Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction
GREGORY H. ROBINSON, THE UNIVERSITY OF GEORGIA
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- Bu səhifədəki naviqasiya:
- OLIVER SACKS, NEW YORK CITY
- CELIA HENRY, CEN WASHINGTON
- JEROME O. NRIAGU, UNIVERSITY OF MICHIGAN, ANN ARBOR
- THALLIUM AT A GLANCE
- GEORGE A. OLAH, UNIVERSITY OF SOUTHERN CALIFORNIA
- ROALD HOFFMANN, CORNELL UNIVERSITY
- SILICON AT A GLANCE
- SANDY GERRARD, ENGLISH HERITAGE
- BASSAM Z. SHAKHASHIRI, UNIVERSITY OF WISCONSIN, MADISON
GREGORY H. ROBINSON, THE UNIVERSITY OF GEORGIA
Even its position on the periodic table--gracefully residing between the only nonmetallic element of group 13, boron, and the oddly mercurial metal of gallium--is perhaps indicative of how extraordinary element 13 is. There is not even complete agreement on the spelling and pronunciation: Americans typically employ aluminum, while significant portions of the remaining English-speaking world prefer aluminium. Aluminum, like most elements, has its share of interesting trivia: For example, a 2.73-kg pyramid of "precious" 1884 aluminum metal sits atop the Washington Monument. While Hans Christian Oersted is acknowledged as the first to isolate aluminum in 1825 in Copenhagen, Denmark, the eminent German chemist Friedrich Wöhler is generally regarded as the first to secure a pure sample of the element by chemical reduction in 1827. The intriguing international tale of the discovery of the economical production of aluminum independently by two young men, the American Charles M. Hall and the Frenchman Paul L. T. Héroult, via electrolysis of alumina dissolved in cryolite, is well documented. Certainly, the aluminum industry as we know it today is due to the creative genius of Hall and Héroult. My fascination with aluminum, however, has less to do with the actual element and more with the relationship between Frank Fanning Jewett and Hall. Jewett, educated at Yale University in chemistry and mineralogy, had a passion for travel. Indeed, he studied briefly at Universität Göttingen, spending time in the laboratory of Wöhler. In 1880, the 36-year-old Jewett was appointed professor of chemistry and mineralogy at Oberlin College. Thus, the stage was set for the well-traveled professor and the prodigious student. Jewett is the (most often) anonymous "professor" at Oberlin who opined to his chemistry class, where Hall was in attendance, that great financial rewards awaited the person who could devise an economical means to produce aluminum metal from its ubiquitous ore. The role of Jewett in Hall's life proved crucial in the seminal discovery that would ultimately spawn the Aluminum Co. of America, Alcoa (2002 revenues of $20.3 billion), and the worldwide aluminum industry. The initiative and drive of Hall remain impressive. For an amateur scientist to doggedly pursue a scientific problem of such magnitude, and ultimately succeed in such an endeavor, is almost unimaginable. Records indicate that Jewett was a counselor, mentor, adviser, and friend to Hall. Furthermore, Jewett often provided materials and laboratory space to the budding entrepreneur. Jewett, a modest man by all accounts, apparently was not interested in sharing the praise, fame, or financial rewards that would soon befall his student. As Oberlin College's Norman C. Craig has so elegantly stated, Jewett was "content to report to his Yale classmates that his greatest discovery was the discovery of a man--Charles Hall" [Chem. Heritage, 15, 36 (1997)]. In my mind, the Jewett-Hall relationship epitomizes the idyllic professor-student dynamic. It is this relationship that I envision when I am working with students: a synergistic pursuit of the scientific unknown. To be sure, the stakes are much lower in my day-to-day struggles in the laboratory. The problems that my students and I face are much smaller in magnitude, and any potential immediate impact is often ambiguous. Nonetheless, the Jewett-Hall relationship drives me in an oddly personal manner as I strive to improve my teaching skills and hone my research capabilities. Might my perspective on this relationship be a rather naive interpretation? Almost certainly. Is this simply an outdated commentary on the contemporary professor-student dynamic? Most clearly. Could this all be little more than a "nonprofessorial" waste of time? Absolutely not! The professor-student dynamic represents much of what I find uniquely attractive in academia. I have observed parallels to the Jewett-Hall relationship in athletics: In tennis, it is that perfectly executed service ace down the middle of the court; in basketball, it is that gracefully arching jumpshot from the corner, hitting "nothing but net"; in golf, it is that splendid tee shot on the par 5, 18th hole--you know, that one shot that keeps bringing you back time and time again. 
Gregory H. Robinson is a distinguished research professor of chemistry at the University of Georgia. His research interests, the organometallic chemistry of the main-group metals, are tempered by his recent obsession with golf. The author acknowledges the gracious assistance of Norman C. Craig (Oberlin College) and Richard K. Hill (University of Georgia) with this essay.
GALLIUM OLIVER SACKS, NEW YORK CITY
Why gallium? Not a common element, not one likely to be lying around the house, but one I was introduced to quite early on, by my Uncle Tungsten (as we used to call my Uncle Dave, who owned a tungsten lightbulb factory). Uncle possessed what was, to my eyes, a quite wonderful thermometer that he used for testing the temperature in the furnaces in his factory. This, he showed me, contained gallium; it was the only stuff such a thermometer could contain, for gallium had the widest temperature range of any metal in the liquid state: It would melt in the heat of the hand but not boil until well over 2,000 °C (higher than the melting point of platinum or of the quartz of which the thermometer itself was made).
A strange optical illusion appears if one melts gallium in a cup: There seems to be a transparent liquid skimming and floating above a silver background. Is this due to the strongly concave meniscus of liquid gallium? I do not think I ever saw this, by contrast, with mercury, which has a strongly convex meniscus. And once melted, gallium may remain liquid, superfluid, for many hours, even if the room temperature is well below its melting point. The liquid may form a skin, wrinkled with fine lines, and when it finally solidifies, it may do so in shallow quadrangular prisms and zigzags like medieval fortifications. I also had a little stick of indium, another element that intrigues me, partly because, like tin and zinc, it emits a "cry" or "squeal" when bent. (With my little bar, it was more like a crackling.) One day, just recently, I carelessly left the indium on top of some gallium I had in a bowl, and I was startled to find, within hours, that the bar seemed to have partly dissolved and that there was now a pool of liquid metal at the bottom of the bowl, despite its being a rather cold day. Clearly, the two elements, merely by being in contact, had fused together to form a eutectic alloy with a substantially lower melting point than that of pure gallium. I was reminded of how Berzelius had been sent samples of metallic sodium and potassium, which were put together in the same container for convenience, and when he opened the package, he found only a pool of liquid metal, the two elements having spontaneously alloyed at room temperature, just as my indium and gallium had. When I came to learn about the periodic table and its history, I was intrigued to learn that gallium was the first element to be predicted by Mendeleev based on its place in Group III (he called it "eka-aluminum"), and how this prediction was vindicated, just six years later, helping to convince Mendeleev's critics of the fundamental truth of his periodic law. Many decades later, I was fascinated to learn that there was a huge pool containing 200 tons of ultrapure liquid gallium deep beneath the Caucasus, an essential part of the Soviet solar neutrino detector. A passionate swimmer, I had fantasies of swimming, or rather floating, on this unique lake of metal. I was shocked when I read, a few years ago, that thieves had come by with siphoning equipment one night and almost managed to steal the whole lot. The great gallium heist was foiled only at the last moment.
Oliver Sacks is a neurologist practicing in New York City. He is the author of "Uncle Tungsten: Memories of a Chemical Boyhood," as well as "Awakenings" and "The Man Who Mistook His Wife for a Hat."
INDIUM CELIA HENRY, C&EN WASHINGTON
"I hadn't performed in 20 years, but for some unknown reason I said yes," he says. He was afraid that getting his equipment to St. Petersburg would be a "nightmare," so he decided to compose a "minimal piece of slow-moving music and use a video to accompany it." Bone found inspiration for that piece in the most unlikely of places--the Periodic Table of the Elements. The combination of music and video is known as "Indium," a 30-minute composition consisting of three distinct sections. In the accompanying video, the scene changes every 10 minutes to match the changes in the music. It falls into the genre of experimental music known as ambient music. Bone got his start in pop music and says that he is still something of a pop songwriter, because most of his pieces are only three or four minutes long. "This was the first time I had ever attempted to compose something that was 30 minutes long and try to keep it interesting and moving," he says. It was written as three separate pieces that Bone then had to flow into each other to make one continuous piece of music. Interestingly, although indium ultimately provided the inspiration for the finished piece, the idea was originally planted by another element. A fan of the Spanish surrealist painter Joan Miró, Bone visited the Fundació Joan Miró in Barcelona, where he saw the "Mercury fountain," which was created by Alexander Calder for the 1937 World's Fair in Paris. In this sculpture, liquid mercury is piped from a pool and flows down a series of curved metal pieces back into the pool. The sculpture was created as a political statement against the seizure of the Almaden mercury mines by Francisco Franco during the Spanish Civil War.
But he didn't want to call a piece "Mercury" because that had already been done, so he looked elsewhere for an appropriate name. "I somehow stumbled across a periodic table and saw the word indium," Bone says. "I'd never heard of it, so I looked up the description of it. The description almost perfectly described the music and the video that I wanted to do." Indium is a soft, lustrous, silvery white metal. It is useful for making alloys with low melting points. For example, an alloy of 24% indium and 76% gallium is a liquid at room temperature. The element was discovered spectroscopically in 1863 by Ferdinand Reich and his assistant Hieronymous T. Richter while they were searching for thallium in zinc ore. The element was named for the brilliant indigo line in its spectrum. The pure metal was first isolated in 1867 by Richter. Until 1924, about a gram represented the entire world supply of the isolated metal. Indium was originally thought to be rare, but it is actually about as abundant as silver. The definition of indium that Bone found "evoked in me the music I was creating." He printed the definition and taped it over the keyboard where he could see it as he worked on the piece. "There was something about the description of that element that seemed to capture what I wanted to create in the music," he says. He took the definition to Jim Karpeichik, a local videographer, and told him that he wanted to create a visual that evolved and moved very slowly. Karpeichik filmed ocean scenes at the beach in Rhode Island, slowed them down, and colorized them. Bone's approach to composing "Indium" was unusual for him. Typically, he improvises on the keyboard, looking for an unusual sound that he hasn't worked with before that can inspire him melodically. The work slowly starts to evolve from there. Often, he doesn't have a title for a work until he's done. "This was going to be a live performance piece, originally," Bone says. "I needed to have some very clear concept of what I was going to do." The music festival fell through when the Russian government declined to fund it, but the festival producers also owned a record company. The Russian record company Electroshock Records released the album "Indium" in December 2002. In the U.S., it is exclusively available from the online ordering sites http://www.eurock.com and http://gemm.com. Another interesting tidbit, given that this is C&EN's 80th or "mercurial" anniversary: The second track on the album is called "Mercurial Wave."
THALLIUM JEROME O. NRIAGU, UNIVERSITY OF MICHIGAN, ANN ARBOR
Public angst and concern were first drawn in the late 1960s to reports of widespread contamination of the Great Lakes' ecosystems with toxic metals. At the time, I was a graduate student at the University of Toronto. Of the various groundbreaking studies on the hazards of heavy metals in the Great Lakes basin, I was captivated by reports that alluded to the fact that symptoms typical of thallium poisoning were being observed in many wildlife populations in the Great Lakes basin, and especially one report that claimed that nine out of the 34 bald eagles found sick or dying in 1971-72 in parts of the basin in the U.S. were poisoned by thallium. For some inexplicable reasons, these early reports on poisoning of wildlife by thallium were ignored by the scientific community and no efforts were made to conduct necessary follow-up studies. Nevertheless, they aroused my curiosity and served as my initiation into the literature on the environmental chemistry and toxicity of thallium. Thallium is an enigmatic element to study because of its highly divergent properties predicated upon its oxidation status. Its paradoxical nature became obvious after its discovery as attempts were made to place the new element in the periodic classification of the elements. The physical properties of elemental thallium, including specific gravity, hardness, appearance, melting point, and electrical conductivity, were found to be similar to those of lead. The chemical properties of many Tl(I) (thallous) salts resemble those of lead, but the valency and other features of thallium were found to be too divergent from the lead family of elements.
Thallium elicits some of the most complex and serious toxicities in living things, involving a wide range of organs and tissues. A recent study with planktonic communities showed that Tl3+ ions, a common form in aquatic environments, are about 34,000 times more toxic than cadmium ions [Environ. Sci. Technol., 37, 2720 (2003)]. This is serious, considering that organisms (especially mammals) at the top of the food chain are the most susceptible to thallium toxicosis. Recently, we have found elevated levels of thallium in Great Lakes fish [Bull. Environ. Contam. Toxicol., 67, 921 (2001)], which has raised the issue of bioaccumulation and biomagnification of this element in aquatic food webs and the potential risks to fish-eating members of the food chain. I keep wondering whether the dismissal of the early reports of thallium poisoning of bald eagles and other wildlife should, in fact, be regarded as a major scientific oversight of our time. Compared with many heavy metals, thallium has a short history and what appears to be a rosy future. Its traditional uses (in rodenticides and insecticides, pigments, wood preservatives, and ore separation; in mercury lamps to increase the intensity and spectrum of the light; as catalysts in chemical syntheses; and so forth) are being phased out in deference to its toxicity. At the same time, there is increasing demand for thallium in the high-technology and future-technology fields. Among the growing uses for thallium are in the semiconductor and laser industry, in fiber (optical) glass, in scintillographic imaging, in superconductivity, and as a molecular probe to emulate the biological function of alkali-metal ions. Recycling of thallium in commercial products is not yet a serious business, and one has to be concerned about the long-term environmental effects of the growing technological applications of one of the most toxic and eccentric metals known to humankind. Jerome O. Nriagu is a professor in the School of Public Health at the University of Michigan, Ann Arbor. He is the editor of "Thallium in the Environment" as well as many other volumes on heavy metals in the environment.
CARBON GEORGE A. OLAH, UNIVERSITY OF SOUTHERN CALIFORNIA
Elemental carbon is found in nature as its allotropes--diamond and graphite--which are of vastly differing abundance and thus also of differing value. In the 1980s, a new group of allotropes called fullerenes or buckyballs--named after R. Buckminster Fuller, who designed famous geodesic domes resembling soccer balls--was recognized, first spectroscopically (for which Robert F. Curl Jr., Sir Harold W. Kroto, and Richard E. Smalley received the 1996 Nobel Prize in Chemistry) and later produced in electric discharge devices using carbon electrodes (Donald R. Huffman and Wolfgang Krätschmer). These new carbon allotropes promise significant applications. Carbon has a remarkable ability to bind with itself to form chains, rings, and complex structures. The variety of carbon compounds with bound hydrogen (hydrocarbons) and other elements (oxygen, nitrogen, and phosphorus, for example), which are generally called organic compounds, is practically unlimited.
The related development and use of superacids, billions of times stronger than sulfuric acid, and superelectrophiles have changed the field of chemistry. The realization that carbon, in many of its electron-deficient species, including carbocations, can simultaneously bind (coordinate) to five, six, and even seven neighboring atoms or groups is significant. This has extended Kekulé's concept of the limiting four valency of carbon to higher coordinate carbon compounds and has opened up new fields of hypercarbon chemistry. My discoveries also provided insights into the electrophilic activation of C–H and C–C single bonds and formed the basis of the development of new and improved hydrocarbon transformations. These have significant applications in the petroleum and chemical industries for improved production of high-octane gasoline (via alkylation and isomerization) and the direct conversion of methane (natural gas) to methanol and higher hydrocarbons without producing syngas. From plant life over the ages, new fossil fuels can be formed. The process is so slow, however, that within our human life span we do not have time for nature to replenish what we are rapidly using up. A challenging new approach that we are pursuing is to reverse the process and produce hydrocarbons from carbon dioxide and water via methanol, thus chemically recycling carbon dioxide. In the laboratory, we already know how to do this, and progress is being made toward bringing about the feasibility of such an approach. The limiting factor is the energy needed for generating hydrogen from water. Using alternative energy sources--but first of all atomic energy, albeit improved and made safer--will eventually give us needed energy. Much is said these days about a "hydrogen economy," emphasizing hydrogen as the clean, inexhaustible fuel of the future. However, the safe handling and dispensing of volatile hydrogen--for which no infrastructure exists--is difficult and costly. I believe a much preferable way of storing hydrogen is in the form of methyl alcohol ("methanol economy"). Methanol is a convenient liquid that can be produced by reduction of carbon dioxide in the atmosphere. It can be catalytically converted into ethylene and propylene and through them to higher hydrocarbons. This can provide an inexhaustible source of hydrocarbon products and fuels, which are now obtained from oil and gas. Furthermore, in recent years, with colleagues at California Institute of Technology and the Jet Propulsion Laboratory, we have also developed a new, direct methanol fuel cell that produces electric power without the need of hydrogen. Thus, methanol is both a fuel and a source of hydrocarbons. By recycling excess CO2 into methanol instead of just storing or sequestering it, we can also mitigate global warming. It is to this effect that a major research effort, with my colleagues associated with the Loker Hydrocarbon Research Institute at the University of Southern California, is directed.
George A. Olah is a professor of chemistry and director of the Loker Hydrocarbon Research Institute at the University of Southern California. He won the 1994 Nobel Prize in Chemistry for his work on carbocation and hydrocarbon chemistry.
SILICON ROALD HOFFMANN, CORNELL UNIVERSITY
The story was into similarity; it invoked no less an icon of analogical thinking than Dmitry Ivanovich Mendeleyev. Silicon was under carbon; surely it could not be denied its Mendeleyevian birthright, to give us all that carbon did? In the heyday of silicon chemistry, it was just a matter of time and chemistry. Or maybe of different ambient conditions; silicon-based life on another planet made such good science fiction sense. I was anxious to learn English, to speak without an accent, to fit in; of the life-enhancing differences, such as that between boys and girls, there were only stirrings. I lapped up that vision of an alternative, yet like, universe. But silicon is different. Yes, one can make Si analogs of hexane and cyclohexane. The differences emerge earlier in reactivity than in structure. Most importantly, unsaturation is strongly disfavored for silicon, as it is for other main-group elements below the first row. At a normal Si–Si distance, p bonding is just not worth very much. So double- and triple-bonded molecules have high-lying filled and low-lying unfilled orbitals. This makes for reactivity to acids, bases, and radicals. Silylenes, SiR2, are also very stable; put these trends together and an equilibrium such as H2Si = SiH2 Just how different silicon is shows up in the equilibrium structure of Si2H2. In a triumph of computational chemistry (not by me, alas), Si2H2 was predicted to have the striking nonclassical dibridged structure, which was then confirmed experimentally. The bond energies of C–E and Si–E (E = an element) are similar (within 50 kJ per mol) for a large variety of Es. The exception is Si–O (and Si–F), which is nearly 100 kJ per mol stronger than C–O. Coupled to the unhappiness of Si with unsaturation, silicon's love of oxygen leads to an overwhelming difference between a world of CO multiple bonds and the universe of silicates. Contrast gaseous CO2 and solid polymeric SiO2. Poor carbon, to have no 2- and few 3-D polymers. Compare clays, micas, zeolites, aerogels, amethyst and carnelian, lapis lazuli, asbestos, porcelain, and glass--just to pick a sampler of silicates. Silicon can certainly support chirality. Its chemistry sports helices. So why has this element essentially been dealt out of life, at least out of animal biochemistry? The abundance of silicon in Earth's crust is second only to oxygen. Much less common elements--take copper or molybdenum--are co-opted by life; one would have thought that nature's evolutionary tinkering surely would have found ample use for silicon. I'm not quite fair. The element, through silica, is of immense utility as a structural material in diatoms and radiolarians. A typical one, Cyclotella cryptica, is 22% SiO2 by dry weight. There are a lot of those little guys, too. Silicates are also essential for higher plants. Horsetails, which once grew into tall trees, accumulate enough silica to make them used as scouring brushes. There's a lot of silica in rice and in grasses in general. But almost none in bacteria and animals. It's a great puzzle. Though there are advocates of the way of clay, it appears that the road to life was through isomeric variability and a balancing act of metastable yet kinetically persistent small organic molecules and one-dimensional polymers. Poor silicon, capable of supporting isomerism but pulled away from Si–Si catenation by that seductive bond to abundant oxygen! How unexpected then (little help from science fiction here) is our here-and-now-and-still-growing in silico world. This is silicon's revenge! Or the evolution of electronics and computers might be seen as only the latest extension of human cultural appropriation of those elements underused in biology. Following (for silicon) that constant of civilization--ceramics. And glass. In the excitement of building a theory, in the desperate and thrilling labor of trying to understand, the primal urge is to simplify. Not just in science, though our craft is prone to special reductive tendencies. Simplifying to excess, we make things the same. But the beauty of this world is as much in the workings of chance and in the infinite variety that being not quite the same provides. So it is for silicon and carbon, as it is for people. Roald Hoffmann is the Frank H. T. Rhodes Professor of Humane Letters at Cornell University. He received the Nobel Prize in Chemistry in 1981.
GERMANIUM BETHANY HALFORD, C&EN WASHINGTON
When Mendeleyev conceived the periodic table, germanium hadn't even been discovered. I imagine him as a father with an overwhelming brood, leaving a place for an errant child, knowing that it has to be around somewhere. It doesn't even get respect from spellcheck, which insists on renaming it as a flower. But as elements go, germanium is a late bloomer. It took the inventions of the transistor and crystal diode in the 1940s to transform it from chemical curiosity to important industrial material. Before 1945, only a few hundred pounds of the element were produced each year. But by the end of the 1950s, annual worldwide production had reached 40 metric tons. Finding steady work in solid-state electronics, germanium held onto its high-tech reputation through the 1970s. And when sister silicon edged in on its territory, germanium moved on to more intriguing applications: polymerization catalysts, fiber-optic communication networks, and infrared night-vision systems. The U.S. government even designated germanium as a strategic and critical material, calling for a 146,000-kg supply in the national defense stockpile for 1987. Like the brainy, adolescent wallflower who grows up to make a fortune as president of a high-tech company, germanium had made a metamorphosis. Its cachet has come at a cost, though. In 1998, the price of 1 lb of germanium was almost $800; 1 lb each of silicon, tin, and lead could be bought for less than $5. My own personal experience with germanium, and its high cost, came the summer before I started graduate school when I worked with noted main-group organic chemist Peter P. Gaspar at Washington University in St. Louis. I was to make several hundred grams of (CH3)2GeCl2 using Gaspar's clever synthetic protocol [Synth. React. Inorg. Met.-Org. Chem., 20, 77 (1990)]. Perhaps because he was reluctant to let an inexperienced student like me work with pricey germanium powder--I had given him a purchase order for several thousand dollars' worth of the stuff--Gaspar suggested I call some chemical suppliers and see if any would be willing to discount their older stock. A little bit cocky, I was hurt at his lack of faith in my synthetic skills. One supplier gave me a bargain price on some decade-old "five nines" pure material that had been taking up space on his shelves. Gaspar had said that less pure material tended to work better in the reaction, so I agreed. I figured the stoichiometry in my notebook for the germanium I needed if it was only 5/9 pure. I told Gaspar of the deal. "I thought I told you that less pure material worked better," he said. "Yeah, it's only five nines pure," I countered. He stared at me, unmistakably annoyed. Irritated at his lack of enthusiasm, I skulked away. Only when the dusty old bottle arrived did I learn that "five nines" pure means 99.999%. Gaspar once showed me a half-kilo ingot of germanium metal that a company had given him in the halcyon days before the government began stockpiling the stuff. I can still remember the sleek, cool feel of the dark metal. By the way Gaspar cradled the bread-loaf-shaped object in his hand, I could tell it was one of his treasures. I was paralyzed by the thought that he might actually ask me to grind the material up and use it. For this essay, I called Gaspar to ask if he still had the ingot or if some unlucky student had to grind it up in the six years that had passed. "I keep it as a curiosity," he said, but conscious of its value, would not elaborate. "It's in a safe place." When I asked Gaspar what he thinks about my theory of germanium as the main group's middle child, he said he could see the parallel, but added--in his wise professorial tone of voice--that germanium has never had plain-Jane status in chemistry. On the contrary, "the most important thing about germanium was the fact that it started as a twinkle in Mendeleyev's eye," he said. "It might have been found much later than it was if it weren't for the fact that it had to be there if the periodic table was correct." Bethany Halford is an assistant editor at C&EN. Five nines of the time spent on this essay went into writing this sentence.
TIN SANDY GERRARD, ENGLISH HERITAGE
I will remember the first three days of that dig for the rest of my life. Dense fog, combined with torrential rain and powerful winds tragically known as the Fastnet Storm, coincided with our arrival. Once the winds had abated, the sun shone, and very rapidly it became apparent as we dissected the mill and the surrounding dressing floors that we were revealing a very complicated and informative story. The mill, throughout its life, from at least 1507 until around 1600, had crushed tin from the nearby opencast quarry, but it had been abandoned and refurbished on several occasions, each time becoming more efficient. A wide range of artifacts provided an insight into the character of the tin processing operations, and domestic rubbish gave us a glimpse into the lives of the tinners who had worked here. The true complexity of the site only became apparent during the post-excavation process and preparation of the final report.
So it was on the last day of January 1983 that I set off with a plane table to tackle the earthworks in the valley bottom next to the mill. I really had no idea whether it would be possible or even worthwhile, as nobody else had ever tried to tackle this type of survey. I certainly did not know when I plotted the first point on the table that the results would be so informative that for the next 20 years I would be spending large blocks of time surveying and interpreting tinwork earthworks all over Cornwall and Devon. The survey work revealed that it was possible to demonstrate exactly how the tinners had used different methods to extract tin. By doing a detailed analysis of the streamwork plans in particular, I could identify the precise methods used to extract the cassiterite and to recognize earthworks of different dates. Much has yet to be achieved regarding the absolute dating of the tinworks, but pollen analysis near the Colliford tin mill indicated that at least part of the tin streamwork was abandoned before 600 to 700 A.D. Many of the surviving streamworks in the southwest of England will be much more recent than this, as most probably belong to the late medieval period (1300 to 1500). The scrutiny of streamwork earthworks has been the most productive aspect of the detailed survey of tinworking remains, but other types of tinwork lend themselves to this form of examination to a greater or lesser extent. Analyses of the surface workings associated with early forms of mining--including the shallow shafts known as lode-back pits and the opencast quarries known as openworks or beams--have provided a valuable insight into the earliest forms of mining. Together with investigations of the contemporary documentation, the surveys have allowed us to build a remarkable picture of the industrial, technological, and social character of early tin exploitation in Britain. Sandy Gerrard is a designation archaeologist for English Heritage. He has directed several archaeological excavations and surveys and has published extensively on the early tin industry.
LEAD BASSAM Z. SHAKHASHIRI, UNIVERSITY OF WISCONSIN, MADISON
One of my classmates was overweight, and we nicknamed him "lead." Later on, when we learned about the symbols of the elements, we referred to him simply as Pb, an abbreviation of the Latin plumbum, from which we also get the word plumbing. A Jesuit archaeologist told me that lead is mentioned in the Bible many times, that alchemists believed it was the oldest metal, and that they tried to change it into gold. He also told me that lead is poisonous. This fascinated me because I wondered how a solid piece of metal could be eaten. The Romans used lead to make pipes, some of which may still be in use. Later, I read that the Romans also used lead in cookware and drinking cups and that lead poisoning may have contributed to the fall of the Empire. In high school and later as a freshman at the American University of Beirut, I learned more about lead. I was astonished to discover that lead is the end product of three series of naturally occurring radioactive elements and that lead has more than two dozen radioactive isotopes. In my sophomore year at Boston University, I learned that lead was in a gasoline additive--an organometallic compound called tetraethyl lead. I learned that lead is a major component of automobile batteries, that it was widely used in paint, and that it is used in soldering. Later, I learned that even in ancient times, some physicians believed that lead was poisonous, but it continued to be used in medicines and cosmetics until the 20th century. Concerns over lead exposure in recent decades reflect heightened societal concerns for health and safety. All developed countries have banned two uses of lead that were once almost universal--tetraethyl lead and lead-based paint for residential use. The use of lead shot for hunting waterfowl has also been banned in the U.S. because bottom-feeding birds can ingest spent shot. Hunters must use steel shot instead. Older U.S. cities that have lead pipes or heavily soldered pipes are replacing them at considerable expense. There is continuing concern about lead glazes still used in some ceramics made in developing countries, but the lead found in fine china and crystal glass does not readily leach out and is not a concern. Until recently, lead poisoning was diagnosed by its symptoms, but it is now diagnosed by analyzing its presence in blood by atomic absorption methods. The mechanism is not well understood, but lead is believed to bind to proteins involved in neurological signaling and development that otherwise would bind to calcium or zinc. Removing lead from the body safely remains a major challenge.
I have been a Wisconsin Badger for more than 30 years. The nickname of the state, and the mascot of the University of Wisconsin, Madison, are derived from lead. Lead mining in southwestern Wisconsin was one of the state's first industries. During the 1830s, lead miners, mostly young, single men, came up the Mississippi and Wisconsin Rivers from St. Louis and points south to spend their summers digging for lead. Not wanting to waste time on frills like housing, they lived in dugouts in the sides of hills. Some residents said they lived like badgers, a derogatory label they adopted with pride. Bassam Z. Shakhashiri is professor of chemistry at the University of Wisconsin, Madison. His annual "Once Upon a Christmas Cheery in the Lab of Shakhashiri" is in its 34th year and is presented on PBS and cable stations.
NITROGEN PETER NAGLER, DEGUSSA AG
Although nitrogen was initially thought to be unreactive, this simple diatomic molecule does have interesting chemical properties. Some elements can "burn" in nitrogen, among them magnesium at 300 °C and lithium even at room temperature, producing crystalline metal nitrides. Complexes of molecular nitrogen with transition metals had been predicted for many years, but the first organometallic compound of dinitrogen, [Ru(NH3)5N2]Cl2, was only reported in 1965. Most of the fascinating chemistry of nitrogen is, however, reserved for its inorganic and especially organic compounds, notably amines, nitro compounds, and their more complex relatives.
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H3Si – SiH, highly endothermic for carbon, is nearly thermoneutral and faces a small activation energy for silicon. Or, should you want something still more striking, it takes less energy to dissociate H2Si = SiH2 into two silylenes than it takes to break the single bond in H3Si–SiH3. Five and six-coordinate structures also become accessible for neutral compounds of Si.
