Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction

As for applications, rubidium is used in a few electronic devices, as a frequency reference in atomic clocks, and to estimate the age of rocks. The element's future is full of potential if it could be more easily isolated. For example, rubidium compounds are being studied for medical uses, such as a potential antidepressant akin to lithium. A rubidium ionic crystal, RbAg₄I₅, has high room-temperature conductivity and could be used in thin films for batteries. Another prospect is Rb₃C₆₀, which is a potential superconducting material.

That brings us to the coup de gr^ace of this essay: what I did on my summer vacation. My family went to Spain to visit relatives on my wife's side. In Ourense, Galicia, in northwest Spain, we came across rubidium at a set of hot springs.

The springs are located next to the Miño River, which flows through downtown Ourense. In the first century, the Romans established a town around the springs and built a strategic bridge over the river. The mineral water has been used since that time for bathing and drinking, and we witnessed some of the locals coming to fill jugs with the water to take home. In the late 1800s, a doctor analyzed the water, and his results are listed on a plaque next to one of the springs' fountains. Among the ions are 0.16 mg per L rubidium, 102.2 mg per L sodium, and 11.2 mg per L calcium. These values are typical for the concentration of various ions in mineral waters worldwide.

Stephen K. Ritter is a senior editor for C&EN's Science, Technology & Education group. He likes finding chemistry in unexpected places.

RUBIDIUM AT A GLANCE

Name: From the Latin rubidius, deepest red, the color of its spectral lines.

Atomic mass: 85.47.

History: Discovered in 1861 by German chemists Robert Bunsen and Gustav Kirchoff while studying the mineral lepidolite.

Occurrence: Soft, silvery white metal. Rubidium can be liquid at ambient temperature if at or above 40 °C.

Appearance: Colorless, odorless gas at room temperature; pale blue as a liquid and a solid; faintly blue, brackish odor as gaseous ozone.

Behavior: One of the most electropositive and alkaline elements. It ignites spontaneously in air and reacts violently with water, setting fire to the liberated hydrogen. It colors flame yellowish-violet. Rubidium can be toxic by ingestion.

Uses: Used in cathode-ray tubes and as a "getter" for vacuum systems.

CESIUM

RICHARD KANER, UNIVERSITY OF CALIFORNIA, LOS ANGELES

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.