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In 1860, Bunsen and Kirchhoff in Germany discovered it through spectral analysis. It is an element of Group IA in Period 6, with the symbol Cs. It is commonly found in lithium ores in nature; a small amount of cesium chloride, CsCl, is also present in carnallite. It has only one stable form, with an oxidation state of +I. Its elemental form is a silver-white metal that is soft and lightweight, malleable, and has a low melting point of (28.40±0.01°C); it is one of the three metals that are liquid at room temperature (the other two being mercury and gallium). It has the lowest electronegativity; it is the most reactive element among the alkali metals, and it reacts violently with oxygen to form a mixture of various oxides. In humid air, the heat generated by this oxidation is sufficient to melt and ignite it. It does not react with nitrogen, but at high temperatures it can react with hydrogen to form fairly stable hydrides. It reacts violently with halogens, causing ignition, and it explodes when reacting with sulfur and phosphorus. It also reacts violently with water, even with ice at temperatures as low as -116°C. This is due to its large ionic radius and strong positive charge. Cesium also undergoes reactions similar to those of other alkali metals, but it is more reactive. CsCl is the main compound. Cesium releases electrons when exposed to light. Cesium is mainly used in the manufacture of phototubes, spectrometers, scintillation counters, electron tubes, military infrared signal lights, as well as various optical and testing instruments. Its compounds are used in the production of glass and ceramics, and can also serve as absorbents in carbon dioxide purification devices and as getters in electron tubes; they are also utilized in trace chemistry. In medicine, cesium salts can also be used as anti-shock agents after taking arsenic-containing drugs, while the isotope cesium-137 can be used to treat cancer. Cesium can be produced by reducing CsCl with calcium.