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Strontium carbonate (SrCO3) is the carbonate salt of strontium that has the appearance of a white or grey powder. It occurs in nature as the mineral strontianite.

Strontium carbonate
SrCO3.jpg
Names
IUPAC name
Strontium carbonate
Other names
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.015.131
EC Number 216-643-7
RTECS number WK8305000
UNII
Properties
SrCO3
Molar mass 147.63 g/mol
Appearance White powder
Odor Odorless
Density 3.74 g/cm3
Melting point 1,494 °C (2,721 °F; 1,767 K) (decomposes)
0.0011 g/100 mL (18 °C)
0.065 g/100 mL (100 °C)
Solubility soluble in ammonium chloride
slightly soluble in ammonia
−47.0·10−6 cm3/mol
1.518
Structure
rhombic
Hazards
Safety data sheet External MSDS data
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentineReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
1
0
Flash point Non-flammable
Related compounds
Other cations
Magnesium carbonate
Calcium carbonate
Barium carbonate
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

Contents

Chemical propertiesEdit

PreparationEdit

Other than the natural occurrence as a mineral, strontium carbonate is prepared synthetically in one of two processes, both of which start with naturally occurring celestine, a mineral form of strontium sulfate (SrSO4). In the "black ash" process, celesite is roasted with coke at 110–1300 °C to form strontium sulfide.[1] The sulfate is reduced, leaving the sulfide:

SrSO4 + 2 C → SrS + 2 CO2

A mixture of strontium sulfide with either carbon dioxide gas or sodium carbonate then leads to formation of a precipitate of strontium carbonate.[2][1]

SrS + H2O + CO2 → SrCO3 + H2S
SrS + Na2CO3 → SrCO3 + Na2S

In the "direct conversion" or double-decomposition method, a mixture of celesite and sodium carbonate is treated with steam to form strontium carbonate with substantial amounts of undissolved other solids.[1] This material is mixed with hydrochloric acid, which dissolves the strontium carbonate to form a solution of strontium chloride. Carbon dioxide or sodium carbonate is then used to re-precipitate strontium carbonate, as in the black-ash process.

UsesEdit

 
Nitric acid reacts with strontium carbonate to form strontium nitrate.

It is widely used in the ceramics industry as an ingredient in glazes. It acts as a flux and also modifies the color of certain metallic oxides. It has some properties similar to barium carbonate.


Strontium carbonate is also used for making some superconductors such as BSCCO and also for electroluminescent materials where it is first calcined into SrO and then mixed with sulphur to make SrS:x where x is typically europium[citation needed]. This is the famous "blue/green" phosphor which is sensitive to frequency and changes from lime green to blue[citation needed]. Other dopants can also be used such as gallium, or yttrium to get a yellow/orange glow instead.

Because of its status as a weak Lewis base, strontium carbonate can be used to produce many different strontium compounds by simple use of the corresponding acid.

Microbial precipitationEdit

The cyanobacteria, Calothrix, Synechococcus and Gloeocapsa, can precipitate strontian calcite in groundwater. The strontium exists as strontianite in solid solution within the host calcite with the strontium content of up to one percent.[3]

ReferencesEdit

  1. ^ a b c Aydoğan, Salih; Erdemoğlu, Murat; Aras, Ali; Uçar, Gökhan; Özkan, Alper (2006). "Dissolution kinetics of celestite (SrSO4) in HCl solution with BaCl2". Hydrometallurgy. 84 (3–4): 239–246. doi:10.1016/j.hydromet.2006.06.001.
  2. ^ MacMillan, J. Paul; Park, Jai Won; Gerstenberg, Rolf; Wagner, Heinz; Köhler, Karl; Wallbrecht, Peter. "Strontium and Strontium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_321.
  3. ^ Henry Lutz Ehrlich; Dianne K Newman (2009). Geomicrobiology, Fifth Edition. CRC Press. p. 177.

External linksEdit