was obtained naturally as the pure element, contained the most mercury, and the highest mercury to sulfur ratio. On the other hand, the metal lead was deemed the least perfect of all because it exhibited a dull luster, was soft and malleable, and hence contained the lowest ratio of mercury to sulfur. Since iron was hard and combustible, it was also classified as an imperfect metal and therefore contained mostly sulfur. Jabir further explained the difference between gold and silver by proposing that silver was made from white sulfur (probably white arsenic) and mercury, whereas gold was composed of purified red sulfur (probably the mineral cinnabar or realgar) and the brightest (purified) mercury. The idea that gold contained red sulfur could have originated from observations of decomposing cinnabar which has a bright red-gold color. Moreover, since it was also known that smelting produced a purer metal, there was proof of Aristotle’s postulate that heat played a necessary role in the formation of metals (in reality, alloys). Such observations probably led alchemists to the theory of transmutation of metals, attempting to produce for example, gold from tin. Jabir is also given credit for the discovery of aqua regia, an acidic mixture that could dissolve gold; it contains HCl and HNO3 in a 3:1 ratio (32). Although Jabir believed that mercury and sulfur were the basic building blocks of all metals, the alchemist’s sulfur was not necessarily the natural element known today, but instead was a name applied to any combustible or volatile substance such as arsenic. Therefore, although some scholars (33) have proposed that Jabir had postulated a three-component theory of metals consisting of mercury, sulfur and arsenic, in reality, Jabir’s white sulfur proposed as a constituent of silver was probably white arsenic (As2O3). Because white arsenic was viewed as being just another form of sulfur, arsenic could not be a third fundamental component of metals. Yet Jabir’s theory was even more obscure than it appeared, because most scholars agree that Jabir’s sulfur and mercury were not to be literally taken as the natural elements known today (34). They instead represented definite principles, idealized elements or philosophic elements exhibited only by purified essences of the real substances, and in the context of the mercury-sulfur theory, these substances would be described as philosophic mercury and philosophic sulfur. Hence in Jabir’s theory, gold was constituted when a mixture of mercury and 145

sulfur ripened in the recesses of the earth through heating, with that change transpiring through the interaction of the principles of philosophic mercury and philosophical sulfur. While thousands of works are typically credited to Jabir, scholars generally believe that many of his less important manuscripts were actually written by his students, collaborators, and later followers, and therefore this group of scholars became known as the Jabirian Corpus (35). Although Jabir’s theory of metals was later discredited, it was an important step forward toward a theory of matter. Subsequent investigations of Jabir’s ideas also helped unravel the distinction between elements, compounds, and alloys.

Mercury in European Alchemy Following Arabic alchemy was the rise of European alchemy (1100 - 1700) made possible by the transfer of Arabic knowledge through translation of the works of Jabir and others into Latin by medieval scholars, who not only edited some of his manuscripts, but changed Jabir’s name to Geber, its Latin equivalent (36). Although it is normally assumed that European alchemists accepted Jabir’s mercury-sulfur theory of metals, William R. Newman (37) has advanced the opposing viewpoint that not all European alchemists accepted the mercury-sulfur theory of metals, but instead believed that mercury and sulfur were palpable, independent substances rather than philosophic constituents of metals. Bombastus von Hohenheim (1493-1541), or Paracelsus as he is generally known, is the next historical figure who initiated a scientific revolution that connects with mercury (38). He was born in Switzerland but worked throughout Europe, especially in Germany. He was an iconoclastic medical scholar who having spurned the current norms of medical practice, publicly burned accepted manuscripts, which did nothing to ingratiate himself with his peers. However, by reinventing medicine, his innovative treatments for disease are today recognized as revolutionary, although they remained controversial during his time. Paracelsus is now honored as being the first medicinal chemist and toxicologist because he was first to comprehend the human body as an organism controlled by natural (alchemical) processes, and furthermore, he argued that medical practitioners should study and use alchemy to invent new medicines. Hence, Paracelsus was a forerunner of today’s applied pharmaceutical chemists because he pursued, introduced, and studied novel alchemical compounds using experimentally obtained knowledge to treat intractable diseases of the time like syphilis. He even advocated that physicians discover and introduce less toxic medicines, like for example calomel (Hg2Cl2) to replace the more toxic mercury metal (39). Having discovered the concept of dosage, he urged his peers to limit their prescribed amounts of known toxic compounds, especially those containing mercury and arsenic, and to experiment with innovative alternative compounds to mitigate their unwanted side effects. He even prepared ointments by mixing mercury compounds with oil. For example, mercury(II) oxide was made into an ointment used to treat eye and skin problems, and mercury(II) chloride was used as a disinfectant (40). 146

Although Paracelsus established a Kuhnian revolution in medicine, his most important alchemical contribution extended Jabir’s mercury-sulfur theory to the so-called tria prima or three philosophic principles to explain not only metals, but all natural phenomena (41, 42). Paracelsus’ tria prima subsumed: mercury (representing the properties of changeability and fluidity), sulfur (the property of combustibility) and salt (the qualities of permanence, hardness, incombustibility). Furthermore, he wrote that since the toxins causing all disease were controlled by these three principles, then understanding them could lead to novel cures. However, Paracelsus’ theory came with an added religious twist - that his three philosophical principles were under God’s control. Moreover, he drew the analogy that mercury, sulfur, and salt in metals, corresponded to spirit, soul, and body in man. Obviously, this new theory of metals may have hindered scientific progress. Paracelsus also wrote detailed narratives about the philosopher’s stone (43). The next scientific revolution that incorporated mercury, and led by Robert Boyle (44), was made possible by the work of Evangelista Torricelli who constructed the first mercury barometer in 1643 (45). Robert Boyle (1627-1691) was born in Ireland and there he began his experimental studies in 1649. In 1654 he moved to Oxford, England where his major scientific works were published. Notwithstanding his amazing inventions and outstanding work in physics, Boyle was primarily an alchemist who revolutionized the field by initiating a Kuhnian paradigm shift, from alchemy to modern chemistry. Boyle guided (46) the transition in six ways: (a) by moving alchemy away from its metaphysical influences, (b) reopening the question of Jabir’s and Paracelsus’ theories of matter centering on mercury, and furthermore, (c) by invoking corpuscular (particle) theory to explain chemistry, (d) believing that the universe can be explained using mathematics, (e) by founding the lab-based technique known as chemical analysis, and (f) by urging the use of sound experimental techniques. What’s more, Boyle was a pure scientist, investigating compounds for the sake of academic knowledge, and not for example, just in the quest for the Philosophers’ Stone. But how did mercury play such an important role in Boyle’s alchemical work? Firstly, Boyle used a mercury manometer in his studies on the relationship between the pressure and volume of air (47). Secondly, Boyle penned a few papers where he claimed to have transmuted gold into mercury, invoking the term chrysopeia (chrysos meaning gold and poiein meaning to make) to describe his work on transmutation (48), as he still clung to that false and increasingly indefensible notion (49). In this regard, Boyle claimed that he converted lead into gold whose transformative agent was a red elixir whose recipe he never divulged (50). Although several authors have indicated that Boyle investigated mercury compounds for their medicinal benefits and tested his potions on himself, there is scant proof of this in the literature. However, Boyle did write about a mercury medicine for use against worms in children: “Infuse one Dram of clean Quicksilver all Night in about two Ounces of the Water of Goats Rue, destil’d’ the common way in a cold Stoll: And afterwards strain and filter it, to sever it from all Dregs that may happen in the making of it. This quantity is given for one Dose.” (51). 147

In 1661, Boyle described the distinction between alchemy and chymistry in his landmark treatise titled The Skeptical Chymist (52), today generally recognized as the first modern chemistry textbook. Boyle also contributed to technical vocabulary by referring to chemists as chymists rather than alchymists. However, some authors (53) debate the contention that Boyle created a modern definition of a chemical element. Recognizing the difference between compounds and mixtures, Boyle advanced one of his most important contributions, that of rejecting existing theories of matter, including both Jabir’s mercury-sulfur theory, and Paracelsus’ mercury-sulfur-salt theory. However, Boyle did not advance a new or alternate theory. Instead, he sidestepped the question leaving it open to study. By doing so he encouraged alchemists to continue experimentation and to later refute the theory (54). Overall, Boyle laid the groundwork for a modern chemical discipline. The next important scientific connection to mercury led to development of the field of thermodynamics when Daniel Gabriel Fahrenheit invented the mercury thermometer (55) in 1714. Mercury even played an important role in the discovery of oxygen (56) by Joseph Priestley (1733 – 1804), a religious iconoclast who founded Unitarianism. In his famous 1774 experiment, Priestley generated nascent oxygen (57) by decomposing red calx (mercuric oxide) under heat produced from sunlight focused with a 12 inch “burning lens”:

Having proven that sulfur was not produced when mercuric oxide decomposed, Priestley’s work challenged these widely accepted scientific theories: (a) Jabir’s mercury-sulfur theory of metals, (b) the Paracelsian tria prima, and (b) phlogiston theory (58), the idea that a substance called phlogiston, a component of all combustible materials, was released during combustion. What’s more, Priestley’s reaction became a research model for the investigation of other metal oxides like PbO. Carl Wilhelm Scheele (59) who lived from 1742 to 1786 is also recognized as having a played role in the discovery of oxygen in 1771, a few years before Priestly; however, he did not publish his results until 1777 (60). Scheele had also generated oxygen from mercuric oxide, and from other metal oxides as well (61). Priestley later collaborated with Antoine Lavoisier (1743 – 1794) and helped motivate the studies that led to the chemical revolution, and the recognition of Lavoisier as the father of modern chemistry (62). However, from the time of Lavoisier, mercury no longer played a central role in the development of the science. Priestley’s brilliant experimental work on mercury(II) oxide was perhaps the last step preceding the protracted collapse of the unstable, untenable edifice of alchemy. Although some scholars (63) do not agree that alchemy’s convoluted progression led to the discipline of chemistry, there can be no doubt that alchemy was impeding scientific progress, and had to be abandoned. 148

Arsenic It is said that the element name arsenic derives from Persian (zarnikh) and Greek (arsenikos) roots, meaning potent (64). Realgar and orpiment (65) are the chief arsenic containing minerals, and both are sulfides. Orpiment (As2S3) is bright yellow in color and is found in volcanic fumaroles, hot springs, and hydrothermal veins. Significant deposits of orpiment have been found in Turkey, northern Albania, and near thermal pools in sulfuric acid caves of Aghia Paraskevi on the Kassandra peninsula of northern Greece (66). Realgar (α-As4S4), derives from the Arabic rahl al ghar meaning powder of the mine; it forms brilliant red crystals and is sometimes called "ruby sulphur" or "ruby of arsenic". Realgar melts between 307 and 320 °C, and burns with a bluish flame releasing fumes of arsenic and sulfur. Since ancient times, Greek and Egyptian civilizations, for example, have mined arsenic from its sulfide ores: arsenopyrite (FeAsS), orpiment (As2S3), and realgar. Thus arsenic is widely distributed around the globe. Although in the 5th century AD, Olympiodorus of Thebes is said to have roasted arsenic sulfide to obtain “white arsenic” (As2O3), Jabir is generally given credit for its discovery (67). Jabir was also aware of elemental arsenic which he deposited on copper as a silvery mirror. Today, most commercial arsenic is obtained by heating arsenopyrite (68). The German scholar and alchemist Albertus Magnus (1193-1280) is given credit for the discovery of elemental arsenic in 1250 (69). He discovered arsenic by conducting a chemical reaction between orpiment and soap under heat (70). He further investigated and discovered some of arsenic’s metallic properties. Later in 1649, the German pharmacist Johann Schroeder (71) prepared arsenic by two different methods: (a) heating orpiment with lime, and (b) reducing arsenious oxide with charcoal, and he published his discoveries in the work known as in De Mineralibus (72). Throughout history, arsenic has been employed as both medicine and poison. Although the Greeks used arsenic to treat ulcers (73), Frith (74) wrote that Nero poisoned his step-brother Tiberius Britannicus to become Roman Emperor. In 1365, the Italian City of Sienna passed a law restricting the sale of red arsenic at pharmacies (75). As a health remedy, in the time of Paracelsus, barbers (surgeons) used arsenic to cure wounds, ulcers and other ailments that did not respond to normal treatment (76). In his medicinal formulations, Paracelsus mixed arsenic with saltpeter to concoct potassium arsenate (77), but he also prepared chloride salts of arsenic (78). Such concoctions were used to treat skin diseases and syphilis. Paracelsus also studied arsenic and wrote about how arsenic can transmute red copper into white copper (79). By 1670, potassium arsenate was sold as a general medicine, and alchemists discovered how to make solutions of arsenic by boiling arsenious acid in alkali (80). By 1686, orpiment was mentioned as a medicine used to treat the plague (81). As an internal medicine, later in 1786, a 1% solution of potassium arsenite (KAsO2) became an often prescribed tonic known as Fowler’s solution (82), used until its toxicity was discovered. Scheele generated arsine gas (AsH3) through the heated reaction of As2O3 in a solution containing zinc metal and nitric acid (83). The garlic odor of evolved arsine became a quick test for arsenic. In 1836, James Marsh took Scheele’s work 149

and developed a forensic test for arsenic by inventing an apparatus that allowed for the digestion of biological samples, such as stomach contents, in the presence of zinc metal and nitric acid to generate arsine, which was collected and passed through a heated tube, where it decomposed into elemental arsenic that formed a characteristic gray mirror (84). Selenium Selenium was discovered by Jöns Jacob Berzelius (1779-1848) in 1817 (85), and in the year 2017, the world celebrated its bicentennial (86). One interesting note anecdote (87) regarding selenium’s discovery illustrates the role of serendipity and empiricism in pure and applied science. Sweden has a strong tradition of mining, and Jacob Berzelius’ laboratory in Stockholm became both an R&D and quality control center. Around 1817, Jacob Berzelius formed a business partnership with Johan Gottlieb Gahn to produce sulfuric acid through the lead chamber process in the city of Gripsholm located due west of Stockholm. For this process, cheap iron pyrite (FeS2) was the source of sulfur (pyrite is still used today to generate about one third of sulfuric acid production). One source of pyrite during the time of Berzelius was the Fahlun Mine, located about 160 miles northwest of Stockholm. However, a quality problem unique to pyrite from the Fahlun mine hindered sulfuric acid production. The problem involved formation of a reddish sludge. Since Berzelius insisted on basing his conclusions on sound laboratory science, he analyzed the red sludge through a number of tedious experiments, and found that when combusted, it emitted the distinct horseradish odor of tellurium. However, since tellurium had never been found in Fahlun Mine ores, he continued a further series of painstaking experiments and in early 1818, having reproduced his experimental results in his Stockholm laboratory, concluded that the sludge must contain a new element, which was characterized by its unique physical- (brilliant grayish luster, sublimation properties, azure-blue flame, horseradish odor) and chemical- properties (its unique chemical reactions with metals, oxygen, hydrogen, sulfur, phosphorus, and various salts). Since his results showed that the new element was like tellurium, Berzelius adopted the Swedish name selenium (Greek: selene, moon) for the new element. Hence the element selenium was accidently discovered through analysis of the red sludge that deterred sulfuric acid production.

Toxicity and Safety Although alchemists investigated, employed mercury as a medicine, and endowed it with metaphysical powers, they undoubtedly self-poisoned, not knowing that mercury would be one day viewed as one of the world’s most toxic agents. Today, mercury is ranked #1 on the list of 275 substances in the Substance Priority List as published by the Agency for Toxic Substances and Disease Registry (88) because it poses very significant threats to human health based on known human exposure and toxicity data. What’s more, arsenic and selenium are ranked #3 and #145th (89) respectively on that same list. Mercury and arsenic have 150

been major players in promoting the odious public perception of the chemistry profession, a public image problem that carried forward from medieval times when alchemists cloaked their work in secret symbols (90), concealed gold in aqua regia, and were sometimes exposed as cheaters or charlatans (91). Arsenic has also secured the dubious distinction of being the poisoner’s choice since antiquity (92). Furthermore, the notorious reputation of mercury even played a role in the now retracted papers linking autism to the measles, mumps, and rubella vaccine that contained an organic mercury compound as preservative (93). After that well publicized episode, mercury preservatives were quickly eliminated from all U.S. vaccines. In one sense, nature’s periodic table has rigged the reputation playing field against chemists, since most elements have a long list of associated health and safety hazards on their safety data sheets, and because the public is not well versed in chemical knowledge. Today, although there are scant medical applications of mercury, arsenic, and selenium, the FDA (94) lists about 130 products that contain mercury, usually as phenylmercuric acetate, phenylmercuric nitrate, mercuric acetate, mercuric nitrate, merbromin, or mercuric oxide yellow. In some countries, mercurochrome (an organic mercury compound) is still sold as an antiseptic to treat minor cuts and scrapes. Arsenic trioxide is used to treat a rare disease, acute promyelocytic leukemia (95). Selenium is mainly used as a dietary supplement because of its antioxidant properties. Both HIV and Crohn’s disease (96) are associated with low selenium levels. Selenium, as selenium sulfide, is also used in medicated shampoos (97). The World Health Organization cites (98–100) many severe, chronic problems associated with mercury, arsenic, and selenium and their respective compounds. For example, mercury, especially methyl- and ethyl- mercury (101), are well known as potent neurotoxins that harm the central nervous systems of both humans and wildlife. Inorganic mercury and its compounds are known to damage renal systems (102). Moreover, mercury also injures the digestive and immune systems, lungs, skin, and eyes (103). On the other hand, arsenic causes cancers of the skin, bladder, and lungs, and in addition, induces developmental effects, neurotoxicity, diabetes, pulmonary disease and cardiovascular disease (104). It has been found that long term exposure to selenium causes discoloration of the skin, pathological deformation and loss of nails, loss of hair, excessive tooth decay, and other possible cancers (105). The Morbidity and Mortality Weekly Report (106) as published by the Center for Disease Control, stated that mercury spillage was the most frequent school accident between 2002 and 2007, and therefore it is essential that all science educators be familiar with its hazards, although it has been essentially banned in the U.S. K-12 educational system. Because mercury is a trace element in coal, today it is released to the environment when coal is combusted. Hence one major source (107) of mercury in the air and drinking water is residue from coal-fired power plants. In addition, environmental mercury also derives from erosion of natural geologic deposits, discharge from refineries and factories, and runoff from landfills making mercury a priority pollutant on a global scale (108). 151

However, mercury has long been imparted to the environment through many types of product manufacturing and human usage. For example, since antiquity, it has been used in cosmetic formulations like mascara. In modern cosmetic products, phenylmercuric acetate was used as a mascara preservative until recently. Other anthropogenic sources of mercury have included thermometers, barometers, blood pressure devices, switches, and chemical catalysts, all of which when disposed of improperly add mercury to the environment (109). In a disconcerting recent phenomenon, mercury has been used to make an amalgam to extract gold in artisanal (small scale) gold mining (110, 111), and the resulting pollution has resulted in tragic ecological impacts, especially in South America. Throughout history, mercury has also been used in various products that played a major role in causing disease. For example, in the 18th century, mercury(II) nitrate poisoning caused mad hatter’s disease (112) and resultant brain damage to numerous workers. However, the most dreadful episode imaginable, of mass mercury poisoning occurred in 1956 when over 10,000 inhabitants near Minamata Japan ate seafood from Minamata Bay (113) that was contaminated with biomagnified methylmercury from a chemical production facility. Almost 1,800 of the officially recognized affected victims suffered an agonizing death. The biomagnification of mercury and selenium is a game changer for pollutants because it causes damaging secondary and third-order effects throughout the food chain. In aquatic ecosystems, for example, bacteria and zooplankton convert mercury (II) to methylmercury that concentrates in fat tissue, most affecting top predators like swordfish which have the highest concentrations (114). More recently, the alarm has been raised that methyl mercury is hyper-accumulating in rice (115). Selenium like mercury, can be biomagnified by a factor of 200,000 (116). Being the 20th most abundant element in the earth’s crust, arsenic leaches into groundwater (117) from its widely distributed natural mineral deposits. On the other hand, anthropogenic sources of arsenic contamination include: wood preservatives, coal-fired power plants, and the manufacture of semiconductors (118). In the U.S., although most public water systems have measured arsenic levels in the 2 to 10 ppb range, Western states, and parts of the Midwest and New England have higher concentrations (119). For example, it has been reported that in Fallon, Nevada arsenic has been found at 80 ppb (120). A tragic example of arsenic groundwater contamination (121, 122) took place in the Ganges Delta of Bangladesh where tube-wells, drilled in the 1960s to provide a source of potable water, were found in the 1990s to be contaminated with arsenic. This case represents the largest mass poisoning in history affecting over 50 million people who were chronically exposed to arsenic at levels over 10 ppb. Moreover, it was discovered that about 1.4 of 4.7 million tube-wells contained arsenic at levels greater than 50 ppb affecting more than 20 million people (123). Deeper tube-wells (more than 150 m in depth) have been found to be safer while shallow wells between 20 and 100 m contain the most toxic water (124).

152

Table 1. MCL and MCLG Data for Hg, As, and Se Mercury μg/L

Arsenic μg/L

Selenium μg/L

1942: U.S. Public Health Service (129) Limit

NA

50

50

1946: Public Health Service (130) Limit

NA

50

50

1962: Public Health Service (131) MCL

NA

10

10

1974: Safe Water Drinking Act (SWDA) MCL

0.002

10

50

SWDA: July 30, 1992 (132) MCL

0.002

50

50

MCLG

0.002

0

50

SWDA: January 23, 2006 (133) MCL

0.002

10

50

MCLG

0.002

0

50

The EPA and Safe Water Drinking Act: History Cases like some of the foregoing instances of mass poisonings caused by natural and anthropogenic sources, led to a Paracelsian paradigm shift regarding public approaches to protecting human health as implemented by WHO (125), and in the U.S., by the establishment of the EPA on December 2, 1970 (126). Moreover, in the U.S., the Safe Water Drinking Act (SWDA) of 1974 was enacted to protect the quality of drinking water from toxins such as mercury, arsenic, and selenium (127). To carry out this mission, the EPA devised and set: (a) a maximum containment level (MCL) and (b) a maximum containment level goal (MCLG) for toxic elements and compounds (128). The MCL is an enforceable standard regulating the maximum level of a contaminant allowed in drinking water, while the MCLG is established to provide a level below which no health risk is expected, but it is not enforceable. The SWDA directs the EPA to set the MCL as close to the MCLG as is technically feasible, requiring public water systems to detect and remove contaminants using suitable treatment technologies that are feasible and affordable. When health benefits are great, methods are cost-effective, and when technology is readily available, MCLs are set close to MCLGs. Currently there are fewer than 100 chemicals for which an MCL has been established. Legal limits, 153

as well as MCL and MCLG values for arsenic, mercury, and selenium are listed below in Table 1 in chronological order according to how they were changed by the Public Health Service (129–131) and the Safe Water Drinking Act (132, 133).

History of EPA Mercury Cold Vapor and Hydride Methods Mercury Boyle’s visionary acumen in establishing the field of chemical analysis is today recognized through the Robert Boyle Prize for Analytical Science (134) as awarded annually by The Royal Society of Chemistry. The value of analytical environmental chemistry is also exemplified in the work done by often neglected analysts whose work in normal science enables the implementation of safety regulations protecting populations from the many health hazards associated with arsenic, mercury, and selenium. EPA approved methods for the determination of arsenic, selenium and mercury can be placed into five groups: graphite furnace atomic absorption (GFAA), inductively coupled atomic plasma (ICP), cold vapor atomic absorption spectrometry (CVAAS), AA-hydride and spectrophotometric. Although graphite furnace (135) has been used to analyze for arsenic and selenium at low detection limits, it has never been a viable choice for mercury, due to its high vapor pressure at room temperature. On the other hand, while inductively coupled atomic plasma (136) analyzes for all three elements, required detection limits for drinking waters cannot be met using it. Spectrophotometric methods have not been used for mercury due to its volatility. Of these five methods, only the classes of cold vapor and hydride methodology can analyze for all three elements at detection limits mandated for drinking waters. In the following sections, the term cold vapor applies to mercury analyses whereas gaseous hydride refers to analyses for arsenic and selenium. The history of cold vapor analytical methods as used by the EPA to analyze mercury in drinking waters will be covered first. In 1963 Poluektov et al. (137) suggested a new analytical method, dubbed cold vapor, as a tool for mercury analysis, but it was Hach and Ott (138) in 1968, who first reported its successful use, analyzing geological samples, both rocks and soils, for mercury to a level of 1 ppb. A schematic of instrumentation typically used is shown in Figure 1. Sample preparation takes place in a BOD bottle where a mixture of various acids and potassium permanganate break down the sample matrix, converting bound mercury to Hg2+. To determine the concentration of mercury using cold vapor atomic absorption (CVAA), mercury ion in the sample is first reduced to its elemental form with stannous chloride:

154

Figure 1. Mercury cold vapor instrument schematic.

A stream of inert gas is then pumped into the sample bottle to carry mercury vapor into and through a 10 cm quartz absorption cell that is placed in the optical path of an AA instrument containing a mercury hollow cathode lamp that emits radiation at 253.7 nm. A uv detector measures the intensity output which is a linear function of the mercury concentration in the sample. The original technique had a detection limit of about 0.2 micrograms/liter, which hasn’t changed over the years. While Hatch and Ott provided the initial breakthrough instrumental method in 1972, Kopp, Longbottom and Lobring (139) published a historic paper, titled “Cold Vapor” Method for Determining Mercury in the Journal of the American Water Works Association that greatly influenced the development of published EPA methods relating to mercury cold vapor. Their method determined total mercury in drinking, surface, ground, sea, brackish waters, and industrial and domestic wastewaters and soils. The paper’s major strength was its systematic approach that solved several critical problems relating to mercury analysis, and laid a framework for EPA’s future step by step analytical methods. For example, it demonstrated that potassium permanganate was ineffective in releasing mercury from organic compounds such as phenyl mercuric benzoate, and showed that potassium persulfate (K2S2O8) was necessary in the digestion of organic samples, on the one hand to release bound mercury, and on the other hand to remove organics whose broad uv absorption bands overlap the 253.7 nm mercury line. It also provided in depth data demonstrating that copper, tellurium, and other heavy metals did not interfere with the analysis, except at very high concentrations, possibly over 10 mg/L. Moreover, it showed that hydrogen sulfide, present in sewage samples, was eliminated by potassium permanganate. The paper also pinned down air flow rates, how to eliminate water vapor interferences, calibration, and provided statistical data to validate the method through reproducibility and spiked sample recovery. One of the authors of the just mentioned 1972 landmark paper, Larry B. Lobring, worked in the Inorganic Chemistry Branch of the Chemistry Research Division Environmental Monitoring System Laboratory located in Cincinnati, Ohio. He coauthored numerous EPA methods and manuals on the determination of metals, including mercury in environmental samples. One of his most 155

substantial contributions in the realm of mercury analysis, was based on his previous work, and resulted in the enduring Method 245.1, Determination of Mercury in Water by Cold Vapor Atomic Absorption Spectrometry (140). This method was specific for drinking waters, and could be applied to analyze for mercury in a wide variety of forms, such as phenyl mercuric acetate (cosmetics), and methyl mercury (fat tissue in sea food). Sample preparation involved converting bound mercury to divalent mercury ion through a two hour digestion in a glass BOD (biological oxygen demand) bottle with a persulfate- permanganate mixture under heating. Subsequent research focused on lowering the detection limit, on removing, circumventing or avoiding matrix interferences, increasing speed through automation, implementing digestion cell modifications to remove interferences and, instrumental modifications. Following development of method 245.1, an automated method, 245.2. was issued in 1974 (141). In 1984, Varian Instruments, later acquired by Agilent, marketed its VGA-76 assembly (142) that allowed for an automated and rapid determination of many samples using EPA approved methods. Method 245.1 went through several revisions. However, all calibration procedures and detection limits (0.2 ppb) remained essentially the same. One of the most striking differences between revisions was that revision 3 offered three options for apparatus design whereas only one is offered through revision 2.3. Revision 2.3 and 3 also introduced large sections devoted to quality control (QC) whereas the original paper did not address the issue. Revision 3 also contained an extensive section on sample collection, preservation and storage whereas minimal information is provided in revision 2.3, and nothing in the original version. As far as the wet digestion procedure goes, the original paper gave a two-step procedure that is expanded to eight steps in revision 2.3, and expanded to more steps in revision 3. As far as water quality, the original papers through revision 2.3 did not contain any information about water purity, but revision 3 required ASTM Type II water. Also, in revision 3, it is stated that the analyst must be vigilant regarding environmental contamination and background interferences arising from unusual sample matrices. This procedure also warned that when doing low level work, the analyst must physically separate both Kjeldahl nitrogen and chemical oxygen demand (COD) determinations from mercury determinations because they contained the reagent mercury(II) sulfate. These interferences were described: sulfide, copper, chlorides, tellurium, chlorine and other volatiles whereas the original interferences specified only: sulfide, copper, chlorides, and volatile organics. In addition, the first published method did not contain a safety section but revision 2.3 contained an expanded safety section, including: (a) notes on mercury toxicity, (b) indications that a fume hood should be used, (c) directions that workers be immunized when human waste is analyzed, and (d) to watch for the evolution of sulfide and cyanide. In 1998, EPA released Method 1631 (143) titled: Mercury in Water by Oxidation, Purge and Trap, and Cold Vapor Atomic Fluorescence to meet the lower mercury water quality criteria (WQC) published through section 304(h) of the Clean Water Act, the clean water programs of 1999 in the National Toxics Rule (144), and in the Final Water Quality Guidance for the Great Lakes System. Bloom (145) and Fitzgerald (146) played major roles in developing the method. 156

Method 1631 established formal procedures to quantitate low level mercury in water at detection levels of 1 parts per-trillion (ppt), although the fluorescence method allows determination of mercury to 0.5 ppt. The only interferents listed are gold and iodide. It utilizes a class-100 clean room. In 1999, after public comments were made, revision B was released, addressing correction of test sample results for reagent blanks. In 2001, revision C added requirements for the reporting and use of field blank results. It also clarified that field blank results must be reported separately, and that a field blank correction must be performed if requested or required by a regulatory authority or in a permit. In 2002, revision E corrected many minor clarifications, technical errors and inconsistencies. Regarding sample preparation, according to Method 1631, samples are digested using bromine monochloride (BrCl) to oxidize bound mercury to Hg2+(aq) which then undergoes sequential reduction with ammonium hydroxide and stannous chloride to convert Hg(II) to volatile Hg(0). Finally Hg(0) is purged from water and preconcentrated onto a gold-coated sand trap, followed by thermal desorption from the trap, and detection by cold-vapor atomic fluorescence spectrometry. Although the major advantage of CVAFS techniques is a low detection limit, there are several disadvantages including: (a) the periodic replacement of gold traps, and (b) that some sample matrices pose challenges, especially those with high volatile organic concentrations. Comparing CVAFS methods, Method 1631 is at least 500 times more sensitive than Method 245.1. Furthermore, Method 1631 is also faster than Method 245.1 because it avoids a lengthy permanganate digestion. Although the last revision of Method 245.1 was completed in 1994, the method is still in use today. Nevertheless, Method 245.1 has generally been supplanted by EPA Method 1631, cold-vapor atomic fluorescence spectrometry. In 2005, EPA also approved a second CVAFS procedure, Method 245.7, that is very similar to 1631, and quantifies mercury in filtered and unfiltered water, and is applicable to drinking water, surface and ground waters, marine water, and industrial and municipal wastewater. Although both methods require use of a CVAFS detector to measure low levels of mercury, Method 1631 uses oxidation, a purge and gold trap isolation, and desorption followed by CVAFS, while Method 245.7 uses liquid-gas separation and a dryer tube for isolating mercury analyte. Method 245.7 suffers from gold, silver and iodide interferences. Since both methods 1631 and 245.7 meet a detection limit of about 1 ppt, they are normally carried out in a “clean room” environment to prevent contamination of samples, and moreover, sampling procedures required must also implement special cleanliness procedures. One advantage of Method 245.7 is that it is faster and simpler than Method 1631 because it avoids a heating step, so that digestion blocks are not required. In August 1998, EPA released Method 1630 (147) for the determination of methyl mercury (CH3Hg) in water by distillation, aqueous ethylation, purge and trap by CVAFS; its range was 0.02 to 5 ng/L. Table 2 lists EPA approved mercury cold vapor methods for water in chronological order, along with their most important advancements. In Table 2 below are summarized the major published revisions in methods 245.1 (148–151), 245.7 (152–154), and 1631 (155–161) in chronological order. 157

Table 2. Timeline of EPA Mercury Cold Vapor Methods 245.1: Mercury in water by CVAAS. (Detection Limit = 0.2 μg/L). Revision

Critical Changes

1.0 in 1972 (148) 2.0 in 1979 (149)

1. Quality control addressed. 2. Includes interlaboratory precision data.

2.3 in 1991 (150)

Emphasizes and improves quality control

3.0 in 1995 (151)

1. Incudes three options for the apparatus. 2. Interferences from COD and Kjeldhal nitrogen analyses. 3. Quality control greatly expanded.

245.7: Mercury in water by CVAFS. (MDL = 1.8 ng/L). Revision

Critical Changes

1.1 in 1996 (152) Draft in 2001 (153)

The draft is based on collaborative work between EPA’s Environmental Monitoring Systems Laboratory, EPA-Region 4, and Technology Applications, Inc. and on peer-reviewed publications.

2.0 in 2005 (154)

1. Addresses matrix spikes and precison. 2. This method is performance based.

1631: Mercury in water by CVAFS. (MDL = 0.2 ng/L). Revision

Critical Changes

Proposal/Draft in 1991/1995 (155, 156)

Result of the Clean Water Act requiring lower Hg detection levels.

Validation in 1996 (157)

Provided guidance for quality control, interferences, technical difficulties, and general use of the method.

CFR in 1998 (158)

Published as law in the federal register.

B in 1999 (156)

Released after public comments.

C in 2001 (160)

New requirements were made regarding field blank results.

E in 2002 (161)

Clarifications and corrections were made for technical errors and inconsistencies.

Today’s analytical challenge is to determine the concentrations and distribution of various mercury species in the hydrosphere. The future of mercury cold vapor method may rely on new instrumental methods in which cold vapor is combined with ICP-MS (162), GC-AFS (163), and IC-ICP-MS (164) that allow determinations of individual mercury species to (10-15) ppq levels. 158

Arsenic and Selenium Analytical methods centering on hydride generation for the determination of arsenic and selenium are described in this section. The Gutzeit method for arsenic analysis was first published as a qualitative test in 1879 (165), that improved on Marsh’s arsenic mirror. In 1907, Sanger and Black (166) modified it into a reliable quantitative measure of arsenic via a hydride method. In their determination, arsenic in the sample is first reduced to As(III) as arsenious acid (H3AsO4) using for example, a reducing agent such as potassium iodide. Hydrogen gas, generated through the reaction of granulated zinc with hydrochloric acid in a second apparatus, is then passed into the arsenious acid solution producing arsine gas:

Arsine then reacts with paper impregnated with mercuric chloride to produce a yellow stain (AsHg3Cl3) whose color intensity is indicative of the quantity of arsenic in the sample. In 1942, Jacobs and Nagler (167) published a paper that reviewed seven existing methods for determining arsenic at low levels, with the object of finding a better quantitative method. The authors then recommended a new method that combined the Gutzeit and molybdenum blue methods. In this modified method, arsine gas is generated using the Gutzeit procedure, and is trapped in a sodium hypobromite solution. The sample is then transferred to a Nessler tube, and ammonium molybdate is added. After 30 minutes, the developed blue color is compared to a standard for a visual determination of arsenic that is quantitative with a detection limit of 0.038 μg/L. Its range was reported as 1.5 to 50 μg. However, both phosphorus and silicon were found to interfere with the determination because they also form blue molybdenum complexes. Following the successful development of the Jacobs-Nagler colorimetric method, in 1942, the U.S. the Public Health Service mandated that it be required for analyzing arsenic. The Public Health Service also established an As and Se drinking water standard at 50 μg/L. In 1952, Vasak and Sedivek (168) reported the silver diethyldithiocarbamate (SDDC) method for quantitatively determining arsenic in water samples. The SDDC colorimetric method for arsenic involved reducing inorganic arsenic to arsine gas using a reducing agent and hydrogen generated from the zinc/acid reaction. The AsH3 hydride generated was then scrubbed though glass wool infused with lead acetate, and then absorbed in SDDC-pyridine solution. Arsine then formed a red complex with SDDC, whose absorbance was easily measured at 535 nm using spectrophotometry. In 1962, three papers were published that established SDDC as a new quantitative method for trace arsenic analysis. Stratton et al. (169) discussed the strengths and weaknesses of the SDDC method established by Vasak and Sedivek. Ballinger et al. (170) reported that the SDDC method for arsenic was superior to alternate methods such as the (a) heteropoly blue method, and (b) the Gutzeit-heteropoly blue methods. Advantages included improved time of analysis, ease of analysis, precision, and accuracy. Also in 1962, Clarke et al. 159

(171) reported that the detection limit for arsenic by the SDDC method was 0.03 mg/L. In 1962, the Public Health Service lowered the arsenic drinking water standard to 0.010 mg/L for which the SDDC colorimetric methods became the approved method. In 1971, based on the previously described work, the EPA issued Method 206.4 (172) to determine arsenic (total, dissolved, suspended) by SDDC to a limit of 10 μg/L. In 1978, Sandhu (173) published a modification of the As-SDDC Method addressing these metal ion interferences: antimony, cobalt, copper, chromium, mercury, and molybdenum. He also studied arsenic recovery in the presence of the afore-mentioned interferents. In 1974, Method 206.3, the atomic absorption gaseous hydride method (174) for determining inorganic arsenic at a detection limit of 2 μg/L was published by the EPA, and was designed to be applicable to drinking waters. Interferences arise when sample matrices contain high concentrations of chromium, cobalt, copper, mercury, molybdenum, nickel or silver. For sample preparation, arsenic is first reduced to As3+ using SnCl2, and then converted to arsine, AsH3, using zinc metal and HCl. The gaseous hydride is then swept into an argon-hydrogen flame of an atomic absorption (AA) spectrophotometer and analyzed at 193.7 nm. The working range of the method is 2 -20 μg /L. In February, 2002, Method 206.3 was removed from EPA’s list of analytical methods because other technologies with lower detection limits became available and were more useful (175). In April 1995 was published the first draft of Method 1632, Determination of Inorganic Arsenic in Water by Hydride Generation Flame Atomic Absorption Absorption, which was developed under the direction of William A. Telliard of the U.S. EPA. (176). This method determines dissolved and total arsenic in the 10 - 200 ng/L concentration range. The analysis involves treating an aqueous sample with 6 M HCl and 4% NaBH4 to form various hydrides such as arsine, AsH3(g). The gas is purged and collected in a liquid nitrogen gas trap packed with 15% OV-3 Chromasorb. Upon desorption, the various hydrides are swept into a flame AA instrument where absorbance is measured at 193.7 nm. The method is based on the work of Braman (177) and Andreae (178). In January, 2001, revision A (179) was published that allowed determination of dissolved, total and arsenic species through hydride generation quartz furnace AA. The following species may be analyzed: total inorganic arsenic (As), arsenite (As3+), arsenate (As5+), monomethylarsonic acid (MMA), and dimethylarsinic acid (DMA). Dissolved and total arsenic can be determined in the 10 to 200 μg/L range. Having released Method 1632, the colorimetric-SDDC methods for arsenic were both removed and did not appear in the Feb. 19, 2015 EPA proposed rules. In the 1996 amendments to the SDWA, Congress directed the EPA to consider a new arsenic regulation. Later, in 2006, arsenic was set at 10 ng/L and Se raised to 50 ng/L. Although GFAA is the preferred method for determining As and Se in drinking water, since the theme of this paper focuses on methods common to Hg, As and Se, furnace methods won’t be discussed further. Analytical methods for the analysis of selenium follow the pattern set by arsenic. In 1958, Danzuka and Ueno (180) published a colorimetric method to determine trace levels of selenium(IV) using 3,3’-diaminobenzidine. This 160

colorimetric method measures selenium at 420 nm in the pH range 6-7, but does not employ a hydride. Method 270.3 (181) was issued in 1974 to determine inorganic selenium in drinking water by the gaseous AA hydride method. It covers inorganic selenium in water (drinking, fresh, saline). In the analysis, selenium is reduced from the +6 oxidation state to the +4 oxidation state using of SnCl2. Zinc is then added to the acidified sample, producing hydrogen gas which reacts with selenium to produce the hydride, SeH3, which is carried into an argon-hydrogen flame of an AA spectrophotometer and analyzed at 196.0 nm. The method determines Se to a 2 ng/L level. Interferences may occur in the presence of high concentrations of chromium, cobalt, copper, mercury, molybdenum, nickel and silver. Table 3 shows approved and retired EPA methods for arsenic and selenium analyses. A compendium of EPA analytical methods is available (182, 183).

Table 3. EPA Methods for Inorganic As and Se Water Instrumentation

Date Issued

Range ng/L

λ nm

LOD ng/L

206.2

As: AA-Furnace

1978

5-100

193.7

1

206.3 retired

As: Ar/H2 AA-GaseousHydride

1974

2-20

193.7

2

206.4 retired

As: Spectophotometric-SDDC in pyridine

1971

535

10

1632

As: 4% NaBH4 Hydride Quartz Furnace AA

1998

0.01-50

193.7

0.003

270.2

Se: AA-Furnace

1978

5-100

196

2

270.3 retired

Se: Ar/H2 AA-Gaseous Hydride

1974

2-20

196

2

Conclusion and Summary Analyzing a historical perspective on arsenic, mercury and selenium demonstrates that Jabir, Paracelsus, Boyle, and the EPA respectively launched scientific revolutions in alchemy, medicine, chemistry, and environmental safety. Hence studying the history of these elements can help educators see the big picture of how these three elements impacted society and influenced scientific thought through the ages. In the age of alchemy, mercury was viewed as a paradigm in itself, but in modern times, both mercury (and arsenic compounds) were finally recognized as rather blunt instruments for treating disease when more precise and effective tools were needed. Before the Safe Water Drinking Act of 1974, few methods existed for the analyses of mercury, arsenic, and selenium. However, after its inception, the need to reduce exposure to arsenic, mercury, and selenium in drinking waters spurred the EPA to invent new methods for their detection. Mercury detection limits were 161

first regulated in 1974 at 2 ng/L (ppb) as determined by cold vapor. In 1991, CVAFS techniques allowed a much lower the detection limit to 0.001 ng/L (1ppt) to be met. Arsenic and selenium were first regulated in 1942. In 2006, the MCL for arsenic was set to zero because no level was considered safe. In 1974, qualitative methods were replaced by spectrophptometric methods when Flame AA-hydride methods were advanced. In 1988 AA hydride quartz furnace allowed arsenic to be determined at 0.01 μg/L. However GFAA is still a superior method allowing As and Se determinations to 1 ng/L.

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121. Chowdhury, U. K.; Biswas, T. R.; Chowdhury, T. R.; Samanta, G.; Mandal, B. K.; Basu, G. C.; Chanda, C. R.; Lodh, D.; Saha, K. C.; Mukherjee, K.; Roy, S.; Kabir, S.; Quamruzzaman, Q.; Chakraborti, D. Groundwater arsenic contamination in Bangladesh and West Bengal. India Environ. Health Perspect. 2000, 108, 393–397. 122. Smith, A. H.; Lingas, E. O.; Rahman, M. Contamination of drinking-water by arsenic in Bangladesh: a public health emergency. Bull. W. H. O. 2000, 78, 1093–1103. 123. Nickson, R.; McArthur, J.; Burgess, W.; Ahmed, K. M.; Ravenscroft, M. Arsenic poisoning of Bangledesh Groundwater. Nature 1999, 401, 545–546. 124. Singh, A. K. Chemistry of arsenic in groundwater of Ganges-Brahmaputra river basin. Curr. Sci. 2006, 91, 599–606. 125. WHO. Guidelines for Drinking Water Quality, 3rd ed.; World Health Organization: Geneva, 2008. http://www.who.int/water_sanitation_health/ dwq/fulltext.pdf (accessed February 12, 2017). 126. EPA History. The Origins of EPA. https://www.epa.gov/history/origins-epa (accessed February 3, 2017). 127. Safe Water Drinking Act (SDWA). https://www.epa.gov/sdwa (accessed May 7, 2017). 128. MCL and MCLG. https://safewater.zendesk.com/hc/en-us/articles/ 211405328-What-do-MCL-MCLG-and-MRDL-mean- (accessed May 7, 2017). 129. Public Health Service. Public Health Drinking Water Standards. Public Health Rep. (1896−1970). 1943, 58, 69–82. 130. Public Health Service Drinking Water Standards. Public Health Rep. (1896–1970). 1946, 61, 371–384. 131. U.S. Department of Health, Education and Welfare, Public Health Service Drinking Water Standards: Public Health Service Publication 956, Revised 1962, Washington 25 DC. https://nepis.epa.gov/Exe/ZyNET.exe/ 2000TP5L.TXT?ZyActionD=ZyDocument&Client=EPA&Index=Prior+to+ 1976&Docs=&Query=&Time=&EndTime=&SearchMethod=1&TocRestrict =n&Toc=&TocEntry=&QField=&QFieldYear=&QFieldMonth=&QField Day=&IntQFieldOp=0&ExtQFieldOp=0&XmlQuery=&File=D%3A%5C zyfiles%5CIndex%20Data%5C70thru75%5CTxt%5C00000002%5C2000TP 5L.txt&User=ANONYMOUS&Password=anonymous&SortMethod=h%7C&MaximumDocuments=1&FuzzyDegree=0&ImageQuality=r75g8/r75g8/x 150y150g16/i425&Display=hpfr&DefSeekPage=x&SearchBack=ZyAction L&Back=ZyActionS&BackDesc=Results%20page&MaximumPages=1& ZyEntry=1&SeekPage=x&ZyPURL (accessed May 7, 2017). 132. Safe Water Drinking Act (SWDA). United States Environmental Protection Agency. https://www.epa.gov/sdwa (accessed February 3, 2017). 133. Tiemann, M. Safe Drinking Water Act (SDWA): A Summary of the Act and Its Major Requirements, Congressional Research Service, Feb. 5, 2014. https:/ /fas.org/sgp/crs/misc/RL31243.pdf (accessed Jan 10, 2016). 134. Robert Boyle Prize for Analytical Science. http://www.rsc.org/ ScienceAndTechnology/Awards/BoylePrize/ (accessed February 8, 2017). 168

135. EPA Method 200.9 Revision 2.2: Determination of Trace Elements by Stabilized Temperature Graphite Furnace Atomic Absorption. Environmental Monitoring Systems Laboratory Office of Research and Development; U.S. Environmental Protection Agency: Cincinnati, OH, 1994. 136. EPA Method 200.8 Revision 5.4: Trace Elements in Water and Wastes by Inductively Coupled Plasma-Mass Spectometry. Environmental Monitoring Systems Laboratory Office of Research and Development; U.S. Environmental Protection Agency: Cincinnati, OH, 1994. 137. Poluektov, N. S.; Vikun, R. A.; Zelyukova, Y. V. Atomic fluorescence spectrometry of mercury: principles and developments. Zh. Anal. Khim. 1963, 18, 33–37. 138. Hatch, W. R.; Ott, W. L. Determination of submicrogram quantities of mercury by atomic absorption spectrophotometry. Anal. Chem. 1968, 40, 2085–2087. 139. Kopp, J. F.; Longbottom, M. C.; Lobring, L. B. Cold Vapor Method for Determining Mercury. J. Am. Water Works Assoc. 1972, 64, 20–25. 140. EPA Method 245.1: Determination of Mercury in Water by Cold Vapor Atomic Absorption Spectrometry. Environmental Monitoring Systems Laboratory Office of Research and Development; U.S. Environmental Protection Agency: Cincinnati, OH, 1974. 141. EPA Method. 245.2: Mercury (Automated Cold Vapor Technique) by Atomic Absorption. Environmental Monitoring Systems Laboratory Office of Research and Development; U.S. Environmental Protection Agency: Cincinnati, OH, 1974. 142. An Automated Vapor Generation Accessory for Atomic Absorption Analysis. http://www.agilent.com/cs/library/applications/AA038.pdf (accessed February 12, 2017). 143. EPA Method 1631: Mercury in Water by Oxidation, Purge and Trap, and Cold Vapor Atomic Fluorescence Spectrometry. 2002. Environmental Monitoring Systems Laboratory Office of Research and Development; U.S. Environmental Protection Agency: Cincinnati, OH, 2002. 144. EPA Method 1630: Methyl Mercury in Water by Distillation, Aqueous Ethylation, Purge and Trap, and Cold Vapor Atomic Fluorescence Spectrometry. Environmental Monitoring Systems Laboratory Office of Research and Development; U.S. Environmental Protection Agency: Cincinnati, OH, 1998. 145. Prestbo, E. M.; Bloom, N. S. Mercury speciation adsorption (mesa) method for combustion flue gas: metodology, artifacts, intercomparison, and atmospheric implications. Water, Air, Soil Pollut. 1995, 80, 145–158. 146. Fitzgerald, W. F.; Gill, G. A. Sub-Nanogram Determination of Mercury by Two-Stage Gold Amalgamation and Gas Phase Detection Applied to Atmospheric Analysis. Anal. Chem. 1979, 15, 1714–1720. 147. EPA Method 1630: Methymercury in Water by Distillation, Aqueous Ethylation, Purge and Trap, and Cold Vapor Atomic Fluoresence Spectrometry. Environmental Monitoring Systems Laboratory Office 169

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Chapter 9

Recent Advancements in the Radiochemistry of Elements Pertaining to Select Nuclear Materials and Wastes Eric S. Eitrheim, Andrew W. Knight, Michael K. Schultz, Tori Z. Forbes,* and Andrew W. Nelson Department of Chemistry, University of Iowa, Iowa City, Iowa 52242, United States *E-mail: [email protected]. Phone: 319-384-1320.

A 2012 National Academy of Sciences report indicated insufficient numbers of post-graduate level radiochemists and nuclear scientists to fulfill the current and future national needs in this sector. This report prompted the creation of a radiochemistry graduate program at the University of Iowa with research efforts in energy, environmental, and medical applications. These research projects, outlined here, have included a focus on understudied elements whose chemistry is still poorly understood. The first study targets methodology to separate gallium (Ga) and plutonium (Pu) isotopes for nuclear fuel applications. Second, we explore the presence of naturally occurring radioactive material (uranium (U), thorium (Th), radium (Ra), lead (Pb), bismuth (Bi), and polonium (Po)) in drill cuttings associated with unconventional drilling operations. Last, the fundamental chemistry of protactinium (Pa) relating to geochronometry and the nuclear fuel cycle is discussed, specifically focusing on separation and dating techniques. The University of Iowa radiochemistry program has aided in this exploration with anticipation that additional research will continue to increase our understanding of these elements in the future.

© 2017 American Chemical Society

Introduction In 2012, the National Academy of Sciences released a report indicating that there were dwindling numbers of radiochemists and nuclear scientists graduating from U.S. graduate programs (1). The numbers of radiochemists and nuclear scientists graduating from Ph.D. programs is particularly low and expected to be well below the demand by employers in the coming decade. In response to the low rate of Ph.D.-trained radiochemists, the University of Iowa launched an interdisciplinary radiochemistry program that included research and training in medical, environmental, and energy applications of radiochemistry (2). The first cohort of three analytical radiochemistry trainees began in 2012, with research focusing on understudied radionuclides that were ripe for new discoveries. In particular, the research focused on gallium (Ga) and plutonium (Pu) with applications in mixed-oxide (MOX) fuels, naturally occurring uranium-series radionuclides in the environment, and fundamental protactinium (Pa) chemistry. The following chapter will highlight what is known, what the University of Iowa cohort of three discovered, and what still remains to be explored for select radioactive elements, including: plutonium (Pu), protactinium (Pa), uranium (U), thorium (Th), radium (Ra), lead (Pb), bismuth (Bi), and polonium (Po).

Plutonium Nuclear Materials and Applications in Nuclear Forensics Plutonium (Pu) nuclear materials originated shortly after discovery of the element by Glenn T. Seaborg and colleagues in December of 1940 (3). This discovery was not initially announced due to the understanding that this new element had application in nuclear weapons technologies (3). Plutonium was initially produced at the Hanford Site in Washington State using a nuclear reactor; most notably reactor B (4). While production and isolation of gram quantities of Pu began almost immediately, additional materials engineering and processing was necessary to stabilize Pu for use in nuclear weapons. One of the most important advancements in this area was the discovery of plutonium-gallium alloys during the Manhattan Project, which effectively stabilized the material for use in nuclear weapons (5). Pure Pu has six allotropes and the thermodynamically stable phase at standard temperatures and pressures is the high density (19.86 g/cm3) α-form. The lowest density phase is δ-Pu, which is approximately 25% less dense (16.0 g/cm3) than the α-phase and can only be stabilized at temperatures between 310 and 425 °C. Super criticality is only reached through a compression based δ-α transformation, thus maintaining the δ-Pu material under standard conditions is crucial for the design of a functioning nuclear weapon. Creating a Pu alloy stabilizes the δ-phase at room temperature and this is typically achieved by adding ~0.8-1 mol% Ga to the Pu metal. Alloying Ga with Pu also improves the mechanical properties by decreasing corrosion rates, improving the ability to cast, and creating low thermal expansion properties leading to better material processing for nuclear material production. 174

Thousands of nuclear weapons have been created using Ga-Pu alloys, and consequently these materials pose a nonproliferation risk (6). According to the International Atomic Energy Agency (IAEA), eight incidences of theft or loss of plutonium-based nuclear materials have been reported globally, excluding Pu-based smoke detectors, from 1993 through 2015 (7). When the authorities confiscate these now-illicit Pu nuclear materials, identifying isotopic and elemental "signatures" linking them to their last legal owner, production location, facility, or method of production is considered crucial for accountability and legal prosecution (8). As Pu used in weapons production has been alloyed with Ga, the elemental signature of interest is the Ga/Pu ratio because it is unique to each specific production facility and can be used to pinpoint the initial material processing facility. This method’s usefulness goes beyond nuclear forensics applications and can also be valuable for nuclear energy-producing industries. Decommissioning of the nuclear arsenal has prompted the U.S. to start dismantling numerous nuclear weapons and using the Pu in MOX (mixed-oxide) nuclear reactors (9). MOX fuel is the mixture of plutonium and uranium oxides and can be used in place of the traditional isotopically-enriched uranium dioxide fuels. However, the initial alloy has significant amounts of Ga, which is incompatible with practical MOX fuel applications even in small quantities, because it can migrate from the fuel into the zirconium (Zr) cladding, making that cladding more brittle and subject to corrosion (10–13). In order to down-blend these weapon pits for MOX applications, removal of Ga to approximately a 105 decontamination factor is needed to avoid these complications, requiring a final Ga concentration of 0.1 ppm within the solid oxide material (14). To use MOX fuel from down-blended Pu-bearing nuclear materials in nuclear reactors, there is a need for advancements in separations technologies. Several methods have been shown to effectively separate Ga and Pu in nuclear weapon pits for application as MOX fuel. For example, a separation of Pu and Ga by ion-exchange chromatography was proposed by DeMuth at Los Alamos National Laboratory that uses aqueous-based ion-exchange chromatography (15, 16). This method focused on large scale separations instead of isotopic analysis, making it less applicable for nuclear forensics applications (15, 16). A second method utilized high-temperature thermal removal of gallium for industrial MOX fuel production and again focused on large quantities necessary in the nuclear energy industry (17). These methods were not designed for, nor are they suitable for, the analytical investigation of gallium, plutonium, or other candidate actinides of interest for nuclear forensics of MOX fuel analysis. In order to address this critical need, methodology was developed at the University of Iowa that provided analysis of various actinide elements and gallium within plutonium nuclear materials, such as nuclear weapon pits and MOX nuclear fuels. Developing a dependable, fast, and accurate analysis method for stable gallium in Pu nuclear materials could be a powerful tool for both nuclear forensics applications and nuclear energy-producing industries. Combining the analysis of Ga with the determination of isotopes of the actinides Pu, U, Th, and Am would increase the efficiency of nuclear forensics investigations. 175

Using extraction chromatography paired with isotope dilution techniques, a complete radiochemical separation of Ga, Pu, Am, U, and Th was developed for nuclear forensics applications of Pu-bearing nuclear materials (18). Extraction chromatographic resins (TEVA and TRU Resins from Eichrom Technologies, LLC) were used to prepare elementally pure fractions for isotope analysis (Figure 1) (18). These chromatographic resins have organic extractants adhered to a solid substrate to allow for a chemical separation. Preparing and separating samples for analysis requires an accurate and precise method that determines the yield of Ga, even though the complex matrix of Pu-bearing nuclear materials is a hindrance to the analysis of gallium concentrations. A notable contribution to developing these methods was the use of a nuclear medical radioisotope, 68Ga, as an improved alternative to trace stable Ga yields in these separations (18). Innovative medical radioisotopes, such as 68Ga, are becoming available and affordable due to recent advancements in their production and isolation (19). Combining ideas from various areas of radiochemistry, such as nuclear medicine and nuclear forensics, can have substantial impacts on methodology development. Utilizing a medical radioisotope to provide stable gallium signatures in nuclear-forensics applications allows quick and accurate determination of Ga concentrations. The 68Ga nuclear medical radiotracer allows precise, stable gallium determination without interfering with the separation and quantification of other isotopes of radioactive actinides.

Figure 1. Tandem column arrangement of the TEVA/TRU developed for the separation and purification of stable Ga and radioactive isotopes of elements Th, Pu, Am, and U. The load and rinse solutions remove common ions and matrix interferences. The columns are then disassembled. Steps 4-6 (TEVA) may be run concurrently with steps 7-9 (TRU). Reproduced with permission from ref. (18). Copyright 2015 Springer. 176

Advancements in Separations and Nuclear Forensics Applications Methods developed at the University of Iowa provide the separation and analysis of Ga, Pu, U, Th, and Am for the purpose of nuclear forensics and nuclear fuel analysis. Further advancements are needed, however, that include certification of the method for use on various nuclear materials. This method’s applications could potentially be extended into many types of nuclear materials (MOX fuels, Pu-Ga alloys, nuclear wastes), but verification is a crucial step in the process. Accessing nuclear materials is not trivial due to nonproliferation security controls; accessing nuclear materials for peaceful research and development is similarly challenging. Performing this separation on digested nuclear fuels (both pre and post burn-up) as well as Pu-nuclear weapon pits would allow for a more complete understanding of this method’s strengths and limitations. Additional quantification of Ga by ICP-MS would also be a useful improvement as these instruments become more widely used for analysis of trace-level metals. The use of 68Ga as a radiotracer could be performed before the analysis of gallium by ICP-MS, and would likewise provide a precise quantification of the gallium present in a nuclear material. Certifying this method for ICP-MS analysis would widen the scope of sample analysis to include compounds like post-irradiation nuclear materials, which would have a dramatically more complex matrix, including fission products(e.g. 137Cs, 90Sr, 131I, 85Kr, 99Tc, 151Sm, 93Zr, etc.) that complicate analysis of the sample by radiotracers (20). Additional modifications to the method may be needed with the addition of fission products. For example, additional method verification would be required for various isobaric interferences, which could accompany fission products arising from post-burnup MOX fuels. Overall, the use of extraction chromatography with the addition of a medical radioisotope tracer helped to advance our methods regarding Pu-bearing nuclear materials. This work highlights the importance of radiochemistry training programs within the university setting, where advancements in one field can be applied to other fields and students involved in interdisciplinary programs can take advantage of a growing number of interesting isotopes used in a wide range of applications.

Concerns about Radioactivity in Liquid Waste from Hydraulic Fracturing When the University of Iowa program launched in 2012, concerns about naturally occurring radium (Ra) isotopes in Marcellus Shale produced fluids had just begun to surface within the scientific community (21–23). A group at the United States Geologic Survey (USGS) published a detailed study that illustrated the potential for 226Ra and 228Ra to concentrate in produced fluids and brines generated by hydraulic fracturing for natural gas in the Marcellus Shale formation (21). This report raised numerous questions about radioactivity in hydraulic fracturing flowback water and produced fluids, such as the presence of 177

naturally occurring radionuclides and environmental impacts of waste disposal. The USGS report was followed by a key study by Warner et al. (2013) that indicate the presence of high levels of Ra isotopes downstream of wastewater treatment plants handling Marcellus Shale produced fluids, suggesting that certain treatment plants were undertreating the radioactivity associated with these fluids (23). After the Warner et al. (2013) and USGS studies, it was clear that there were elevated levels of radioactivity in these fluids and that there was the potential for certain isotopes (226Ra and 228Ra) to contaminate the environment. Despite these two observations, there were no validated methods to detect and characterize radionuclides in hydraulic fracturing-produced fluids and brines. The problem presented by this lack of methodology was twofold: (1) without comprehensive methodology for alpha, beta, and gamma emitters, it was unclear whether other radioactive elements were present in the fluids, and (2) the radioactivity concentrations of reported Ra isotopes remained questionable.

Method Development By the summer of 2013, the U.S. Environmental Protection Agency (USEPA) had funded a project at the University of Iowa to develop and validate a method to detect gross alpha and beta particles in these fluids (24). Development of the gross alpha and beta method required a technique that could isolate key radionuclides, including U, Th, Pa, Ra, Pb, Bi, and Po isotopes. Numerous methods to isolate these radionuclides in drinking water exist (e.g. EPA 900.0 (25) & ASTM method #D7283) (25, 26); however, these methods are intended for waters with low concentrations of dissolved solids. Water samples with higher levels of dissolved solids are generally unsuitable for a “gross alpha and beta” method due to 1) the need to remove interfering stable elements that would absorb or attenuate the ionizing radiation, and 2) the challenges of separating the numerous radionuclides from stable elements. The sample of Marcellus Shale produced fluids provided to the University of Iowa had exceptionally high concentrations of dissolved solids (~278,000 mg/L) (27). Furthermore, the fluids contained high levels of barium (Ba) salts (>9,000 mg/kg) (28). The problem with Ba for radiochemical separations is that Ba and Ra have very similar chemistry and often co-elute or co-precipitate during chemical separations (28–31). The high levels of dissolved solids (in particular the high concentrations of Ba) suggested that a low-cost, rapid, practical separation of Ra from the Marcellus Shale fluids would not be possible. Given that Ra had been previously established as a key radionuclide in Marcellus Shale produced fluids and flowback (21, 32), this suggested that a single gross alpha/beta method was impractical. Thus, we had to split the sample up to measure key alpha emitters (238U, 234U, 232Th, 230Th, 228Th, 210Po) by alpha spectrometry and key beta emitters by gamma spectrometry (228Ac, 212Bi, 212Pb, 214Pb, 214Bi). Note that most key beta emitters of concern in this case also emitted gamma rays (ex: 210Pb; 46 keV) (33), or had short-lived daughters that could be measured by gamma spectrometry, which would allow for modeling their activity (ex: 226Ra released 214Bi and 214Pb). 178

Radiochemical Disequilibrium During the development of the isotope-specific method, we discovered several interesting radiochemical parameters. First, we observed that U levels in the produced fluids were exceptionally low, suggesting that U was insoluble. This is to be expected given that the Marcellus Shale is a reducing zone, meaning that U would be found in the +4 valence state and immobile (34). The second discovery was that 234U radioactivity concentrations exceeded 238U radioactivity concentrations. This result was surprising at first, given that a system as old as the Marcellus Shale is expected to be in secular equilibrium (meaning 238U and 234U should have the exact same radioactivity concentration; for further reading on radioactive equilibrium we recommend the following references) (35–37). The discrepancy in radioactivity levels is believed to be attributed to a phenomenon known as alpha recoil enrichment, which results in higher-than-expected concentrations of decay products (38, 39). The third discovery was that Ra activities, as measured by chemical separations, were greatly underestimated. We attributed this discovery to Ba interference during chemical separations, which is further discussed in the following references (27, 29). Lastly, and perhaps of greatest interest to us, was the apparent absence of 226Ra decay products from the fluids (40). Given the solubility of Pb in these fluids (27), we expected 210Pb to be detectable; however, models suggested that 210Pb was absent at the time of extraction and thereby insoluble in the fluids. More studies are needed to fully understand the chemical reasoning for the 210Pb insolubility in these fluids, but we currently hypothesize that this is caused by the partitioning of 222Rn gas (41). The absence of 210Po was expected due to the lack of its parent 210Pb.

Radium Decay Products As we were investigating the absence of 226Ra decay products in these fluids, Warner et al. (2013) documented high levels of Ra isotopes downstream of a wastewater treatment facility processing Marcellus Shale produced fluids but the study did not address Ra decay products.To further explore the environmental transport and mobility of Ra decay products, particularly 210Po and 210Pb, we embarked on a year-long field study near Mannington, West Virginia (42). Local citizens expressed concerns that the Northern West Virginia Water Treatment Facility was discharging undertreated Marcellus Shale waste into the Hibbs Reservoir. We later discovered through a Freedom of Information Act (FOIA) request that this facility only received coal wastes and that all documented discharges appeared in the range of issued permits (43). Regardless, we saw this as a unique opportunity to investigate the fate and transport of 226Ra decay products in the environment. After three sampling trips, we discovered that 210Po radioactivity concentrations were approximately 2.3 ± 0.4 (n=12) fold greater than the parent 226Ra (42). In one sample, the 210Po radioactivity concentration exceeded that of cleanup goals for surface sediments at USEPA CERCLA sites. We were hesitant to publish this result without an appropriate local control site, but many of the lakes in the surrounding area were impacted by oil and natural gas or coal extraction. Therefore, we sampled a lake at F.W. Kent State Park near 179

Oxford, Iowa that was of similar size, but had no documented oil and natural gas or coal extraction in its watershed to serve as a control site (44). We were shocked to discover that 210Po exceeded 226Ra concentrations by 2.8 ± 0.5 (n=5) at the Iowa site and that the activity concentration was very similar to what we observed at impacted Hibbs Reservoir in West Virginia. The only apparent source of 210Po for F.W. Kent Lake is background 222Rn, suggesting that 210Po accumulates in lakes in areas with high background 222Rn.

Drill Cuttings Although the bulk of discussion on radioactivity related to drilling activities in the Marcellus Shale has focused on liquid waste, we understood that all the radioactivity in liquid waste ultimately came from solids in the subsurface—i.e., the source rock necessarily containing the parents 238U and 232Th (40). As we were finishing our research on liquid wastes, reports were emerging in popular press outlets that radioactivity levels in solid wastes were potentially at elevated and actionable levels. Specifically, several news sources noted that radioactivity alarms at landfills were activated by trucks carrying drill cuttings (solid waste) (45). These radioactivity alarms were characterized by dose rates, but did not define the radiochemical profile of the drill cuttings. In early 2014, we began efforts to obtain drill cuttings so that we could characterize the disequilibrium status of these cuttings, but were not able to obtain materials until a drilling company directly supplied us with them in late 2015. During this time, Pennsylvania and West Virginia released studies that characterized certain aspects of these drill cuttings (gamma spectrometry-based results); however, the cuttings’ radiochemical profiles were not released due to a lack of methodology for alpha spectrometry. Thus we saw a critical need to develop methodologies for examining these materials and to distribute these methods to other researchers through the peerreview process. Due to the chemical characteristics and high organic content of the cuttings, we tailored these methods from techniques previously developed for asphalt. The method development process is described further in Eitrheim et al. (2016) (41). The most interesting aspect of the study is that 226Ra decay products were in disequilibrium with the supporting parent radionuclides. This continual trend of 238U/210Po disequilibrium in environmental radioactivity measurement raises several questions: Where is 210Po going?; How is it getting there?; What are the health implications of 210Po exposure? The following section of this chapter will explore these questions regarding the fate, transport, and health implications of 210Po exposure.

Future Directions of Research: Enhanced Understanding of 210Po in Natural Environments Polonium Radiochemistry and Sources Observations from our unconventional drilling studies suggest that 210Po poses an uncharacterized environmental health concern. Moreover, Po is among 180

the least characterized and poorly studied elements in the periodic table, in part due to the inherent challenges studying short-lived radionuclides. There are no stable isotopes of Po; all known isotopes are radioactive and have relatively short half-lives. The longest-lived isotope is the artificially produced 209Po (half-life 109 years) (46) and the longest lived naturally occurring isotope is 210Po (half-life of 138 days) (33). 210Po is the most important isotope of Po with respect to environmental nutrient cycling and public health. As the final radioactive product in the 238U decay series, 210Po is ubiquitous in nature and found along long-lived 238U series radionuclides. For example, 210Po is often observed in geologic and industrial sources of 238U (such as phosphate rocks, marine black shales, and uranium ores), 226Ra (phosphogypsum, hydraulic fracturing wastes, pipe scale), and 210Pb (coal fly ash, tobacco leaves, marine organisms) (34, 40, 47–50). Despite its high concentration in some environmental media (concentration factors greater than 10-fold over parent radionuclides), there are significant challenges with studying 210Po in the environment. The short half-life of 210Po results in very small quantities; masses of 210Po rarely exceed a few femtograms per gram or liter of environmental materials (51). For comparison, the half-life of 238U and 210Po is 4.5 billion years and 138 days, respectively, which means that with the same radioactivity concentration there will be 14 billion more atoms of 238U than 210Po. Laboratory-based studies on 210Po chemistry and biochemistry are thus largely based on trace-level masses due to the radiological hazards and challenges of obtaining gram quantities of 210Po. One gram of 210Po has a specific activity of 166 terabecquerels (TBq) per gram (~4500 Curies per gram), which corresponds to a lethal dose of 210Po of less than 10 micrograms. Furthermore, gram quantities of Po lack environmental relevance, i.e., trace-level experiments are more informative of environmental processes. Such trace levels of Po (10 femtograms or less) dictate that radiochemical methods must be used as masses are below the reasonable detection limits of conventional chemical characterization techniques (35). Polonium Water Chemistry Po has a unique aqueous chemistry that has variable redox chemistry that is not completely understood within natural systems. Its most stable valence state in the environment is tetravalent Po (Po(IV)), which forms strong hydroxide complexes to produce insoluble Po(OH)4. Under oxidizing conditions, Po is very particle-reactive and readily adsorbs to mineral surfaces (51). Reduction to divalent species can occur in natural waters, leading to the formation of more soluble forms and higher mobility. Organisms such as sulfate-reducing bacteria can take advantage of this redox process, as the Po(IV) likely serves as a reducible electron acceptor for either dissimilatory or assimilatory metabolic processes that produce the more mobile Po(II) oxidation state. However, studies have suggested that the mobilization of Po occurs in a narrow range of biogeochemical and geochemical conditions and in more reducing conditions resulting in the formation of insoluble Po(0) forms. The unique chemistry of Po is exploited during radiochemical separations and measurements, as Po is the only known naturally occurring radionuclide that autodeposits onto Ag, Ni, and Cu metals 181

under reducing, acidic conditions. While the basic understanding of Po chemistry has been explored, behavior in more complex environmental systems deserves additional investigations. Polonium Health Concerns Although very small mass quantities of 210Po are present in the environment, presents significant public health hazards. 210Po is known to bioaccumulate in the body and target the soft tissues, where it can deliver significant doses of high-energy alpha particles (52). Both acute exposures (ingestion of few micrograms) and extremely low-level environmental exposures (ingestion of few femtograms per day) are of concern (53, 54). As environmental health practitioners become more aware of the ubiquity and cancer hazards of 210Po, there are increasing concerns over possible routes for exposure (53). Although some of this attention may be attributed to the assassination of Alexander Litvinenko in 2006, acute radiotoxic exposures to 210Po are exceptionally rare (55). On the other hand, environmental and occupational exposures for 210Po are commonplace, with major pathways of exposure including inhalation (for examples, smoking tobacco products, 222Rn contaminated buildings) and ingestion (consumption of aquatic species, ingestion of groundwater) (47, 51, 56). Of all exposure pathways, ingestion of 210Po from contaminated groundwater is the least characterized and represents the most critical need for future research. 210Po

Polonium in Groundwater: A Critical Need for Research Contamination of groundwater by 210Po can occur through both anthropogenic and natural processes but the extent of the problem is unclear. 210Po may accumulate in groundwater due to activities associated with extractive industries, such as phosphate fertilizer production and unconventional drilling for natural gas (40, 57, 58). 210Po can also accumulate in groundwater due to natural processes as first reported by the United States Geologic Survey (USGS) when they discovered naturally high levels of 210Po in groundwater—the major source of water for local dairy operations in Fallon, NV (40, 57, 58). The extent of natural 210Po contamination of groundwater is currently unknown for multiple reasons. First, 210Po has been historically neglected in groundwater monitoring regimens, as it is not directly regulated by any federal agency (58). Second, most groundwater users are not required to test groundwater for any radionuclides. Lastly, as a pure alpha-emitter, 210Po cannot be measured by routine radioactivity dose assessment techniques (such as via a Geiger counter) (35). Despite these challenges, several agencies including the USGS and others in Minnesota and Iowa have begun to monitor for 210Po in select groundwater sampling sites with preliminary results anticipated very soon (59, 60). Regional scale environmental monitoring can provide important clues to the geochemistry and biogeochemistry of 210Po. Students in University of Iowa Radiochemistry program will take part in future studies on the fate and transport of 210Po and with the help of local, regional, and international collaborations to further our understanding of the chronic health implications of exposure to low doses of this material. 182

Protactinium: The Element of Witchcraft Protactinium (Pa) is a naturally occurring actinide element, yet its chemistry continues to mystify scientists. Of the naturally occurring elements, Pa was one of the last seven to be discovered, largely in part to a general misunderstanding of its chemistry (61, 62). The chemistry of Pa has been described as puzzling (63), peculiar (64), mysterious (65), and even witchcraft (65), and much of these descriptions still apply today (65). In the late 1950s, a researcher in the National Academy of Sciences, Harold Kirby, postulated that in that decade, much of the mystery surrounding the chemistry of Pa “will be eliminated through the development of a convenient radiotracer (65).” Yet our understanding of its chemistry is still shrouded in mystery and ongoing concerted efforts are only beginning to understand the basic chemical and physical properties of Pa. At the University of Iowa, investigations were driven to establish a better understanding of the chemistry of Pa and apply this new knowledge to applications in nuclear forensics, nuclear energy, and environmental radiochemistry.

The Chemistry of Protactinium and Applications The challenge of understanding the chemistry of Pa lies in its unique chemical and nuclear properties. Interestingly, Pa exists in a critical location within the actinide series (An), where an energy crossover occurs between the 6d and 5f orbitals and computational studies have suggested that these states will be nearly degenerate (66). This leads to distinct bonding characteristics that are distinctive to Pa and chemistry that deviates drastically from other early series actinides like Th, U, Np and Pu (64). A recent study investigating the enthalpy of isomerization of AnO2(H2O)+ to AnO(OH)2+ (for An: Pa-Pu) revealed that this process is distinctively low in energy for Pa in contrast to the other An elements where the formation of AnO2(H2O)+ is considerably more energetically favorable (66). Within this context, the research at the University of Iowa aimed to develop a detailed understanding of the trace-level chemical properties of Pa (67). We focused on the unique properties of Pa in trace-level concentrations to improve our understanding of the chemical behavior and bonding characteristics of Pa (67–69). Pa is not widely used as a radiotracer due to its high radioactivity, low abundance, and poorly understood chemistry, so its primary geochemical application is the assessment of 231Pa/235U ratios (Figure 2) (68, 70–73). Small amounts of 235U occur naturally in geologic deposits where it will decay via alpha-particle emission to 231Pa and the isotopic ratio can be used to determine the time that has elapsed since the initial fractionation (70). While there are numerous dating methods available to geologists to date a variety of sample types, the 230Th/238U and 231Pa/235U isotopic ratios are viable options for determining the precise ages of quaternary materials (events occurring up to 250,000 years ago or ~7 half-lives of 231Pa). The 231Pa/235U chronometer is more sensitive with samples of younger ages due to its shorter half-life compared to 230Th and can be used for more precise measurements within this time period. Most notably, 231Pa/235U has been used to date a human skull from Qafzeh, Israel and Neanderthal bones in 183

Israel (74), which helped confirm that humans and Neanderthals did not coexist in the Levant during the same period (75). Furthermore, these two isotopic ratios can be used together to determine sedimentation rates or provide extremely precise ages (76–78).

Figure 2. Radioactive ingrowth of 231Pa to secular equilibrium with 235U and to secular equilibrium with 238U. This figure does not take into account the isotopic abundance of the U isotopes.

230Th

The 231Pa/235U isotopic ratio has also proven itself to be valuable for the agedating of U-special nuclear material for nuclear forensics analysis (79, 80). For these samples, the materials are relatively young (4 M HCl), Pa is selectively extracted into the organic phase whereas the other actinides (Th, U, Np, and Am) remain in the aqueous phase. Slope analysis experiments with respect to the extractant alcohol and the ligand (Cland NO3-) showed a 2:1 relationship between the extractant molecules and the central Pa metal. Furthermore, these studies revealed the approximate speciation of Pa with respect to the ligands. In HCl, the extracted species of Pa appears to be a hexachloro complex, PaCl6-, suggesting that the extraction occurs via the formation of HPaCl6 complex acid and stabilized by the H-bonding network of the alcohols of the hydroxyl groups. In HNO3, Pa speciation is less obvious because NO3- can act as a mono- or bidentate ligand, and slope analysis cannot distinguish between the two forms. The analysis suggests the formation of a 2:1 Pa-nitrate species, but the additional coordinating species (hydroxyl, water, oxo) are not known and must be verified with additional spectroscopic or scattering techniques (84).

Figure 4. Separation methods for the separation of Pa from Np for the tracer preparation, and separation of Th, Pa, and U for isotope dilution alpha spectrometry.

Future Work with Pa Pa was discovered over 100 years ago, but our current understanding of Pa chemistry remains at an introductory level. Current research is focusing on the fundamental chemistry of aqueous complexes (85, 86) and rapid methods 187

for isotopic analysis (73, 79), but the overall progress is slow. Only a few investigators around the world are actively engaged in Pa research. In a recent perspective, Wilson highlights the importance of investigating Pa as an avenue to understanding the chemical dependencies of 6d and 5f electronic structures, but notes, “future contributions to (Pa) chemistry may come from where Meitner and Hahn first found it, in silico” (64).

Conclusions As discussed above, the forecasted shortage of radiochemists and radiochemistry graduate programs led to the development of curriculum and expertise at University of Iowa. The first cohort of radiochemistry students focused on Ga and Pu with applications in MOX fuels, naturally occurring uranium-series radionuclides in the environment, and fundamental Pa chemistry. Advancing the complete separation of gallium and actinides in nuclear materials for nuclear forensics applications was accomplished yet developments regarding ICP-MS and certification of these methods is still needed. Research in uranium-series radionuclides highlighted the importance of sound methodology and comprehensive analysis of decay products, including 210Po. From our observations of 210Po in aqueous environments, we have begun to probe deeper into the geochemical cycling of 210Po in the Upper Midwest. Our understanding of Pa has also broadened, potentially allowing better applications for isolation, purification, and separation of Pa for isotopic analysis with applications ranging from nuclear forensics to geochronology. Still a better understanding of fundamental Pa chemistry is needed, which may be aided in the future by better access to this element. Overall, the University of Iowa Radiochemistry program has contributed to the understanding of these various elements. There is high hope that radiochemistry advancements will continue to contribute to future discoveries regarding elements and their chemistry.

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36. Choppin, G. R. R., J.; Liljenzin, J-O.; Ekberg, C. Radiochemistry and Nuclear Chemistry, 4th ed.; Academic Press, Elsevier: Amsterdam, The Netherlands, 2013. 37. Peate, D. W.; Hawkesworth, C. J. U series disequilibria: Insights into mantle melting and the timescales of magma differentiation. Rev. Geophys. 2005, 43 (1), RG1003. 38. Fleischer, R. L. Isotopic disequilibrium of uranium: alpha-recoil damage and preferential solution effects. Science 1980, 207 (4434), 979–81. 39. Osmond, J. K.; Cowart, J. B.; Ivanovich, M. Uranium isotopic disequilibrium in ground water as an indicator of anomalies. Int. J. Appl. Radiat. Isot. 1983, 34 (1), 283–308. 40. Nelson, A. W.; Eitrheim, E. S.; Knight, A. W.; May, D.; Mehrhoff, M. A.; Shannon, R.; Litman, R.; Burnett, W. C.; Forbes, T. Z.; Schultz, M. K. Understanding the radioactive ingrowth and decay of naturally occurring radioactive materials in the environment: an analysis of produced fluids from the Marcellus Shale. Environ. Health Perspect. 2015, 123 (7), 689. 41. Eitrheim, E. S.; May, D.; Forbes, T. Z.; Nelson, A. W. Disequilibrium of Naturally Occurring Radioactive Materials (NORM) in Drill Cuttings from a Horizontal Drilling Operation. Environ. Sci. Technol. Lett. 2016, 3 (12), 425–429. 42. Nelson, A. W.; Eitrheim, E. S.; Knight, A. W.; May, D.; Wichman, M. D.; Forbes, T. Z.; Schultz, M. K. Polonium-210 accumulates in a lake receiving coal mine discharges—anthropogenic or natural? J. Environ. Radioact. 2017, 167, 211–221. 43. West Virginia Department of Environmental Protection, Freedom of Information (FOIA) request. West Virginia Department of Environmental Protection, Ed. 2016. 44. Freidhoff, B.; Nelson, A. W. Personal communication, 2015. 45. McMahon, J. Fracking Truck Sets Off Radiation Alarm At Landfill. https://www.forbes.com/sites/jeffmcmahon/2013/04/24/fracking-truck-setsoff-radiation-alarm-at-landfill/#4af56ee036e3 (accessed March 1, 2017). 46. Colle, R.; Laureano-Perez, L.; Outola, I. A note on the half-life of 209Po. Appl. Radiat. Isot. 2007, 65 (6), 728–30. 47. Shannon, L. V.; Cherry, R. D.; Orren, M. J. Polonium-210 and lead-210 in the marine environment. Geochim. Cosmochim. Acta. 1970, 34 (6), 701–711. 48. Tso, T. C.; Harley, N.; Alexander, L. T. Source of Lead-210 and Polonium210 in Tobacco. Science 1966, 153 (3738), 880–882. 49. Hull, C. D.; Burnett, W. C. Radiochemistry of Florida phosphogypsum. J. Environ. Radioact. 1996, 32 (3), 213–238. 50. Lauer, N. E.; Hower, J. C.; Hsu-Kim, H.; Taggart, R. K.; Vengosh, A. Naturally Occurring Radioactive Materials in Coals and Coal Combustion Residuals in the United States. Environ. Sci. Technol. 2015, 49 (18), 11227–33. 51. Persson, B. R. R.; Holm, E. Polonium-210 and lead-210 in the terrestrial environment: a historical review. J. Environ. Radioact. 2011, 102 (5), 420–429. 191

52. Ansoborlo, E.; Berard, P.; Den Auwer, C.; Leggett, R.; Menetrier, F.; Younes, A.; Montavon, G.; Moisy, P. Review of chemical and radiotoxicological properties of polonium for internal contamination purposes. Chem. Res. Toxicol. 2012, 25 (8), 1551–64. 53. Seiler, R. L.; Wiemels, J. L. Occurrence of (210)Po and Biological Effects of Low-Level Exposure: The Need for Research. Environ. Health Perspect. 2012, 120 (9), 1230–1237. 54. Scott, B. R. Health Risk Evaluations for Ingestion Exposure of Humans to Polonium-210. Dose-Response 2007, 5 (2), 94–122. 55. Maguire, H.; Fraser, G.; Croft, J.; Bailey, M.; Tattersall, P.; Morrey, M.; Turbitt, D.; Ruggles, R.; Bishop, L.; Giraudon, I.; Walsh, B.; Evans, B.; Morgan, O.; Clark, M.; Lightfoot, N.; Gilmour, R.; Gross, R.; Cox, R.; Troop, P. Assessing public health risk in the London polonium-210 incident, 2006. Public Health 2010, 124 (6), 313–8. 56. McDonald, P.; Cook, G. T.; Baxter, M. S. Natural and Artificial Radioactivity in Coastal Regions of UK. In Radionuclides in the Study of Marine Processes; Kershaw, P. J., Woodhead, D. S., Eds.; Elsevier Science Publishing Co., Inc.: New York, 1991; pp 329–339. 57. Harada, K.; Burnett, W. C.; LaRock, P. A.; Cowart, J. B. Polonium in Florida groundwater and its possible relationship to the sulfur cycle and bacteria. Geochim. Cosmochim. Acta 1989, 53 (1), 143–150. 58. Nelson, A. W.; Knight, A. W.; Eitrheim, E. S.; Schultz, M. K. Monitoring radionuclides in subsurface drinking water sources near unconventional drilling operations: a pilot study. J. Environ. Radioact. 2015, 142, 24–28. 59. May, D.; Schultz, M. K. Natural Radioactivity in Drinking Water Resources in Iowa – a Pilot Study; Radiobioassay and Radioanalytical Conference, RRMC, Iowa City, Iowa, Oct 25−30, 2015. 60. Minnesota Department of Health. Radionuclide Assessment Project. http://www.health.state.mn.us/divs/eh/risk/guidance/dwec/radionproj.html (accessed February 15, 2017). 61. Morss, L. R. The Chemistry of the Actinide and Transactinide Elements; Springer: Dordretch, The Netherlands, 2010; Vol. 1. 62. Pal’shin, E. S.; Myasoedov, B. F.; Davydov, A. V. Analytical Chemistry of Protactinium; Ann Arbor-Humphrey Science Publishers, Inc: Ann Arbor, Michigan, 1970. 63. Siboulet, B.; Marsden, C. J.; Vitorge, P. What Can Quantum Chemistry Tell Us About Pa(v) Hydration and Hydrolysis? New J. Chem. 2008, 32 (12), 2080–2094. 64. Wilson, R. E. Peculiar Protactinium. Nat. Chem. 2012, 4 (7), 586–586. 65. Kirby, H. W. The Radiochemistry of Protactinium, (National Academy of Sciences National Research Council); Nuclear Series (NAS-NS 3016); National Academy of Sciences - National Research Council: Washington, DC, 1959. 66. Kaltsoyannis, N. Covalency Hinders AnO2(H2O)+ → AnO(OH)2+ Isomerisation (An = Pa-Pu). Dalton Trans. 2016, 45 (7), 3158–3162. 67. Knight, A. W.; Eitrheim, E. S.; Nelson, A. W.; Peterson, M.; McAlister, D.; Forbes, T. Z.; Schultz, M. K. Trace-Level Extraction Behavior of Actinide 192

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Elements by Aliphatic Alcohol Extractants in Mineral Acids: Insights into the Trace Solution Chemistry of Protactinium. Solvent Extr. Ion Exch. 2016, 34 (6), 509–521. Knight, A. W.; Eitrheim, E. S.; Nelson, A. W.; Nelson, S.; Schultz, M. K. A Simple-Rapid Method to Separate Uranium, Thorium, and Protactinium for U-series Age-Dating of Materials. J. Environ. Radioact. 2014, 134 (0), 66–74. Knight, A. W. N. A. W.; Eitrheim, E. S.; Forbes, T. Z.; Schultz, M. K. A Chromatographic Separation of Neptunium and Protactinium Using 1-octanol Impregnated onto a Solid Phase Support. J. Radioanal. Nucl. Chem. 2015 (307), 59–67. Bourdon, B.; Turner, S.; Henderson, G. M.; Lundstrom, C. C. Introduction to U-series Geochemistry. Rev. Mineral. Geochem. 2003, 52, 1–21. Peate, D. W.; Hawkesworth, C. J. U-Series Disequilibria: Insights into Mantle Melting and the Timescales of Magma Differentiation. Rev. Geophys. 2005, 43 (1), 1–43. Pickett, D. A.; Murrell, M. T.; Williams, R. W. Determination of Femtogram Quantities of Protoactinium in Geologic Samples by Thermal Ionization Mass-Spectrometry. Anal. Chem. 1994, 66 (7), 1044–1049. Regelous, M.; Turner, S. P.; Elliott, T. R.; Rostami, K.; Hawkesworth, C. J. Measurement of Femtogram Quantities of Protactinium in Silicate Rock Samples by Multicollector Inductively Coupled Plasma Mass Spectrometry. Anal. Chem. 2004, 76 (13), 3584–3589. Yokoyama, Y.; Falgueres, C.; deLumley, M. A. Direct dating of a Qafzeh Proto-Cro-Magnon skull by non destructive gamma-ray spectrometry. Cr. Acad. Sci. Ii. A. 1997, 324 (9), 773–779. Schwarcz, H. P.; Simpson, J. J.; Stringer, C. B. Neanderthal skeleton from Tabun: U-series date by gamma-ray spectrometry. J. Hum. Evol. 1998, 34 (3), A18–A19. Koornneef, J. M.; Stracke, A.; Aciego, S.; Reubi, O.; Bourdon, B. A New Method for U-Th-Pa-Ra Separation and Accurate Measurement of U-234, Th-230, Pa-231, Ra-226 Disequilibria in Volcanic Rocks by MC-ICP-MS. Chem. Geol. 2010, 277 (1-2), 30–41. Richards, D. A.; Dorale, J. A. Uranium-Series Chronology and Environmental Applications of Speleothems. Uranium-Series Geochemistry. 2003, 52, 407–460. Sims, K. W. W.; Pichat, S.; Reagan, M. K.; Kyle, P. R.; Dulaiova, H.; Dunbar, N. W.; Prytulak, J.; Sawyer, G.; Layne, G. D.; Blichert-Toft, J.; Gauthier, P. J.; Charette, M. A.; Elliott, T. R. On the Time Scales of Magma Genesis, Melt Evolution, Crystal Growth Rates and Magma Degassing in the Erebus Volcano Magmatic System Using the U-238, U-235 and Th-232 Decay Series. J. Petrol. 2013, 54 (2), 235–271. Eppich, G. R.; Williams, R. W.; Gaffney, A. M.; Schorzman, K. C. U-235/Pa231 Age Dating of Uranium Materials for Nuclear Forensic Investigations. J. Anal. Atom. Spectrom. 2013, 28 (5), 666–674.

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80. Morgenstern, A.; Apostolidis, C.; Mayer, K. Age Determination of Highly Enriched Uranium: Separation and Analysis of 231Pa. Anal. Chem. 2002, 74 (21), 5513–5516. 81. United States Nuclear Regulatory Commission. U. Storage of Spent Nuclear Fuel. https://www.nrc.gov/waste/spent-fuel-storage.html (accessed March 1, 2017). 82. United States Government Accountability Office. U. Disposal of High-Level Nuclear Waste. http://www.gao.gov/key_issues/ disposal_of_highlevel_nuclear_waste/issue_summary (accessed March 1, 2017). 83. Sill, C. W. Preparation of Protactinium-233 Tracer. Anal. Chem. 1966, 38 (11), 1458. 84. Le Naour, C.; Trubert, D.; Di Giandomenico, M. V.; Fillaux, C.; Den Auwer, C.; Moisy, P.; Hennig, C. First Structural Characterization of a Protactinium(V) Single Oxo Bond in Aqueous Media. Inorg. Chem. 2005, 44 (25), 9542–9546. 85. De Sio, S. M.; Wilson, R. E. Structural and Spectroscopic Studies of Fluoroprotactinates. Inorg. Chem. 2014, 53 (3), 1750–1755. 86. De Sio, S. M.; Wilson, R. E. EXAFS Study of the Speciation of Protactinium(V) in Aqueous Hydrofluoric Acid Solutions. Inorg. Chem. 2014, 53 (23), 12643–12649.

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Chapter 10

Element 118: Teaching A New Element to New Students Justin Pothoof,1 Grace Nguyen,1 Dawn Archey,2 E. Prasad Venugopal,1 and Mark A. Benvenuto*,1 1Department

of Chemistry and Biochemistry, University of Detroit Mercy, 4001 W. McNichols Road, Detroit, Michigan 48221-3038, United States 2Department of Mathematics and Software Engineering, University of Detroit Mercy, 4001 W. McNichols Road, Detroit, Michigan 48221-3038, United States *E-mail: [email protected].

The paper discusses how the discovery and naming of the newest elements can be brought into the discussion of the general chemistry class, with an emphasis on Element 118. These two exercises challenge students to use their knowledge of general chemistry as well as mathematics and physics to explain the existence of element 118, and to predict some of its physical properties.

Introduction To the youthful eyes of new students, the periodic table looks like a hydra of numbers and symbols. Unfortunately for those with that point of view, this monster is indispensable to the understanding of science and the reactions that occur on a daily basis in our bodies, on the Earth, and throughout the universe. Therefore, an understanding of the elements, their properties, and their relationship to other elements is stressed in introductory chemistry courses. Perhaps though there are better methods in which to explore, understand, and interpret this monster of science. The current style of teaching the periodic table, elements, and their nuclei gives a skewed vision of their true existence. For instance, Bohr atom images have continuously been used to teach the idea of an atom to students since it was first constructed. The use of this image inaccurately teaches students that electrons © 2017 American Chemical Society

are the same size as protons and neutrons and that they exist in a continuous orbit around the nucleus, which is composed of the neutrons and protons of an atom and account for 99% of the mass in an atom. More accurately, electrons exist within a cloud, but their location cannot be accurately determined. Realistically, what a student is focused on is memorizing elemental symbols and understanding the correlation of atomic number with proton number for their next quiz. It would be fruitful for the method of teaching to evolve with the periodic table. Students are told that the nucleus contains the protons and neutrons, as stated before, and simply accept the fact rather than interpret it. The question that can be asked is: “What does a nucleus look like?” and to have an image of that, the understanding of the forces at the nuclear level should first be known (1–4). Simply stated, the strong force, or nuclear force, serves to hold the nucleus together; it is an attractive force between protons and neutrons, neutrons and neutrons, and, strangely enough, protons and protons. The electrostatic force acts in the opposite direction as the positive charged protons repel one another. Ideally, the nucleus would be shaped in a way that separates protons away from each other enough to minimize electrostatic forces and maximizes nuclear forces. The periodic table is constantly expanding, with four recently-named newly synthesized elements: nihonium (Nh, 113), muscovium (Mc 115), tennessine (Ts, 117) and oganesson (Og, 118) (5–7). The evolution of the table can with relative ease be correlated to the importance of understanding elemental functionality within groups on the table. Seeing where the elements are placed on the table gives indications as to what characteristics and properties they possess, or can at least give a prediction of them. Ideally, the time in which we are living is intriguing enough to make students contemplate the nature of these newly discovered elements in some other-than-class time. This can also be considered wishful thinking on the part of educators, though. Yet teachers can try to challenge young minds by presenting them with this prompt: what does a cup of element 118 look like? This question has of late created a unique opportunity for students to use their knowledge and creativity to produce responses.

The Traditional Presentation of Periodicity The periodic table has certainly been taught to students in one form or another for over a century. For the student seeing this for the first time, it is fair to say there is a lot of memorization involved. The nomenclature of the elements does not follow any system, at least not in any western language, with several names going back to ancient times, a large number of them having been assigned in the eighteenth and nineteenth century, and the transuranic elements and heavier being assigned after the Second World War. Curiously, an examination of the periodic table in Mandarin Chinese does show something of a systematic approach to it, with the wind character being used in the names of all elements that are gases at ambient temperature, and the character for gold in the names of the elements that are metals (this may be because the first transliterations of a growing periodic table into Chinese were undertaken in the nineteenth century as a collaboration between scientists within the Chinese Empire and western missionaries. 196

In whatever way the periodic table is now presented and learned by students and others, memorization has a primary role, and discussions of periodicity often include some explanation of nuclear particles, protons and neutrons, as tiny balls or hard spheres with the protons being positively charged, and the neutrons possessing mass but no charge. Physicists debate the nature and “look” of the nucleus, but most teachers recognize that it is easy to conceptualize a nucleus as some aggregation of small, hard spheres surrounded by shells or clouds of electrons. The continued build-up of nuclear particles, and their orbiting electrons with negative charges to balance the positively charged nucleus, aids in understanding the table in terms of trends in the columns and rows of different elements.

Student Exercise: A Greater Understanding of the Nucleus Nuclear chemistry is not often stressed in the undergraduate chemistry curriculum. Rather, in the general chemistry and inorganic chemistry classes, an understanding of nucleophilicity, electron movement, and loss, gain, and sharing of electrons is heavily emphasized. While some alpha and beta particle interactions may be covered in a general chemistry classroom setting, the understanding of nuclear transmutations tends to be limited to this, an exercise that is really little more than counting nucleons. So the idea of an exercise in which students developed a model of a 118 proton nucleus, using nothing more than geometry and geometric shapes, has an appeal. It can deepen student understanding of how nucleons can come together, and how they can do so with minimal proton – proton interactions. Element 118 was the element of focus in this exercise because it is one of the most recent, least understood, and possibly controversial elements to be synthesized (since less than 10 atoms were detected). Element 118, which has recently been named “oganneson,” was first produced by a team of Russian and American scientists in 2002 at the Joint Institute of Nuclear Research in Dubna, Moscow, Russia; already world renowned for the successful synthesis of nobelium, copernicium, and elements 113, 114, 115, 116, and 117. Element 118 was produced by collisions of californium-249 atoms and calcium-48 ions. It was found to be relatively unstable with a half-life of 0.89 milliseconds. Its rate of decay helps illustrate the idea of an island of stability when compared to a period of time where the half-life of an isotope spikes to 0.16 seconds (flerovium), before dropping back to 1.9 milliseconds (copernicium). Using nothing but a basic geometric model, students were asked to construct a nucleus containing 118 protons, starting with a proton as a central point from which a three-dimensional shape can grow, as well as starting with a neutron as the central point. Using the first starting point, and placing neutrons at the x, y, and z coordinates results in an octahedron with neutrons at the outer six points. This shape has eight faces, which are the next logical positions at which to place nucleons, which would now be protons. Figure 1 shows an example of this. Placing a proton on each of the eight faces of the octahedron results in the outermost points being a cube. The now larger cube has six faces, which can again 197

have neutrons placed on them. This is also shown in Figure 1, and shows the first layer of a repeating pattern, one that can be represented: 1-6-8-6-8… and that can be repeated until 118 protons are in the layers of this growing shape. When students are presented with this exercise, they do eventually come to the realization that just under 15 layers of octahedron – cube repeating layers will eventually occupy all 118 protons, if the starting point is a proton, and the points of each cube are protons. If the starting point is a neutron, the points of each octahedron are occupied by protons, and with only six points at each layer of protons, it will require almost 20 layers to place all 118 protons. Since this is an exercise to aid students in understanding atomic nuclei in some greater depth, it does not matter if the start point for the shape is a proton or a neutron. The idea is that students construct a model of a nucleus that utilizes or occupies the required number of protons. The two different possible outcomes actually encourage class discussion about what a nucleus actually looks like (a concept that is still debatable among physicists). Class discussions about this exercise routinely pull analogies from the macroscopic world. Students can “see” this model better when they think of how fruit is stacked in a grocery store, or how cannon balls are stacked at a war memorial or national park. It is emphasized that these analogies are linked to a belief that protons and neutrons exist as tiny, hard spheres, which is not something upon which all physicists and chemists agree. Yet it provides students with a starting point to a greater understanding of the properties of a heavy atomic nucleus. While students do not tend to turn in a graphic image of a repeating point-octahedron-cube figure that expands to include 15 layers, many of them do produce diagrams that look much like larger versions of Figure 1. In some cases the assignment is hand written, in others some graphics program has been used. Routinely though, an explanation is provided along with the graphic, indicating how many repeat layers of octahedra and cubes are required to reach a total of 118 protons.

Figure 1. Constructing A 118 Proton Nucleus Starting from A Central Proton.

Student Exercise: Predicting Element 118 Properties Another idea presented to students is that the elements exist within groups or families on the periodic table, and within these families are shared properties. Many students are told this and accept it as true, but are never given the opportunity to apply and test this idea. 198

The reports and press coverage of the four new elements include Element 118, but because of the small number of atoms made, and the short half-life of them, there has been no chemical or physical behavior to report, as of the present. Thus, a student exercise was developed whereby certain physical properties of the known noble gases could be used to make predictions of the properties of Element 118. These are comparisons of: 1. 2. 3. 4.

Melting points Boiling points Molar volumes Atomic densities

Students were asked to find the known values for each, then construct graphs and extrapolate to estimate the property for Element 118. No restrictions were placed on what type of graph was to be constructed, since line graphs, scatter plots, and bar graphs all provide similar enough representations of the physical property in question that it is easy to see. Figure 2 and 3 show the first two of these four physical properties, with the temperature, in Kelvin, as the y-axis.

Figure 2. Melting Points for Elements in Column 18 of the Periodic Table (y-axis is degree K).

Both of the physical properties graphed in Figures 2 and 3 are well known to students, and they come to a freshmen-level chemistry class with a familiarity of temperatures, knowing for instance the melting and boiling point of water, the temperature of a winter and summer day in their locale, and the ambient temperature of a household. This graphing exercise allows them to predict what the melting and boiling point of Element 118 might be, and even to come to the 199

realization that if enough of it were ever made, it could in theory be cooled to the point of being what might be called a noble liquid or a noble solid.

Figure 3. Boiling Points for Elements in Column 18 of the Periodic Table (y-axid is degree K).

Graphing molar volumes and molar densities reinforce an understanding of the physical properties of these elements, including Element 118. Although it is perhaps obvious that students have less basic understanding of molar volume of different materials, or how dense different gases are (a statement that may also be true for the scientists and the general public, as well), the exercise is still a useful one for training students how to make a predictive comparison.

Conclusions These two student exercises were initiated only in 2016, but appear to have proven that even in the freshmen-level general chemistry class, a deeper understanding of the periodic table and the atomic nucleus can be taught than that which comes simply from rote memorization (1, 2). Thus, an understanding of the table becomes more than just how to memorize pieces of the table. Our exercise in using geometric shapes to construct a 118 proton nucleus makes connections between the concepts learned in a chemistry class and those in mathematics classes. As well, the fact that there is more than one “correct” answer to such an exercise opens a larger discussion of what a nucleus actually looks like. Our second exercise uses tools students are already familiar with – graphing – to teach how properties of an element can be predicted. Overall, both exercises bring relevance to the teaching of the periodic table, and can be easily incorporated into chemistry, mathematics, and physics classes. 200

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Battersby, S. Pear-shaped nucleus boosts search for new physics. http://www.nature.com/news/pear-shaped-nucleus-boosts-search-for-newphysics-1.12952 (accessed 27 February 2017). Cook, N. D. Models of the Atomic Nucleus, 2nd ed.; Springer: New York, 2010. Gaffney, L. P.; Butler, P. A.; Scheck, M.; Hayes, A. B.; Wenander, F. Studies of pear-shaped nuclei using accelerated radioactive beams. Nature. 2013, 497 (7448), 199–204. Synopsis: Nucleus is Surprisingly Pear Shaped, APS Physics. https:// physics.aps.org/synopsis-for/10.1103/PhysRevLett.116.112503 (accessed 27 February 2017). IUPAC is naming the four new elements nihoonium, moscovium, tennessine, and oganesson. https://iupac.org/iupac-is-naming-the-four-new-elementsnihonium-moscovium-tennessine-and-oganesson/ (accessed 12 February 2016). The element maker: Yuri Oganessian and the story of the new elements. Chemistry World 2017, 1, 22–32. ExtremeTech, Names for four new superheavy elements formally proposed. https://www.extremetech.com/extreme/229868-names-for-four-newsuperheavy-elements-formally-proposed (accessed 27 February 2017).

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Chapter 11

Chemicals, Their Element Names, and Their Place in Society. Also Known as “Why Did I Choose To Name My Organization Palladium Science Academy?” George W. Ruger Jr.* Founder, Palladium Science Academy, Modena, New York 12548, United States *E-mail: [email protected].

While contemplating names for my new start-up company, Palladium Science Academy, I considered my connection to Green Chemistry education along with my enthusiasm for inspiring children and adults to learn more about science. Over the years, some historical practices such as alchemy and more recent trends such as Chemical-Free marketing helped shape my view of the public perception of chemistry. Sometimes that perception is accurate, but other times — such as with Chemical-Free marketing — it is not. Palladium Science Academy was created to promote valuable education while being entertaining at the same time. A fast journey of scientific topics of interest will be presented.

© 2017 American Chemical Society

Introduction and Historical Perspective Throughout history, people have been inspired to learn about the world around us. We have studied nature, the stars, our fellow man, anything that would-be scientists and philosophers thought might help to explain who we are and what are we doing in the universe. I will focus on two examples that help to showcase how we have arrived at our understanding of the sciences. First we will be given a brief view of alchemy, as told through art and a historical perspective provided by the Chemical Heritage Foundation (CHF) in Philadelphia, PA. This will be followed by a review of chemistry sets available over the decades, which are now in CHF’s archives. Attending their facility at my first ACS National Meeting in 2008, I have gained an appreciation for many historical aspects of the sciences that are often overlooked in today’s textbooks. Alchemy as a science is often unappreciated. Traditionally, we think of alchemists as people who searched for a way to turn lead into gold, or other experiments that seem unrealistic. Alchemy was not historically considered a true science by many, and even today it is mostly overlooked. Artists and poets created works of art around alchemists and their laboratories, although unfortunately, these were not always the most flattering portrayals. In Figure 1, Rijcke-Armoede (Rich Poverty) shows an alchemist toiling away in his lab, while his family in the background looks quite distraught. From the title and the objects in the picture, the artist appears to convey that the alchemist is neglecting his family.

Figure 1. Rijcke-Armoede (Rich Poverty) 1632 Adriaen van de Venne. Courtesy of Fisher Scientific International, Chemical Heritage Foundation Collections. 204

Figure 2 is The Bald Alchemist with His Assistant 17th century. Caption provided by Chemical Heritage Foundation “The Bald Headed Alchemist, after David Teniers II, 17th century, oil on panel. The subject of this painting is an elderly alchemist at his furnace stirring a mixture in a crucible while also reading a manuscript. His young assistant patiently awaits further instructions while squeezing the bellows to keep the fire burning. The tools alchemists needed to experiment are depicted, from alchemical glassware and distillation apparatus to manuscripts.” In Figure 3, An Alchemist and His Laboratory, late 17th century, we see an alchemist looking over a manuscript and taking notes. The laboratory appears well maintained. The table is tidy. The objects on the back shelf are organized. There is a skull on the table and what appears to be an animal cage hanging in the background. Some artwork of this era features biology elements such as these.

Figure 2. The Bald Alchemist with his Assistant, 17th Century. After David Teniers II. Courtesy of Fisher Scientific International, Chemical Heritage Foundation Collections. 205

Figure 3. An Alchemist in His Laboratory, late 17th Century. Follower of Gerard Dou. Gift of Roy Eddleman, Chemical Heritage Foundation Collections.

In Figure 4, An Alchemist At Work, 17th century, the artist paints a very different picture. We still see a manuscript and items used by an alchemist, including clay pots and other apparatus. However, the artist also depicts broken clay pot pieces. Items are piled up on the floor and even the shelves are not well kept. As are many pieces of artwork of this period, the alchemists are showed in a poor light. They are seen as not well organized and they are often depicted as being poor. Some artwork shows straw laying around the laboratories along with many pieces of broken equipment, symbolizing poverty. 206

Figure 4. An Alchemist at Work, 17th century. Mattheus van Helmont, Flemish. Gift of Roy Eddleman Chemical Heritage Foundation Collections.

While alchemy may not have achieved its goal of turning base metals into something valuable such as gold, there were a number of positive accomplishments. The way that the alchemists attempted to carefully study their materials and take intricate notes can be seen as the beginning of the way that scientists still work today. Documentation and replication of experiments are seen as vital to the work of modern science. These traits can be seen in the texts and even the artwork of the period. Without these humble beginnings, the current scientific process might never have taken shape. Another factor shaping our view of chemistry over the decades can be seen in the toys made available to the public by various corporations. By focusing on chemistry sets made available between 1920 and 1980, we can see a shift in how chemistry was viewed both as a science and what a scientist looked like. The collection photographed here, along with others, are on display at the Chemical Heritage Foundation in Philadelphia. 207

In Figure 5, a chemistry set from The Porter Chemical Company, we see a family around a table. The kids are performing various experiments while the parents are observing. On the box, it states that there are no dangerous chemicals in the kit. As we know from reviewing the types of materials that were in the kits of this era, this was not always the case. Some chemistry sets did contain materials that would not be considered safe by today’s standards. Sodium ferrocyanide was used to create Prussian blue dye, but is slightly toxic and no longer used. Potassium nitrate was included in some sets, which could be used to make smoke bombs.

Figure 5. Chemcraft No. 1 “Chemical Outfit” set, ca 1928:. Courtesy of Chemical Heritage Foundation Collections.

Chemistry was more than just a way for kids to pass the time. Chemistry kits were seen as a path to a career in chemistry. “Coming out of the Depression, that was a message that would resonate with a lot of parents who wanted their children to not only have a job that would make them money but to have a career that was stable. And if they could make the world a better place along the way, then even better,” says Rosie Cook, registrar and assistant curator at the Chemical Heritage Foundation in Philadelphia (1). 208

In Figure 6, the Porter Chemcraft set shows a number of bottles containing stock chemicals. There is also a scale and some test tubes along with some instruments for transferring materials. One of the booklets inside relates to outer space. During this time period, the space race was on. There was a large amount of interest in being the first country to make it into outer space, land on the moon, and explore the solar system. This was an excellent opportunity for toy manufacturers to market materials while also encouraging learning about different scientific fields.

Figure 6. Porter Chemcraft Senior Set No. 6105 (Interior) ca. 1957:. Courtesy of Chemical Heritage Foundation Collections.

In Figure 7, the Gilbert Lab Technician Set for Girls, ca. 1958, prominently shows a girl using a microscope, with a mother figure observing. Not shown here is the Gilbert Chemistry Set for Boys as the picture is not available. Here we see a differentiation between the kit marketed for boys and a separate set marketed for girls. This is a whole other topic which will not be discussed here. 209

Figure 7. Gilbert Lab Technician Set for Girls, 1958:. Courtesy of Chemical Heritage Foundation Collections.

In Figure 8, Mr. Wizard’s Experiments in Chemistry Set was marketed by a popular TV show at the time, watched by many children, including myself. Seen in the picture are several stock chemicals and basic scientific equipment found in earlier chemistry sets to allow the user to transfer materials and to observe those materials in a beaker or test tubes.

Figure 8. Mr. Wizard’s Experiments in Chemistry Set, ca. 1973:. Courtesy of Chemical Heritage Foundation Collections. 210

More recent chemistry sets have unfortunately been more reading-based and less hands-on chemistry based. This has been one of the reasons why the exhibit at the Chemical Heritage Foundation has been so popular. It is also one of the reasons why individuals and organizations in the general public have such a difficult task in relating the sciences, and chemistry in particular, to younger generations. Over the years, kids are being exposed to fewer opportunities to explore the world around them from a scientific point of view. That opportunity is being replaced with more readily available information, but the hands-on experience is lacking for many children and even adults.

Current Trends in the Perception of Chemistry Decades ago, the chemistry community was considered highly valuable along with other branches of STEM, even though that term was not yet in use. Dupont’s slogan starting in the 1930’s was “Better Things for Better Living, Through Chemistry.” Products were being developed, and the consumer had many more choices each year as companies strived to fill store shelves with new and innovative products. Many chemical companies such as Dow and BASF began to flourish. These companies continue to lead industry today. In addition to direct chemical companies, other industries were able to benefit from novel products based on chemistry. The automobile industry still continues to use the most efficient materials available. By using plastics instead of metals in some parts of automobile construction, the weight of the vehicle was able to be reduced, yielding a large cost savings to the company and a more economical vehicle for the consumer. Even things that are often overlooked, such as the paints and pigments used in auto construction have changed drastically over the years. The airline industry was able to benefit as well, as lighter weight planes are more fuel efficient which reduces cost. Even credit cards tried to cash in on the positive perception of chemistry. Taking the names of precious metals, different card options were made available. Gold cards were offered as a way for them to try to entice “better customers” to sign up with them. These “better customers” could receive higher credit limits and lower interest rates than their average customer. The Gold Card name was proudly displayed on these credit cards. Over time, Platinum cards and even Titanium cards have been issued. Credit card issuers tried to one-up each other until these names lost their value along the way. Over time, the pendulum has swung back around, and chemistry has received a negative perception in some circles. The public perception has changed to unfavorably view chemistry due in part to campaigns such as Chemical-Free (2). The concept of Chemical-Free is a marketing campaign where companies advertise that their products do not contain any chemicals. Many indicate that their products contain natural products, but no chemicals. Even someone with only a basic knowledge of chemistry knows that chemicals comprise almost everything that surrounds us. This certainly includes all the products on store shelves. The marketing plan was genius, as people embraced the idea of purchasing materials without chemicals in them and sought out items which made that claim on the 211

front of their labels. However, as required by law, on the back label of many products on store shelves, a list of ingredients indicates that those products are in fact made up of chemicals. The campaign unfortunately made many people become weary of chemicals. This negative public perception sometimes even applies to those who work in the scientific field. The Chemical-Free sentiment may have been better served if there was a focus on accuracy. For instance, if the intention was to say that the products in question were pesticide-free or carcinogen-free, then the label could be verified. Those items might even stand out on the shelf against their counterparts, assuming those counterparts did in fact contain amounts of pesticides or carcinogens. However, linking “chemicals” to things that are bad in marketing campaigns had a negative effect on the perception of chemistry in general. In addition, some materials are sold that are intentionally filled with pesticides, including items designed to help protect crops in the field. In these cases, the products need to be able to perform the functions that they are claiming on the label. The same can be said for carcinogens or other categories of chemicals. My own journey on the Anti-Chemical Free crusade began while reading C&EN articles (3) and Twitter posts on the topic. As a starting point, former C&EN reporter Carmen Drahl and I asked our Twitter followers to send us pictures of products that were sold using a Chemical-Free label on the front and a second picture showing the ingredient listing on the back of the label. We received pictures of sunscreens, organic fertilizers, even shopping bags, among others. One clever person even posted a picture of a vacuum cleaner. We co-authored a poster on the topic at the ACS meeting in Philadelphia August 2012. The poster received a lot of attention and comments. Unfortunately, our audience at the meeting were other scientists who know the benefit of chemistry to our lives. A better target audience would have been the general public, and I saw it as my mission to share this information. In order to promote the importance and value of chemistry to the general public, 2011 was declared the International Year of Chemistry (IYC) (4). Celebrations were scheduled around the United States and the world to showcase the positive role that chemistry has in the lives of ordinary citizens. As ACS Mid-Hudson Local Section Chair in 2011, I was able to organize events during IYC. Perhaps the program that stood out the most in my area was a collaboration with the local baseball team, the Hudson Valley Renegades. I have hosted events with them over the years, bringing science or engineering activities assisted by some local engineering students from West Point Military Academy. During games in the 2011 season, the Renegades handed out 2000 baseball cards to fans with a flashback player on the front and information about a non-profit on the back. Two of those games featured the ACS Mid-Hudson Local Section with a logo and website information, plus some basic facts about ACS. During one of those games, I was invited to give a pre-game interview which appeared on the jumbotron and a local sports channel, as well as the opportunity to throw out the first pitch. Figure 9 shows me before taking the field. My chemistry skills are far better than my athletic skills. However, the interview went well and was redistributed online over Facebook and Twitter. Positive feedback was received 212

and the interview was even noted in a committee report by the ACS Committee on Public Relations and Communications by staff member Nancy Blunt.

Figure 9. First Pitch before Renegades home game.

My Why: The Reason I Got Started and Why I Chose Palladium The largest driving force behind why I continue to conduct my science outreach activities is my belief that chemistry can be the answer to many of our challenges in society. Although it can be argued that the implementation of chemistry also has its share of blame in leading to accidents and pollution, if we look to move forward and look for better ways to use chemistry to our advantage, then chemistry needs to be a part of the solution. As I discovered Green Chemistry, complete with the 12 Principles (5), I learned that chemistry can be made inherently safer. As John Warner, one of the co-founders of Green Chemistry, has stated numerous times in lectures across the country, “A molecule is just a molecule. It is not inherently good or evil.” With this way of thinking, scientists can search for ways to create new products and new ways of understanding the mechanisms of reactions while at the same time take into account safety and minimizing risk. This way of thinking truly does make chemistry work for the people. 213

Over the years, as I searched for new and interesting ways to engage students and the public in general, I discovered many resources that Beyond Benign (6) has to offer. Beyond Benign was co-founded by John Warner and Amy Cannon. They have information and lab-based activities for K-12, college level, academics, and industry. Many science professionals have used their materials to educate students and the general public about greener lab practices. In addition, actions such as thinking critically about solving potential problems before they start, or improving products in the pipeline are major components of the concepts behind the Principles of Green Chemistry. There came a point in time when I wanted to go beyond just doing a few ACS-inspired outreach activities. I wanted to see how much I could accomplish by jumping into the field and promoting myself among others who were introducing science activities with kids. To do so, I was going to have to separate myself from everyone else, or I was just going to be lost in the crowd. A proper name would certainly go a long way toward accomplishing my goal. As I started to think about possible names, I wanted to embrace the essence of what I was trying to do. Something involving Green Chemistry should be in there somewhere, since as time went on I was incorporating more green activities into my roster. However, “Green Science” seemed a little too simplistic. As names bounced around in my head, element names started to pop up. Combining elements and Green Chemistry, there was a concept to help narrow down the choice: the use of catalysts. Many catalysts are used in chemical reactions to help speed up the reaction. Upon looking at possible names, one stood out. There was palladium, an element used quite often in literature. Palladium is useful as a catalyst in part because its properties are similar to platinum, located in the same column on the Periodic Table. It is also less expensive than platinum. A useful element, recognized catalyst, represented in the literature, it seemed like a great choice. Shortly after this revelation, Palladium Science Academy became an official entity, after paying a small fee at the County Clerk’s office. With a new name, new branding, and a renewed sense of purpose, I started marketing my new organization, Palladium Science Academy. Scientists thought it was great! Vendors thought I was interested in or used palladium, and wanted to sell me their stock containing the metal. The public, well, many did not quite understand the symbolism, and several could not pronounce the word palladium. So it was a bit of an uphill battle from the beginning. Recognizing the opportunity and challenge, I began my mission of spreading the word about the great possibilities of the sciences, while helping people pronounce palladium. Figure 10 shows my banner outside one of my first events, in New Paltz, NY. Building on my network of organizations with an interest in educating or entertaining children, I began working with schools, libraries, and summer camps. From young children in preschools up through high school students, Palladium Science Academy has something for everyone. It is always a challenge to come up with new and innovative ideas. Technology changes and people’s expectations change at a similar pace. While there are activities that might have seemed engaging twenty years ago, today if kids have already seen them on the internet several times, those same activities are no longer able to captivate their interest. Simple things like mixing Mentos and Diet Coke are still a hit because of the 214

mess factor. Many kids have already seen this activity but they are willing to watch it over and over again because of the messy reaction.

Figure 10. Banner for Palladium Science Academy outside an event.

Some of the most meaningful work that I have been doing with science outreach has involved middle school and high school students. To show them potential career options, I have participated for a number of years with the local Mid-Hudson branch of the American Society of Civil Engineers in the Hudson Valley of New York, where a number of volunteers teach math classes at Cornwall High School for a day. During the class period, we show the students what we do for a living. Most show engineering concepts and designs as the group is made up primarily of civil engineers with a few other types of engineers mixed in. For my part, I show the students a little bit of what it means to be a chemist. I demonstrate some chemical reactions, such as acid-base chemistry with indicator solutions that change color depending on the pH of the different materials. Since the main goal of this yearly project is to show students the relevance of what they are learning in math class is to a career, I share some simple calculations that an analytical chemist might use. 215

Conclusion In addition to bringing science to the public through various outreach and educational opportunities, I have also been fortunate enough to be able to publish articles or be featured in several publications. A full page article was written on Palladium Science Academy for the Times Herald Record (7) which also printed a few additional comments and photos from events at the Dutchess County Chamber of Commerce Expo in 2014. Events relevant to the outreach activities that I have participated in have been posted to the blog The Analyzer Source. As part of an ongoing project, I have written articles that focus on STEM events for Happy Hudson Valley, an online news source that focuses on positive newsworthy events in the Hudson Valley, NY. I have covered topics including an annual Pi Day article (8), events in the science and engineering fields, mostly focused on children learning about new activities, and some general interest stories. This is perhaps the best way that I have succeeded in bringing chemistry to the people. Over time, perhaps the public will even become more knowledgeable of the Periodic Table.

References 1.

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Zielinski, S. The rise and fall and rise of the chemistry set. The Smithsonian. http://www.smithsonianmag.com/science-nature/the-rise-and-fall-and-riseof-the-chemistry-set-70359831/ (accessed March 18, 2017). Drahl, C. Hey Burt’s Bees, Who’re You Callin’ Ugly? Chemical & Engineering News. http://cenblog.org/2008/10/hey-burts-bees-whore-youcallin-ugly/ (accessed September 28, 2017). Drahl, C. Wiping “chemical-free” off the marketing map. https://storify.com/ carmendrahl/shaming-chemicalfree (accessed March 18, 2017). United Nations. Celebrating the International Year of Chemistry. http:// www.un.org/en/events/chemistry2011/index.shtml (accessed September 25, 2017). American Chemical Society page adapted by Paul Anastas and John Warner. https://www.acs.org/content/acs/en/greenchemistry/what-is-greenchemistry/principles/12-principles-of-green-chemistry.html (accessed March 15, 2017). Beyond Benign. http://www.beyondbenign.org/ (accessed March 15, 2017). Walsh, J. Mad Scientist spreads the gospel of chemistry. The Times Herald Record. http://www.recordonline.com/article/20141025/NEWS/141029595 (accessed March 18, 2017). Ruger, G. Yes you still need to know this. Engineering Day brings real life application to the classroom. Happy Hudson Valley. http:// www.happyhudsonvalley.com/news/yes-you-still-need-know-engineeringday-brings-real-life-application-classroom (accessed March 18, 2017).

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Editors’ Biographies Mark A. Benvenuto Mark Benvenuto received his education at the Virginia Military Institute and the University of Virginia (BS and PhD, respectively), and did a post-doctoral fellowship at the Pennsylvania State University. He also served a four-year term of service between his undergraduate and graduate education as a lieutenant in the United States Army, spent mostly in Mannheim, West Germany. He joined the University of Detroit Mercy as a faculty member in inorganic chemistry in 1993, and has been department chair since 2001. Mark has taught freshman-level chemistry to science and engineering students virtually every semester since he has been at the University of Detroit Mercy, and has been voted the UDM Science Teacher of the Year, by the students, five times. He was also awarded the Michigan College Science Teacher of the Year in 2003 by the Michigan Science Teachers Association. He has been active in local and national ACS activities for two decades, and is a Class of 2015 ACS Fellow. Mark maintains research interests in two broad areas: coordination chemistry, specifically the development of multi-dentate ligands with unusual coordinating abilities; in the analysis of trace materials in archaeological objects as well as food supplements and personal care products via energy dispersive X-ray fluorescence spectrometry.

Tracy Williamson Tracy Williamson received her B.A. in chemistry in 1985 from Hamilton College and her Ph.D. in organic chemistry in 1992 from the University of Delaware. Tracy is currently the immediate past chair of the American Chemical Society (ACS) Division of Environmental Chemistry and is a member of the ACS Committee on Nomenclature, Terminology, and Symbols. She is a fellow of the American Chemical Society and was formerly a member of the ACS Committee on Environmental Improvement. Tracy is currently the Chief of the Industrial Chemistry Branch in the Office of Pollution Prevention and Toxics at the U.S. Environmental Protection Agency.

© 2017 American Chemical Society

Indexes

Author Index Archey, D., 195 Benvenuto, M., ix, 111, 195 Burdette, S., 119 Costa, M., 1 Eitrheim, E., 173 Fontani, M., 1 Forbes, T., 173 Karol, P., 41 Knight, A., 173 Kolopajlo, L., 141 McFarland, B., 67

Nelson, A., 173 Nguyen, G., 111, 195 Orna, M., 1 Pothoof, J., 111, 195 Restrepo, G., 95 Ruger, G., 203 Schultz, M., 173 Thornton, B., 119 Venugopal, E., 83, 195 Williamson, T., ix

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Subject Index A All-American errors, lost elements advent of the periodic table, 1869, 11 Bakerian Lecture of 1883, 14 Baskerville, Charles, 20f Baskerville, Charles, 1901, 16 carolinium compounds, samples, 19f Harvard University Museum of Natural History Earth and Planetary Sciences Gallery, meteorite display, 16f J. Lawrence Smith Medal, 17f Mendeleev, Dmitrii, monumental image, 13f Ottoman Empire, employment, 15 Smith, J. Lawrence, 1877, 13 thorium, chemistry, 18 atomic number and isotopes, 1913 and 1915, 21 Allison, Fred, 30f discovered in 1902, presumed naturally occurring elements, 28t false American elements, 32 genuine element 61, 26 Hopkins, B Smith, 1926, 22 Hubbard, H. D., periodic table, 33f magneto-optic diagram and connections, 31f professor Allison, Fred, 1930, 27 professor Hopkins, B Smith, 24f state-of-the-art instrumentation, University of Florence Physics Department in Arcetri, 25 chemical revolution to the periodic table, 1789-1869, 4 book jacket, 5f Chandler, Charles Frederick, element, 7 Chandler, Charles Frederick, portrait, 8f Chandler’s activities, some examples, 9 Genth, Frederick Augustus, 6f Genth, Friedrich August, element, 5 rare earth element compounds, display, 11f conclusion, 33 all-American elemental errors, summary, 34t early errors and early elements, before 1789, 3

erroneous element discoveries before 1789, 4t introduction, 1 Traité élémentaire de chimie, cover of the 1789 edition, 2f Always by-products, sludge elements, abundance, prevalence, and uses, 112 elements, products made, 113t matter, distribution, 112t other materials, elements co-produced, 114t Mendeleev’s “eka” elements and their uses, 115 mining and mining by-products, 115 Arsenic, selenium and mercury, analytical methodologies alchemy to chemistry Arabic alchemy, mercury, 142 conclusion and summary, 161 EPA and Safe Water Drinking Act, 153 Hg, As, and Se, MCL and MCLG data, 153t EPA mercury cold vapor and hydride methods, history arsenic and selenium, 159 EPA mercury cold vapor methods, timeline, 158t inorganic As and Se water, EPA methods, 161t mercury, 154 mercury cold vapor instrument schematic, 155f mercury in environmental samples, 155 oxidize bound mercury, bromine monochloride (BrCl), 157 European alchemy, mercury, 146 arsenic, 149 selenium, 150 introduction, 141 toxicity and safety, 150 modern cosmetic products, 152 Available elements, periodic table ancient environment, how oxygen affected, 76 available and functional, life requires elements, 71 elements in the modern ocean, concentrations, 74f large-scale chemical predictability, 73

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percentage of modern atmospheric levels, oxygen levels over time, 72f biochemical periodic table, environmental availability restricted, 75 old and new forms of the elements used by life, redox potential, 76f elemental abundance in different environments, 68 all elements in the universe, relative abundance, 70f preference for oxygen over sulfur, elements sorted, 71f natural history, unique perspective of chemistry, 77

B Binary compounds and their stoichiometries, 95 materials and methods, 97 different elements between Ni and Nj, number, 100f reactions ρ1 to ρ6, hypergraph, 98f some decomposition reactions of binary compounds, hypergraph, 99f two decomposition reactions ρ and ρ’, 97f results and discussion, 100 binary compounds, 94 elements explored, 100f 94 chemical elements, similarity landscape, 101f similar elements, largest cluster, 103 transition metal cluster, 102

E Early universe, elemental history early cosmology, role of radioactivity, 84 elemental abundances and modern cosmology, 89 primordial element abundances, comparison, 91f primordial nucleosynthesis and stellar thermodynamics, 86 Element 118 periodicity, traditional presentation, 196 student exercise, predicting element 118 properties, 198 column 18 of the periodic table, boiling points for elements, 200f

column 18 of the periodic table, melting points for elements, 199f student exercise, understanding of the nucleus, 197 constructing A 118 proton nucleus, 198f Elements, periodic table alpha decay, 51 stable or near-stable heavy and superheavy nuclei, 51f electronic structure, 58 finite-size nucleus, one-electron total energy, 60f one-electron systems, total orbital energy, 59f orbital radial density distributions with and without relativistic inclusion in copernicium, comparison, 61f systems containing eighteen electrons, total orbital energies, 61f future, 62 eka-radium, extended Periodic Table, 64f extended Periodic Table of Fricke et al, 63f extended Periodic Table through Z=172 following Pyykko, 63f Mendeleev-Seaborg extended Periodic Table prediction, 63f new elements, unsuccessful routes, 62t heavy and superheavy elements, production, 53 limits of stability, Z,N composition grid, 56f nuclide Z,N grid, 58f path of neutron captures, Z,N grid tracing, 54f transuranium elements, yields, 57f Z,N nuclide grid showing stable isotopes, 55f nuclear shell structure, 51 measured masses and liquid drop model predictions, difference between, 52f polonium isotopes, alpha-decay energies, 52f polonium isotopes and hassium isotopes, alpha-decay energies, 53f nuclear stability, liquid drop model, 47 alpha-particle emission and the rate (half-life) of such a nuclear transition, relationship between the energy available, 50f average binding energy per nuclear particle, 48f

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perception of chemistry, current trends, 211 Anti-Chemical Free crusade, 212 Chemical-Free marketing campaign, 211 Renegades home game, Pitch, First, 213f Why I Chose Palladium, 213 Palladium Science Academy, banner, 215f Principles of Green Chemistry, 214

millibarns as a function of energy of the compound system, fusion reaction yields, 49f nucleus, 47 periodic table, 41 elements as a function of time, cumulative number, 43f periodic table, upper limit, 43t predictions, history, 42t transuranium elements in nature, 44 nuclear composition, grid, 46f transuraniums in nature, summary, 46t

R I Isotopes, chemistry’s decision point, 119 chemistry, isotopes and the edge, 120 deuterium, reevaluate isotope chemistry, 129 epilogue, 133 geochemistry, 131 isotopes, physical reality, 126 philosophical interregnum, 128 radioactivity, another challenge to the periodic table, 121 Soddy, Aston, and the mass Spectrometers, 124

P Palladium Science Academy introduction and historical perspective, 204 alchemist at work, 17th century, 207f alchemist in his laboratory, late 17th century, 206f bald alchemist with his assistant, 17th century, 205f Chemical Outfit set, 208f Gilbert Lab Technician Set for Girls, 1958, 210f Mr. Wizard’s Experiments in Chemistry Set, ca. 1973, 210f Porter Chemcraft Senior set no. 6105, 209f Rijcke-Armoede 1632 Adriaen van de Venne, 204f

Radiochemistry, recent advancements, 173 drill cuttings, 180 future work with Pa, 187 hydraulic fracturing, radioactivity in liquid waste, 177 method development, 178 new discoveries, 185 protactinium in extraction systems, trace-level chemistry, 186 separation of Pa from Np, separation methods, 187f separations and purifications, advances, 186 nuclear forensics, plutonium nuclear materials and applications, 174 extraction chromatography, 176 TEVA/TRU, tandem column arrangement, 176f 210Po natural environments, enhanced understanding polonium health concerns, 182 polonium in groundwater, 182 polonium radiochemistry and sources, 180 polonium water chemistry, 181 protactinium, 183 protactinium and applications, chemistry, 183 231Pa, radioactive ingrowth, 184f 231Pa and 230Th, ingrowth, 185f radiochemical disequilibrium, 179 radium decay products, 179 separations and nuclear forensics applications, advancements, 177

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