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When I teach introductory inorganic chemistry, one of my favorite experiments is to toss alkali metals into a beaker of water. Lithium sizzles, sodium sparks, and potassium bursts into flames, so merely holding up a vial of cesium causes quite a stir in the classroom. Although many of the students might like to see the violent explosion that would ensue when cesium hits water, those in the front of the room are especially relieved when I just pass around the sealed cesium vial.

I then add phenolphthalein to one of the beakers, producing the characteristic pink color of base and explain that the reaction of alkali metal with water forms alkali hydroxide and hydrogen. The fireworks are created from the exothermicity of the reaction igniting the hydrogen gas. This occurs much more rapidly as one goes down the column of alkali metals, since as size increases the ionization potential decreases. Thus, cesium is the most reactive of the alkali metals. Note that the alkali-in-water experiment is carried out wearing safety glasses and with a clear plastic blast shield to protect the students.

When the undergraduates actually hold the sealed glass vial containing cesium, most are surprised to see a golden reflective material, the only other metal besides gold and copper that is not silvery in color. Even more remarkable is that the cesium begins to melt as it makes it way around the classroom. Cesium melts just above room temperature at 28.6 ºC, giving it the second lowest melting point relative to mercury (m.p. = –38.7 ºC). Cesium readily alloys with the other alkali metals, and a composition of 41% Cs, 47% K, and 12% Na produces the lowest melting metallic alloy known (m.p. = –78 ºC).

Cesium was the first element to be discovered using spectroscopic means by Robert W. Bunsen and Gustav R. Kirchhoff in 1860, the year after they invented the spectroscope. Cesium, from the Latin caesius, meaning heavenly blue, was named after the color of the most prominent line in its spectrum (= 455.5 nm). It can be identified qualitatively in a flame test from the pale violet light given off by the electronic transitions in the excited metal atoms. Natural cesium consists of a single stable isotope, Cs–133. It occurs chiefly as a hydrated aluminosilicate mineral known as pollucite, 2Cs2O2Al2O39SiO2H2O, mined in the Bernic Lake region of Manitoba.

Cesium is the largest naturally occurring element; it has an atomic radius of 2.65 Å. With its low ionization potential (376 kJ per mol), it readily gives up its only valence electron to produce ionic salts. One of these, cesium chloride, forms a basic structure type that I discuss in the introductory inorganic course. Although most ionic lattices consist of an array of larger close-packed anions with smaller cations in the interstices, in CsCl the cesium cations (1.88 Å) are actually slightly larger than the chlorine anions (1.67 Å). This leads to a structure that can be described most accurately as simple cubic cesium cations with chlorine anions occupying every eight-coordinate cubic site. This looks analogous to a body-centered cubic structure composed of just one type of atom.

BUBBLE, BUBBLE ... Cesium reacts with water-phenolphthalein in solution.
Cesium, like the other alkali metals, readily dissolves in liquid ammonia to produce "solvated electrons" and cesium cations. The solvated electrons possess a characteristic blue color caused by transmitted light and are paramagnetic. At high concentrations the cesium-ammonia solution becomes quite viscous and turns bright gold due to reflected light. On exposure to oxygen, cesium readily forms seven different oxides ranging from Cs7O to CsO3.

Researchers at Bell Labs in the early 1990s reported that doping buckminsterfullerene, C60, with the alkali metal potassium led to room-temperature conductivity and low-temperature superconductivity. In collaboration with physics and chemistry colleagues, my group produced the first pure samples of K3C60 and determined the structure to be face-centered cubic C60 with potassium in all octahedral and tetrahedral sites. K3C60 has a superconducting transition temperature of 19.3 K that surprisingly decreases as pressure is increased. To apply "negative" pressure, we replaced K with Rb, which produced an increase in the transition temperature to 29.6 K for Rb3C60. Unfortunately, substituting three cesiums for the rubidiums led to disproportionation back to C60 and the insulating body-centered cubic phase, Cs6C60, rather than a superconducting material. However, when a small amount of Cs is substituted into Rb3C60, a superconducting transition temperature above 30 K is achieved, the highest known for a carbon-based material.

Therefore, a small amount of cesium goes a long way, whether in doping C60 or in a vial to pass around class. On a warm day, cesium appears as liquid gold. Just watch out for its explosive personality!

Richard Kaner is a professor of chemistry and biochemistry at the University of California, Los Angeles. He is a recipient of the ACS Exxon Fellowship in Solid-State Chemistry; the Buck-Whitney Award from the ACS Eastern New York Section; and Dreyfus, Guggenheim, Packard, and Sloan Fellowships.


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Copyright © 2003 American Chemical Society

Name: From the Latin caesius, heavenly blue. The metal is characterized by two bright blue lines in its spectrum.
Atomic mass: 132.91.
History: Discovered in 1860 by German chemists Robert Bunsen and Gustav Kirchhoff.
Occurrence: Primarily obtained from the mineral pollucite.
Appearance: Silvery gold, soft, ductile metal.
Behavior: The most electropositive and alkaline element. Liquid around room temperature. Cesium reacts explosively with cold water, and reacts with ice at temperatures above –116 ºC. Cesium is fairly toxic.
Uses: Used as a catalyst promoter, as a "getter," in radiation monitoring equipment, and in atomic clocks.

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