The Chemistry of the Rare Gases
|Discovery of Rare Gases||Oxidation Numbers and Position in the Periodic Table|
|Synthesis of Xenon Compounds||Rare-Gas Compounds in Lasers|
In 1892 Lord Rayleigh found that oxygen was always 15.882 times more dense than hydrogen, no matter how it was prepared. When he tried to extend this work to nitrogen, he found that nitrogen isolated from air was denser than nitrogen prepared from ammonia. William Ramsey attacked this problem by purifying a sample of nitrogen gas to remove any moisture, carbon dioxide, and organic contaminants. He then passed the purified gas over hot magnesium metal, which reacts with nitrogen to form the nitride.
When he was finished, Ramsey was left with a small residue of gas that occupied roughly 1/80th of the original volume. He excited this gas in an electric discharge tube and found that the resulting emission spectrum contained lines that differed from those of all known gases. After repeated discussions of the results of these experiments, Rayleigh and Ramsey jointly announced the discovery of a new element, which they named argon from the Greek word meaning the "lazy one" because this gas refused to react with any element or compound they tested.
Argon did not fit into any of the known families of elements in the periodic table, but its atomic weight suggested that it might belong to a new group that could be inserted between chlorine and potassium. Shortly after reporting the discovery of argon in 1894, Ramsey found another unreactive gas when he heated a mineral of uranium. The lines in the spectrum of this gas also occurred in the spectrum of the sun, which led Ramsey to name the element helium (from the Greek helios, "sun").
Experiments with liquid air led Ramsey to a third gas, which he named krypton ("the hidden one"). Experiments with liquid argon led him to a fourth gas, neon ("the new one"), and finally a fifth gas, xenon ("the stranger").
These elements were discovered between 1894 and 1898. Because Moissan had only recently isolated fluorine for the first time and fluorine was the most active of the known elements, Ramsey sent a sample of argon to Moissan to see whether it would react with fluorine. It did not. The failure of Moissan's attempts to react argon with fluorine, coupled with repeated failures by other chemists to get the more abundant of these gases to undergo chemical reaction, eventually led to their being labeled inert gases. The development of the electronic theory of atoms did little to dispel this notion because it was obvious that these gases had very symmetrical electron configurations. As a result, these elements were labeled "inert gases" in almost every textbook and periodic table until about 30 years ago.
In 1962 Neil Bartlett found that PtF6 was a strong enough oxidizing agent to remove an electron from an O2 molecule.
Bartlett realized that the first ionization energy of Xe (1170 kJ/mol) was slightly smaller than the first ionization energy of the O2 molecule (1177 kJ/mol). He therefore predicted that PtF6 might also react with Xe. When he ran the reaction, he isolated the first compound of a Group VIIIA element.
A few months later, workers at the Argonne National Laboratory near Chicago found that Xe reacts with F2 to form XeF4. Since that time, more than 200 compounds of Kr, Xe, and Rn have been isolated. No compounds of the more abundant elements in this group (He, Ne, and Ar) have yet been isolated. However, the fact that elements in this family can undergo chemical reactions has led to the use of the term rare gases rather than inert gases to describe these elements.
Compounds of xenon are by far the most numerous of the rare-gas compounds. With the exception of XePtF6, rare-gas compounds have oxidation numbers of +2, +4, +6, and +8, as shown by the examples cited in the table below.
Compounds of Xenon and their Oxidation Numbers
There is some controversy over whether the rare gases should be viewed as having the outermost shell of electrons filled (in which case they should be labeled Group VIIIA) or empty (in which case they should be labeled Group 0). We believe these elements should be labeled Group VIIIA because they behave as if they contribute eight valence electrons when they form compounds.
The synthesis of most xenon compounds starts with the reaction between Xe and F2 at high temperatures (250400oC) to form a mixture of XeF2, XeF4, and XeF6.
The positively charged XeFn+ ions are then made by reacting XeF2, XeF4, or XeF6 with either AsF5, SbF5, or BiF5.
Oxides of xenon, such as XeOF2, XeOF4, XeO2F2, XeO3F2, XeO2F4, XeO3, and XeO4 are prepared by reacting XeF4 or XeF6 with water. The XeO64- ion, for example, is produced when XeF6 dissolves in strong base.
|2 XeF6(s)||+||16 OH-(aq)||XeO64-(aq)||+||Xe(g)||+||O2(g)||+||12 F-(aq)||+||8 H2O(l)|
Some xenon compounds are relatively stable. XeF2, XeF4, and XeF6, for example, are stable solids that can be purified by sublimation in a vacuum at 25oC. XeOF4 and Na4XeO6 are also reasonably stable. Others, such as XeO3, XeO4, XeOF2, XeO2F2, XeO3F2, and XeO2F4, are unstable compounds that can decompose violently.
The principal use of rare-gas compounds at present is as the light-emitting component in lasers. Mixtures of 10% Xe, 89% Ar, and 1% F2, for example, can be "pumped," or excited, with high-energy electrons to form excited XeF molecules, which emit a photon with a wavelength of 354 nm.