The noble gases, also known as rare or inert gases form Group 18 of the Periodic Table, embedded between the alkali metals and the halogens. The elements helium, neon, argon, krypton, xenon, and radon are the members of this group. In 1785 English physicist and chemist Henry Cavendish performed an experiment in which he passed electric sparks through an air bubble enclosed by a soap solution (NaOH). While nitrogen and oxygen were absorbed by the solution, about 1/120th of the volume of the original bubble remained; it is now known that the residual gas was mainly argon. However, it was a century later before argon was finally recognized as a new element. In 1894 English physicist John William Strutt noticed that nitrogen produced from air had a slightly higher density than that from nitrogen compounds.
Sir William Ramsay, together with Strutt, repeated the Cavendish experiment and identified argon as the unreactive species. The liquefaction of air in 1895 by Carl von Linde allowed Ramsay the further discovery of neon, krypton, and xenon. Extraterrestrial helium had been discovered earlier (in 1868), based on its spectral lines in the Sun. Ramsay realized that the new elements did not fit into the contemporary periodic system of the elements and suggested that they form a new group, bridging the alkali metals and the halogens. The last member of the family, radon was discovered in 1900 by Ernest Rutherford and Frederick Soddy as a decay product of radium.
The chemical inertness of the noble gases is based on their electronic structure. Each element has a completely filled valence shell. In fact their inertness helped to develop the key idea of a stable octet. The atomic sizes of the noble gas elements increase from top to bottom in the Periodic Table, and the amount of energy needed to remove an electron from their outermost shell, the ionization energy decreases in the same order. Within each period, however, the noble gases have the largest ionization energies, reflecting their chemical inertness. Based on increasing atomic size, the electron clouds of the spherical, nonpolar, atoms become increasingly polarizable, leading to stronger interactions among the atoms (van der Waals forces). Thus, the formation of solids and liquids is more easily attained for the heavier elements, as reflected in their higher melting points and boiling points. As their name implies, all members of the family are gases at room temperature and can, with the exception of helium, be liquefied at atmospheric pressure.
Until 1962 only physical inclusion compounds were known. Argon, krypton, and xenon form cage or clathrate compounds with water (clathrate hydrates) and with some organics such as quinol. The host molecules are arranged in such a way that they form cavities that can physically trap the noble gas atoms, referred to as "guests." The noble gas will be released upon dissolution or melting of the host lattice. In 1962 the first chemical noble gas compound, formulated as XePtF6, was synthesized by Neil Bartlett. This result spurred intense research activity and led to the discovery of numerous xenon and krypton compounds. In 2000 the formation of the first argon compound, argon fluorohydride (HArF), was reported by Leonid Khriachtchev and colleagues.
Science is frequently a collaborative discipline. But sometimes, one person, working alone in a laboratory, makes a stunning discovery, one that changes the way scientists look at their field. One such man was Neil Bartlett, who changed the face of chemistry while working alone in the chemistry department of the University of British Columbia. Bartlett demonstrated in 1962 that the "inertness" of the group VIII elements was a result of the reagents used in experiments, not a fundamental law of nature. Bartlett's proof that the rare gases were not chemically inert meant that all existing textbooks had to be rewritten. The American chemical society and the Canadian society for chemistry designated the work of Neil Bartlett and the reactive noble gasses an international historic chemical landmark at the University of British Columbia on May 23, 2006.
Neil Bartlett was born September 15, 1932 in Newcastle-upon-Tyne, United Kingdom. One of his earliest, formative memories was of a laboratory experiment he conducted in a grammar school class as twelve year old. In the experiment, he mixed a solution of aqueous ammonia (colorless) with copper sulfate (blue) in water, causing a reaction which would eventually produce "beautiful, well-formed crystals." From that moment "I was hooked," writes Bartlett, who yearned to know why the transformation took place. He could not have known that the event would vaguely foreshadow his famous experiment decades later in which he produced the world's first noble gas compound following a similarly stunning chemical reaction. He began to immerse himself in chemistry to the extent that he built his own makeshift laboratory in his parent's home, complete with flasks and beakers and chemicals he purchased at a local supply store. That curiosity carried over into academic success and eventually earned him a scholarship for his undergraduate education.
Bartlett attended King's College in Durham (U.K.), where he received his Bachelor of Science degree in 1954 and his doctorate in 1958. That year Bartlett was appointed a lecturer in chemistry at the University of British Columbia in Vancouver, Canada, where he remained until 1966, eventually reaching the rank of full professor. In 1966 he became a professor of chemistry at Princeton University while also serving as a member of the research staff at Bell Laboratories. In 1969, he joined the University of California, Berkeley, as a professor of chemistry, retiring in 1993. From 1969 to 1999 he also served as a scientist at the Lawrence Berkeley National Laboratory. Bartlett became a naturalized U.S. citizen in 2000.
Bartlett's fame goes beyond the inert gas research to include the general field of fluorine chemistry. He holds a special interest in the stabilization of unusually high oxidation states of elements and applying these states to advance chemistry. Bartlett is also known for his contributions toward understanding thermodynamic, structural, and bonding considerations of chemical reactions. He helped develop novel synthetic approaches, including a low-temperature route to thermodynamically unstable binary fluorides, including NiF4 and AgF3. He discovered and characterized many new fluorine compounds and also produced many new metallic graphite compounds, including some that show promise as powerful battery materials.
Scientists had always believed that noble gases, also known as inert or rare gases, were chemically unable to react. Helium, neon, argon, krypton, xenon, and radon (all gases at room temperature) were viewed as the "loners" of the Periodic Table. Their inertness became a basic tenet of chemistry, published in textbooks and taught in classrooms throughout the world. Conventional scientific wisdom held that the noble gas elements could not form compounds because their electronic structure was extremely stable. For all except helium, the maximum capacity of the outer electron shell of the noble gas atom is eight electrons. For helium, that limit is just two electrons. These electron arrangements are especially stable, leaving the noble gases without a tendency to gain or loose electrons. This led chemists to think of them as totally unreactive. A few chemists questioned the absolute inertness of the noble gases. Among those scientists were Walter Kossel in 1916 and Nobel-prize winning chemist Linus Pauling in 1933. They predicted that highly reactive atoms such as fluorine might form compounds with xenon, the heaviest of the noble elements and whose electrons, they observed, were not as tightly bound as those of the lighter gases.
In 1961 Neil Bartlett was teaching chemistry at the University of British Columbia in Vancouver, Canada. Some years earlier, while experimenting with fluorine and platinum, he had accidentally produced a deep-red solid whose exact chemical composition remained a mystery. With the assistance of his graduate student Derek Lohmann, he vigorously pursued the identity of the red solid. After much research, they eventually found that the known gaseous fluoride, platinum hexafluoride (PtF6), was able to oxidize oxygen and produce the red solid, which he and Lohmann had identified as O2+PtF6-.
What was most unusual about this compound was that it contained oxygen in the form of positively charged ions, although oxygen usually has a net negative charge. Oxygen normally pulls electrons from other atoms and is thus called an oxidizing agent or oxidant. But Bartlett believed that in this case, the PtF6 component was a more powerful oxidizing agent than even oxygen and was extracting electrons from oxygen, leaving oxygen with a net positive charge. Even though PtF6 was first prepared some years earlier by researchers at Argonne National Laboratory, its oxidizing power had not been recognized until Bartlett's research. It was this development that led Bartlett to theorize that if PtF6 could oxidize oxygen, then it might also be able to achieve the "impossible" task of oxidizing xenon, whose ionization potential (energy required to remove an electron) was very similar to that of oxygen.
In March of 1962, Bartlett concocted a simple experiment to test his hypothesis. He set up a glass apparatus containing PtF6, a red gas in one container and xenon, a colorless gas in an adjoining container, separated by a seal. Here's his recollection of the ensuing experiment, which he conducted while working alone in his laboratory: "Because my co-workers at that time (March 23, 1962) were still not sufficiently experienced to help me with the glassblowing and the preparation and purification of PtF6 [platinum hexafluoride] necessary for the experiment, I was not ready to carry it out until about 7 p.m. on that Friday. When I broke the seal between the red PtF6 gas and the colorless xenon gas, there was an immediate interaction, causing an orange-yellow solid to precipitate. At once I tried to find someone with whom to share the exciting finding, but it appeared that everyone had left for dinner!"
The reaction took place at room temperature "in the twinkling of an eye" and was "extraordinarily exhilarating," recalls Bartlett. He was certain that the orange-yellow solid was the world's first noble gas compound. But convincing others would prove somewhat difficult. The prevailing attitude was that no scientist could violate one of the basic tenets of chemistry: the inertness of noble gases. Bartlett insisted that he had, to the amusement and disbelief of some of his colleagues! The proof was in the new compound he had made. That orange-yellow solid was subsequently identified in laboratory studies as xenon hexafluoroplatinate (XePtF6), the world's first noble gas compound.
Within months, other chemists successfully repeated the experiment. Although the intricate chemical details behind the reaction would take years to clarify and the formula of the colorful solid was later modified as [XeF]+[PtF5]-, the significance of the experiment remained clear. Spurred by Bartlett's success, other scientists soon began to make new compounds from xenon and later, radon and krypton. With Bartlett's simple experiment, the old "law" of the unreactivity of the noble gases had been vanquished. The new field of noble gas chemistry, with its exciting possibilities, had been launched.
The American Chemical Society designated the research of Neil Bartlett on the noble gases as an International Historic Chemical Landmark in a ceremony on May 23, 2006 at the University of British Columbia in Vancouver. The text of the plaque on the campus of UBC reads in English and French: In this building in 1962 Neil Bartlett demonstrated the first reaction of a noble gas. The noble gas family of elements - helium, neon, argon, krypton, xenon, and radon - had previously been regarded as inert. By combining xenon with a platinum fluoride, Bartlett created the first noble gas compound. This reaction began the field of noble gas chemistry, which became fundamental to the scientific understanding of the chemical bond. Noble gas compounds have helped create anti-tumor agents and have been used in lasers.
Sir William Ramsay, together with Strutt, repeated the Cavendish experiment and identified argon as the unreactive species. The liquefaction of air in 1895 by Carl von Linde allowed Ramsay the further discovery of neon, krypton, and xenon. Extraterrestrial helium had been discovered earlier (in 1868), based on its spectral lines in the Sun. Ramsay realized that the new elements did not fit into the contemporary periodic system of the elements and suggested that they form a new group, bridging the alkali metals and the halogens. The last member of the family, radon was discovered in 1900 by Ernest Rutherford and Frederick Soddy as a decay product of radium.
The chemical inertness of the noble gases is based on their electronic structure. Each element has a completely filled valence shell. In fact their inertness helped to develop the key idea of a stable octet. The atomic sizes of the noble gas elements increase from top to bottom in the Periodic Table, and the amount of energy needed to remove an electron from their outermost shell, the ionization energy decreases in the same order. Within each period, however, the noble gases have the largest ionization energies, reflecting their chemical inertness. Based on increasing atomic size, the electron clouds of the spherical, nonpolar, atoms become increasingly polarizable, leading to stronger interactions among the atoms (van der Waals forces). Thus, the formation of solids and liquids is more easily attained for the heavier elements, as reflected in their higher melting points and boiling points. As their name implies, all members of the family are gases at room temperature and can, with the exception of helium, be liquefied at atmospheric pressure.
Until 1962 only physical inclusion compounds were known. Argon, krypton, and xenon form cage or clathrate compounds with water (clathrate hydrates) and with some organics such as quinol. The host molecules are arranged in such a way that they form cavities that can physically trap the noble gas atoms, referred to as "guests." The noble gas will be released upon dissolution or melting of the host lattice. In 1962 the first chemical noble gas compound, formulated as XePtF6, was synthesized by Neil Bartlett. This result spurred intense research activity and led to the discovery of numerous xenon and krypton compounds. In 2000 the formation of the first argon compound, argon fluorohydride (HArF), was reported by Leonid Khriachtchev and colleagues.
Science is frequently a collaborative discipline. But sometimes, one person, working alone in a laboratory, makes a stunning discovery, one that changes the way scientists look at their field. One such man was Neil Bartlett, who changed the face of chemistry while working alone in the chemistry department of the University of British Columbia. Bartlett demonstrated in 1962 that the "inertness" of the group VIII elements was a result of the reagents used in experiments, not a fundamental law of nature. Bartlett's proof that the rare gases were not chemically inert meant that all existing textbooks had to be rewritten. The American chemical society and the Canadian society for chemistry designated the work of Neil Bartlett and the reactive noble gasses an international historic chemical landmark at the University of British Columbia on May 23, 2006.
Neil Bartlett was born September 15, 1932 in Newcastle-upon-Tyne, United Kingdom. One of his earliest, formative memories was of a laboratory experiment he conducted in a grammar school class as twelve year old. In the experiment, he mixed a solution of aqueous ammonia (colorless) with copper sulfate (blue) in water, causing a reaction which would eventually produce "beautiful, well-formed crystals." From that moment "I was hooked," writes Bartlett, who yearned to know why the transformation took place. He could not have known that the event would vaguely foreshadow his famous experiment decades later in which he produced the world's first noble gas compound following a similarly stunning chemical reaction. He began to immerse himself in chemistry to the extent that he built his own makeshift laboratory in his parent's home, complete with flasks and beakers and chemicals he purchased at a local supply store. That curiosity carried over into academic success and eventually earned him a scholarship for his undergraduate education.
Bartlett attended King's College in Durham (U.K.), where he received his Bachelor of Science degree in 1954 and his doctorate in 1958. That year Bartlett was appointed a lecturer in chemistry at the University of British Columbia in Vancouver, Canada, where he remained until 1966, eventually reaching the rank of full professor. In 1966 he became a professor of chemistry at Princeton University while also serving as a member of the research staff at Bell Laboratories. In 1969, he joined the University of California, Berkeley, as a professor of chemistry, retiring in 1993. From 1969 to 1999 he also served as a scientist at the Lawrence Berkeley National Laboratory. Bartlett became a naturalized U.S. citizen in 2000.
Bartlett's fame goes beyond the inert gas research to include the general field of fluorine chemistry. He holds a special interest in the stabilization of unusually high oxidation states of elements and applying these states to advance chemistry. Bartlett is also known for his contributions toward understanding thermodynamic, structural, and bonding considerations of chemical reactions. He helped develop novel synthetic approaches, including a low-temperature route to thermodynamically unstable binary fluorides, including NiF4 and AgF3. He discovered and characterized many new fluorine compounds and also produced many new metallic graphite compounds, including some that show promise as powerful battery materials.
Scientists had always believed that noble gases, also known as inert or rare gases, were chemically unable to react. Helium, neon, argon, krypton, xenon, and radon (all gases at room temperature) were viewed as the "loners" of the Periodic Table. Their inertness became a basic tenet of chemistry, published in textbooks and taught in classrooms throughout the world. Conventional scientific wisdom held that the noble gas elements could not form compounds because their electronic structure was extremely stable. For all except helium, the maximum capacity of the outer electron shell of the noble gas atom is eight electrons. For helium, that limit is just two electrons. These electron arrangements are especially stable, leaving the noble gases without a tendency to gain or loose electrons. This led chemists to think of them as totally unreactive. A few chemists questioned the absolute inertness of the noble gases. Among those scientists were Walter Kossel in 1916 and Nobel-prize winning chemist Linus Pauling in 1933. They predicted that highly reactive atoms such as fluorine might form compounds with xenon, the heaviest of the noble elements and whose electrons, they observed, were not as tightly bound as those of the lighter gases.
In 1961 Neil Bartlett was teaching chemistry at the University of British Columbia in Vancouver, Canada. Some years earlier, while experimenting with fluorine and platinum, he had accidentally produced a deep-red solid whose exact chemical composition remained a mystery. With the assistance of his graduate student Derek Lohmann, he vigorously pursued the identity of the red solid. After much research, they eventually found that the known gaseous fluoride, platinum hexafluoride (PtF6), was able to oxidize oxygen and produce the red solid, which he and Lohmann had identified as O2+PtF6-.
What was most unusual about this compound was that it contained oxygen in the form of positively charged ions, although oxygen usually has a net negative charge. Oxygen normally pulls electrons from other atoms and is thus called an oxidizing agent or oxidant. But Bartlett believed that in this case, the PtF6 component was a more powerful oxidizing agent than even oxygen and was extracting electrons from oxygen, leaving oxygen with a net positive charge. Even though PtF6 was first prepared some years earlier by researchers at Argonne National Laboratory, its oxidizing power had not been recognized until Bartlett's research. It was this development that led Bartlett to theorize that if PtF6 could oxidize oxygen, then it might also be able to achieve the "impossible" task of oxidizing xenon, whose ionization potential (energy required to remove an electron) was very similar to that of oxygen.
In March of 1962, Bartlett concocted a simple experiment to test his hypothesis. He set up a glass apparatus containing PtF6, a red gas in one container and xenon, a colorless gas in an adjoining container, separated by a seal. Here's his recollection of the ensuing experiment, which he conducted while working alone in his laboratory: "Because my co-workers at that time (March 23, 1962) were still not sufficiently experienced to help me with the glassblowing and the preparation and purification of PtF6 [platinum hexafluoride] necessary for the experiment, I was not ready to carry it out until about 7 p.m. on that Friday. When I broke the seal between the red PtF6 gas and the colorless xenon gas, there was an immediate interaction, causing an orange-yellow solid to precipitate. At once I tried to find someone with whom to share the exciting finding, but it appeared that everyone had left for dinner!"
The reaction took place at room temperature "in the twinkling of an eye" and was "extraordinarily exhilarating," recalls Bartlett. He was certain that the orange-yellow solid was the world's first noble gas compound. But convincing others would prove somewhat difficult. The prevailing attitude was that no scientist could violate one of the basic tenets of chemistry: the inertness of noble gases. Bartlett insisted that he had, to the amusement and disbelief of some of his colleagues! The proof was in the new compound he had made. That orange-yellow solid was subsequently identified in laboratory studies as xenon hexafluoroplatinate (XePtF6), the world's first noble gas compound.
Within months, other chemists successfully repeated the experiment. Although the intricate chemical details behind the reaction would take years to clarify and the formula of the colorful solid was later modified as [XeF]+[PtF5]-, the significance of the experiment remained clear. Spurred by Bartlett's success, other scientists soon began to make new compounds from xenon and later, radon and krypton. With Bartlett's simple experiment, the old "law" of the unreactivity of the noble gases had been vanquished. The new field of noble gas chemistry, with its exciting possibilities, had been launched.
The American Chemical Society designated the research of Neil Bartlett on the noble gases as an International Historic Chemical Landmark in a ceremony on May 23, 2006 at the University of British Columbia in Vancouver. The text of the plaque on the campus of UBC reads in English and French: In this building in 1962 Neil Bartlett demonstrated the first reaction of a noble gas. The noble gas family of elements - helium, neon, argon, krypton, xenon, and radon - had previously been regarded as inert. By combining xenon with a platinum fluoride, Bartlett created the first noble gas compound. This reaction began the field of noble gas chemistry, which became fundamental to the scientific understanding of the chemical bond. Noble gas compounds have helped create anti-tumor agents and have been used in lasers.