noble gases

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 [url=]argon[/url] 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 [url=]radium[/url] .

Physical and Chemical Properties
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 [url=]helium[/url] , be liquefied at atmospheric pressure.(from chemistryrexplain)

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