Radium bromide

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Radium bromide[1][2]
Ra bromid.jpg
Radium bromide
IUPAC name
radium bromide
Other names
radium bromide
3D model (JSmol)
ECHA InfoCard 100.030.066
Molar mass 385.782 g/mol
Appearance white orthorhombic crystals
Density 5.79 g/cm3
Melting point 728 °C (1,342 °F; 1,001 K)
Boiling point 900 °C (1,650 °F; 1,170 K) sublimes
70.6 g/100 g at 20°C
Related compounds
Other anions
Radium chloride
Main hazards Radioactive, highly toxic
GHS pictograms The exploding-bomb pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)The skull-and-crossbones pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)The health hazard pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)The environment pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 4: Very short exposure could cause death or major residual injury. E.g., VX gasReactivity code 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g., fluorineSpecial hazard RA: Radioactive. E.g., plutoniumNFPA 704 four-colored diamond
Special hazard RA: Radioactive. E.g., plutonium
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Radium bromide is the bromide salt of radium, with the formula RaBr2. It is produced during the separation of radium from uranium ore. This inorganic compound was discovered by Pierre and Marie Curie in 1898, which sparked a huge interest in radiochemistry, especially radiotherapy. Since radium oxidizes rapidly in air and in water, the salt form is the preferred chemical form to work with.[3] Even though the salt form is more stable, radium bromide is still a dangerous chemical that can explode under certain conditions.[4]


After the Curies discovered radium, many people tried to get as much radium as possible to use it for medical purposes. However, it was not long until radium and radium salts were thought of as magical substances and were added to food, drinks, clothing, toys and more.[5] Respectable journals in the early 1900s published statements claiming that radium had no toxic effect.

Scientists were racing against each other trying to be the first to successfully use radium to cure health problems. Radium salts, including radium bromide, were most often used by placing the chemical in a tube that was then passed over or inserted into the problem site or tissue in the body. Many of the first scientists to try to determine radium's uses were affected by their exposure to the radioactive material. Pierre Curie went so far as to self-inflict a severe chemical skin reaction by applying a radium source directly to his forearm.[6] All types of therapeutic tests were performed for different skin diseases including eczema, lichen and psoriasis. It was not thought of until later that radium could be used to treat cancerous diseases.

The main problem with the explosion of interest in radium was the lack of radium on earth itself. In 1913, it was reported that the Radium Institute had four grams of radium total, which at the time was more than half the world supply.[5] Countries all over the world set out to make as much radium as possible, a time consuming and expensive task. It was reported in Science magazine in 1919 that the United States had produced approximately 55 grams of radium since 1913, which was also more than half the radium produced in the world at the time.[7] A principal source for radium is pitchblende, which holds a total of 257 mg of radium per ton of U3O8.[3] Obviously with so little product recovered from such a large amount of material, it was difficult to find a large quantity of radium. This was the reason radium bromide became one of the most expensive materials on earth. It was desired by many all around the world and there was very little to go around. In 1921, it was stated in Times magazine that radium cost 17,000,000,000 Euros per ton compared to gold at 208,000 Euros and coal, as diamond, at 400,000,000 Euros.[5]

A common reason that radium was added into almost everything at the time, including water and toys, was that people saw radium as fun, not harmful. Radium bromide was found to induce phosphorescence at normal temperatures.[8] This led to the US army making and supplying luminous watches to all soldiers. It also allowed for the invention of a spinthariscope, which soon became the item everyone wanted in order to show that they were up to date on the new science discoveries at the time.[9]


Radium bromide is a luminous chemical that causes the air surrounding it, even when encased in a tube, to glow and demonstrate all bands of the nitrogen spectrum. It is possible that the alpha radiation on the nitrogen in the air causes this luminescence. Radium bromide is also very dangerous since the crystals occasionally explode. Helium gas evolved from alpha particles can accumulate within the crystals which cause them to weaken and rupture. Another property of radium bromide is its ability to crystallize from an aqueous solution. It forms a dihydrate, very similar to barium bromide.[4]


The chemical radium bromide is not found in nature; it is created. To extract radium from uranium or pitchblende ores, the most common practice is called the Curie method which it involves two major stages. The first stage is a chemical treatment to concentrate the radium as a combination of radium and barium. This is done by treating the ore with a barium salt and sulfuric acid which help the uranium, iron, copper, and other components become soluble. A residue containing gangue, barium, radium, and lead sulfates will be left over. The mixture will then be treated with sodium chloride to remove lead and convert radium and barium into carbonates that are insoluble in hydrochloric acid. The second step requires fractional crystallization to separate the barium from the radium.[3] Because of the differences in solubility between radium and barium in bromide or chlorine, those two chemicals are chosen for the fractional crystallization. Once the radium is separated, radium bromide can be formed by dehydrating the dihydrate with a dry air stream at 200°C.[4] Another method is to heat radium chloride to red heat with a dry hydrogen bromine gas stream.[4]


Radium and radium salts were commonly used for treating cancer. However, many treatments that were used in the early 1900s are not used anymore because of the harmful effects radium bromide exposure caused. Some examples are anaemia, cancer, and genetic mutations.[5] Health effects from harmful amounts of radium exposure are tricky to diagnose because it is a cumulative health hazard. It cannot be diagnosed until years after exposure. In summary, there are not many practical uses for radium bromide anymore.

See also[edit]


  1. ^ Lide, David R. (1998). Handbook of Chemistry and Physics (87 ed.). Boca Raton, FL: CRC Press. pp. 4–78. ISBN 0-8493-0594-2.
  2. ^ Chemical Compounds (Inorganic); B-Table, Record No. 2630. International Critical Tables of Numerical Data, Physics, Chemistry and Technology (1st Electronic Edition). 2000
  3. ^ a b c Babcock, A.B., Jr. Survey of Processes for Radium Recovery from Pitchblende Ores. AEC Research and Development Report. 23 Feb 1950. No. NYO—112
  4. ^ a b c d Kirby,H.W; Salutsky, Murrell L. The Radiochemistry of Radium. Energy Citations Database Dec 1964.[1]
  5. ^ a b c d Harvie, David I. The Radium Century. Endeavor 1999 Vol. 23, Issue 3: 100-105
  6. ^ Dutreix, Jean; Pierquin, Bernard; Tubiana, Maurice. The Hazy Dawn of Brachytherapy. Radiotherapy and Oncology (49) 1998 223-232
  7. ^ Voil, Charles H. Radium Production. Science 17 March 1919 Vol. 49, No 1262: 227-228
  8. ^ 100 and 50 years ago. Nature 24 July 2003 Vol. 424, Issue 6927: 381
  9. ^ Schwarcz, Joe. A Dazzling display in a little jar. The Gazette: Saturday Extra; The Right Chemistry pg B5