Cesium

Typical wastewaters

• nuclear fuel reprocessing discharges to sea caesium-137 as the free hydrated Cs⁺ ion Sellafield discharged 1.3 TBq in 2022 against a permit limit of 17 TBq
• reactor accident water (Fukushima cooling and contaminated water) caesium-137 as Cs⁺, 823 Bq/L in one example at pH 12, in high salinity water from seawater cooling treated by a zeolite stage then nickel ferrocyanide, more than 1000 fold reduction
• reactor emissions and weapons test fallout in raw water supplies caesium-137 and caesium-134 as Cs⁺, passing conventional treatment guidance level 10 Bq/L; ion exchange 40 to 70 percent, reverse osmosis above 70 (WHO Table 9.4)

1 · Identity

Symbol, number

Cs, 55

Oxidation states in water

+1 only, as the large, weakly hydrated Cs⁺ ion; it does not hydrolyse, forms no precipitates in natural water and is held only by cation exchange on clays and by the lattice cavities of hexacyanoferrates.

Note

The element entry covers the metal and the caesium-137 gamma source. In water caesium is an inert cation; the whole story is radiocaesium and how selectively it can be exchanged.

2 · Occurrence in water

Natural sources

Traces from weathering of potassium minerals; stable caesium in water is at the microgram per litre level or below and is not sourced here.

Anthropogenic sources

Caesium-137 and caesium-134 released as fission products in reactor emissions, nuclear weapons tests and accidents (WHO Table 9.2); reprocessing discharges (Sellafield discharged 1.3 TBq of caesium-137 to sea in 2022 against a permit limit of 17 TBq); the Fukushima treatment campaign has processed 482,000 m₃ of contaminated water (Rauwel 2019).

3 · Speciation

Cs⁺ is a free hydrated cation at every pH; it is not precipitated, complexed or oxidised. Selectivity comes from size: Prussian blue analogues capture Cs⁺ in a cavity created by a hexacyanoferrate vacancy whose size matches the hydrated radius of Cs⁺, exchanging it for K⁺ or H⁺; competition follows K⁺ above Ca²⁺ above Mg²⁺ above Na⁺ (Rauwel 2019).

Solubility

All common caesium salts are freely soluble.

Hydrolysis

None.

Complexation

Negligible; the hexacyanoferrate lattice is a size selective host, not a ligand.

Precipitates

None in natural water; caesium hexacyanoferrates are the engineered solids.

4 · Role in treatment

5 · Removal and control

6 · Analytics

Sampling pitfalls

Acidify to keep Cs⁺ off the bottle wall and sediment; note the potassium-40 background when gross beta is used as the screen (WHO subtracts potassium-40 from measured potassium).

7 · Regulatory limits

Limits change, and many are set locally. Treat these as the published values to start from, not as your compliance target: check the standard in force at your site and the numbers written into your own permit.

8 · Health and environmental effects

Toxicity

Stable caesium is of low toxicity and unregulated; caesium-137 is a whole body gamma and beta dose, distributed like potassium in soft tissue (the guidance level is 10 Bq/L, WHO).

Bioaccumulation

Follows potassium into muscle of fish and animals; the food chain, not water, is the dominant pathway after fallout (not sourced here beyond WHO's fission product note).

Ecotoxicity

not relevant

Flags

• Stable caesium occurrence in water was not sourced; only radiocaesium is quantified.
• The hexacyanoferrate exchange equation is an idealised formula written here from the review's description of K⁺ exchange.
• The Fukushima figures (823 Bq/L, 1000 fold, 482,000 m₃) are quoted from a materials review, not from the plant operator.
• Sellafield discharge is a yearly load, not a water concentration.

Gaps

• No source read gives stable caesium in groundwater, surface water, seawater or wastewater.
• Zeolite selectivity coefficients and the clay (illite frayed edge) sorption that governs caesium in sediments are not sourced.
• The hexacyanoferrate dissolution and cyanide release at high pH is a known drawback not quantified in the sources read.
• No GCC document mentions caesium.
• No exchange stoichiometry or selectivity coefficient for caesium on zeolite or strong acid resin was read; the resin equation is the standard monovalent exchange written in the book's notation.

Sources

Identity

Name and symbol

Cesium, Cs

Atomic number

55 protons

Position

group 1 · period 6 · s-block · alkali metal

CAS number

7440-46-2

Atomic structure

Atomic mass

132.905 u

Electron configuration

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s¹
[Xe] 6s¹

Electrons per shell

2, 8, 18, 18, 8, 1

Valence electrons

1 outer shell

Physical properties

State at room temperature

Solid

Melting point

301.59 K (28.44 °C)

Boiling point

944 K (670.85 °C)

Density

1.93 g/cm3

Appearance

pale gold

Thermal conductivity

35.9 W/(m·K)

Electrical resistivity

205 nΩ·m at 20 °C

Electrical conductivity

4.88 MS/m

Crystal structure

bodycentredcubic

Molar heat capacity

32.21 J/(mol·K)

Chemical properties

Oxidation states

+1

Electronegativity

0.79 (Pauling Scale)

Ionisation energy

3.894 eV
1st 375.7, 2nd 2,234.3, 3rd 3,400 kJ/mol

Electron affinity

0.472 eV

Atomic radius

empirical 244, covalent 244, van der Waals 343 pm

Ionic radius

Cs⁺ 167 pm

Reactivity

The most electropositive and most alkaline stable element: one 6s electron so easily lost that the soft gold-coloured metal is pyrophoric in air and explosive in water, and it must be kept sealed under inert gas or vacuum.

with water

Reacts explosively with cold water, and even with ice down to -116 C, giving caesium hydroxide and hydrogen: $\ce{2 Cs + 2 H2O -> 2 CsOH + H2}$.

with oxygen, air

Ignites spontaneously in air and burns mainly to the superoxide: $\ce{Cs + O2 -> CsO2}$, with Cs2O, Cs2O2 and coloured suboxides in limited oxygen.

with acids

Reacts explosively with dilute acids to the caesium salt and hydrogen: $\ce{2 Cs + 2 HCl -> 2 CsCl + H2}$.

with halogens

Burns in the halogens to the caesium halides: $\ce{2 Cs + Cl2 -> 2 CsCl}$.

Typical compounds

• CsOH caesium hydroxide the strongest base known; attacks glass
• CsCl caesium chloride density-gradient centrifugation; its own crystal structure type
• CsF caesium fluoride hygroscopic fluoride source in organic chemistry
• CsO₂ caesium superoxide main product of burning the metal
• HCOOCs caesium formate dense oil-well drilling fluid, the largest use
• CsI caesium iodide scintillator for radiation detection

Occurrence, production and use

Crustal abundance

3 milligrams per kilogram

Oceanic abundance

3×10-4 milligrams per liter

Occurrence and sources

Cesium, an alkali metal, occurs in lepidolite, pollucte (a hydrated silicate of aluminum and cesium), and in other sources. One of the world's richest sources of cesium is located at Bernic Lake, Manitoba. The deposits are estimated to contain 300,000 tons of pollucite, averaging 20% cesium.

It can be isolated by elecytrolysis of the fused cyanide and by a number of other methods. Very pure, gas-free cesium can be prepared by thermal decomposition of cesium azide.

• pollucite (caesium aluminosilicate, 5 to 32 percent Cs2O) Tanco (Bernic Lake), Manitoba; Bikita, Zimbabwe (depleted 2018); Sinclair, Australia (mined out 2019); Alaska, Maine, South Dakota
• lepidolite and other lithium pegmatite minerals by-product of lithium production; Namibia project due 2026
• brines and geothermal waters Chile, China; geothermal systems in China, Germany, India
• crustal and oceanic abundance about 3 ppm (BGS figure via RSC); 3 mg/kg crust and 0.0003 mg/L seawater (PubChem)

Extraction, production

Pollucite processing to caesium chemicals

caesium minerals are the feedstock for caesium compounds and metal; caesium chloride is named by USGS as the intermediate for metal production; no chemistry is stated

Electrolysis of molten caesium cyanide (Setterberg)

RSC states the metal was first obtained this way; the anode product is not named, so no equation is written

Uses

Cesium has the second lowest melting point of all metallic elements, which limits its uses. Cesium readily combines with oxygen and is used as a getter, a material that combines with and removes trace gases from vacuum tubes. Cesium is also used in atomic clocks, in photoelectric cells and as a catalyst in the hydrogenation of certain organic compounds. Since it is easily ionized and has a high mass, cesium ions may one day be used as a propellant in ion engines on spacecraft.

Cesium reacts violently with water and ice, forming cesium hydroxide (CsOH). Cesium hydroxide is the strongest base known and will attack glass. Cesium chloride (CsCl) and cesium nitrate (CsNO3) are cesium's most common compounds and are primarily used in the production of other chemicals.

Because of it has great affinity for oxygen, the metal is used as a "getter" in electron tubes. It is also used in photoelectric cells, as well as a catalyst in the hydrogenation of certain organic compounds.

The metal has recently found application in ion propulsion systems. Cesium is used in atomic clocks, which are accurate to 5 s in 300 years. Its chief compounds are the chloride and the nitrate.

• Oil and gas drilling: caesium formate brine for high-pressure, high-temperature well drilling and completion the primary application by gross weight; about 11,000 t in use (usgs-mcs2025-cesium)
• Chemicals: caesium-promoted vanadium catalysts for high-purity sulfuric acid; caesium catalysts for methyl methacrylate; caesium carbonate in alkylation
• Electronics, optics and timekeeping: caesium atomic clocks for GPS, telecom and internet synchronisation; caesium iodide and bromide scintillators, X-ray image intensifiers, infrared optics; photoelectric cells, vacuum tubes, special optical glass
• Medicine and radiation: caesium-131 brachytherapy sources; caesium-137 industrial gauges and sterilisation, being phased out of blood irradiators by 2027; caesium chloride in isopycnic centrifugation
• Water treatment and energy: caesium sulfates in water treatment and fuel cells; caesium hydroxide electrolyte in alkaline batteries

Safety, toxicity

GHS classification, signal word Danger

• H260 In contact with water releases flammable gases which may ignite spontaneously Substances and mixtures which in contact with water, emit flammable gases
• H314 Causes severe skin burns and eye damage Skin corrosion/irritation
• H318 Causes serious eye damage Serious eye damage/eye irritation

Discovery and name

Discovered by

Gustav Kirchhoff and Robert Bunsen

Discovered

1860

First isolated

Carl Setterberg

Named by

not in sources

Origin of the name

bluish grey, for its spectral colours

The metal is characterized by a spectrum containing two bright lines in the blue along with several others in the red, yellow, and green wavelengths. It is silvery white, soft, and ductile. It is the most electropositive and most alkaline element.

Cesium, gallium, and mercury are the only three metals that are liquid at room temperature. Cesium reacts explosively with cold water, and reacts with ice at temperatures above -116C. Cesium hydroxide, the strongest base known, attacks glass.