|Name, symbol||Caesium, Cs|
|Caesium in the periodic table|
|Standard atomic weight (Ar)||132.90545196(6)|
|Group, block||(alkali metals); s-block|
|Electron configuration||[Xe] 6s1|
|2, 8, 18, 18, 8, 1|
|Melting point||301.7 K (28.5 °C, 83.3 °F)|
|Boiling point||944 K (671 °C, 1240 °F)|
|Density near r.t.||1.93 g/cm3|
|when liquid, at||1.843 g/cm3|
|Critical point||1938 K, 9.4 MPa|
|Heat of fusion||2.09 kJ/mol|
|Heat of||63.9 kJ/mol|
|Molar heat capacity||32.21 J/(mol·K)|
|Oxidation states||+1, −1 (a strongly basic oxide)|
|Electronegativity||Pauling scale: 0.79|
1st: 375.7 kJ/mol |
2nd: 2234.3 kJ/mol
3rd: 3400 kJ/mol
|Atomic radius||empirical: 265 pm|
|Covalent radius||244±11 pm|
|Van der Waals radius||343 pm|
|Crystal structure||Body-centred cubic (bcc)|
|Thermal expansion||97 µm/(m·K) (at 25 °C)|
|Thermal conductivity||35.9 W/(m·K)|
|Electrical resistivity||205·10-9 Ω·m (at 20 °C)|
|Young's modulus||1.7 GPa|
|Bulk modulus||1.6 GPa|
|Brinell hardness||0.14 MPa|
|CAS Registry Number||7440-46-2|
|Naming||From Latin caesius - sky blue, for its spectral colours|
|Discovery||Robert Bunsen and Gustav Kirchhoff (1860)|
|First isolation||Carl Setterberg (1882)|
Caesium or cesium is an alkali metal with the chemical symbol Cs and atomic number 55. It is the heaviest, stable alkali metal.
Caesium is an extremely reactive metal an will spontaneously ignite in air to form caesium oxides and hydroxides. The reaction with water is explosive, capable of shattering the water's glass container.
- Cs + H2O → CsOH + ½ H2
Caesium also forms two unusual acid nitrates, CsNO3·HNO3 and CsNO3·2HNO3
Caesium hydroxide is extremely corrosive and attack glass much faster than the other alkali hydroxides.
Caesium is a soft alkali metal, silvery-gold in color. It has a very low melting point of 28 °C, meaning it can be a liquid at near room temperature or if held in one's hand. It is the least electronegative element and is extremely reactive and even pyrophoric. It will react with water even at temperatures as low as −116 °C.
Cesium ampoules can be purchased from GalliumSource, however it is extremely expensive, a 100 g ampoule is 2400 $ and a 1 g one is 140$. Another seller, SmartElements sells cesium ampoules cheaper, a 10 g ampoule is 179 € while a 1 g ampoule is only 79 €.
Very small amounts of cesium alloys are used in the cathode of the electron gun from the cathode ray tube TVs. The exact composition of the alloy varies, depending on the generation of tube, as not all of them have cesium.
Cesium compounds, such as cesium formate, are used as drilling fluids.
Because it's extremely reactive, isolating pure caesium metal is extremely difficult. One way is to reduce caesium halides with a reactive metal such as calcium, barium, or lithium at 700-800 °C, followed by distillation of the caesium metal. YouTube vlogger thunderf00t has performed this preparation, as it seems to be much cheaper to perform this extraction than it is to buy the metal (but it is far more convenient to just buy the metal).
Electrolytic preparation of elemental cesium is extremely difficult. Caesium chloride for example, melts at at 645 °C, while cesium metal boils at 671 °C, so without a sensitive temperature controller there's a risk of boiling the metal. This boiling, however, tends to drive isolation reactions forward by increasing entropy.
- Alkali water explosion
- Caesium hydroxide synthesis
- Cesium auride synthesis
Caesium metal is extremely reactive and pyrophoric in air. Cesium will not usually catch fire just by being exposed to air, but friction, heating, or exposure to water can trigger a fire or explosion.
While it can be stored under mineral oil, it will oxidize much faster than lithium, sodium or potassium, so it is best stored in vacuum containers or argon ampoules. The best way to handle metallic cesium is in a glove box filled with inert gas, such as argon.
Cesium metal cannot be safely neutralized by dissolving it in isopropanol, like sodium, as the reaction is just as violent as the one with water. It is best to leave it in air (or in a CO2 atmosphere, as there's a less risk of fire) to turn into oxide, hydroxide and carbonate, that can be safely disposed of. Due to its rarity, it's best to recycle caesium.
- Safety in the Chemistry and Biochemistry Laboratory By André Picot, P. Grenouillet, p. 213