Barium carbonate
| Barium carbonate | |
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Other names
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| Identifiers | |
| CAS number | 513-77-9  |
| ChemSpider | 10121 |
| UNII | 6P669D8HQ8 |
| EC number | 208-167-3 |
| UN number | 1564 |
| RTECS number | CQ8600000 |
| Jmol-3D images | Image 1 Image 2 |
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| Properties | |
| Molecular formula | BaCO3 |
| Molar mass | 197.34 g/mol |
| Appearance | white crystals |
| Odor | odorless |
| Density | 4.286 g/cm3 |
| Melting point |
811 °C |
| Boiling point |
1360 °C (decomp) |
| Solubility in water | 0.0024 g/100 mL (20 °C) |
| Solubility | decomposes in acid |
| Refractive index (nD) | 1.60 |
| Thermochemistry | |
| Std enthalpy of formation ΔfH |
−1219 kJ·mol−1[1] |
| Standard molar entropy S |
112 J·mol−1·K−1[2] |
| Hazards | |
| MSDS | ICSC 0777 |
| EU Index | 056-003-00-2 |
| EU classification | Harmful (Xn) |
| R-phrases | R22 |
| S-phrases | (S2), S24/25 |
| NFPA 704 |
 0
2
1
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| Flash point | Non-flammable |
| LD50 | 418 mg/kg, oral (rat) |
| Related compounds | |
| Other cations | Magnesium carbonate Calcium carbonate Strontium carbonate |
 (verify) (what is:  / ?)Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
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| Infobox references | |
Barium carbonate (BaCO3), also known as witherite, is a chemical compound used in rat poison, bricks, ceramic glazes and cement.
Witherite crystallizes in the orthorhombic system. The crystals are invariably twinned together in groups of three, giving rise to pseudo-hexagonal forms somewhat resembling bipyramidal crystals of quartz, the faces are usually rough and striated horizontally.
The mineral is named after William Withering, who in 1784 recognized it to be chemically distinct from barytes.[3] It occurs in veins of lead ore at Hexham in Northumberland, Alston in Cumbria, Anglezarke, near Chorley in Lancashire and a few other localities. Witherite is readily altered to barium sulfate by the action of water containing calcium sulfate in solution and crystals are therefore frequently encrusted with barytes. It is the chief source of barium salts and is mined in considerable amounts in Northumberland. It is used for the preparation of rat poison, in the manufacture of glass and porcelain, and formerly for refining sugar. It is also used for controlling the chromate to sulfate ratio in chromium electroplating baths.[4]
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Preparation [edit]
Barium carbonate is made commercially from barium sulfide either by treatment with sodium carbonate at 60 to 70 °C (soda ash method) or by passing carbon dioxide at 40 to 90 °C.
In the soda ash process, solid or dissolved sodium carbonate is added to barium sulfide solution, and the barium carbonate precipitate is filtered, washed and dried.[5]
Reactions [edit]
Barium carbonate reacts with acids such as hydrochloric acid to form soluble barium salts, such as barium chloride:
- BaCO3 (s) + 2 HCl (aq) → BaCl2 (aq) + CO2 (g) + H2O (l)
However, the reaction with sulfuric acid is poor, because barium sulfate is highly insoluble.
Uses [edit]
Barium carbonate is widely used in the ceramics industry as an ingredient in glazes. It acts as a flux, a matting and crystallizing agent and combines with certain colouring oxides to produce unique colours not easily attainable by other means. Its use is somewhat controversial since some claim that it can leach from glazes into food and drink. To provide a safe means of use, BaO is often used in fritted form.
In the brick, tile, earthenware and pottery industries barium carbonate is added to clays to precipitate soluble salts (calcium sulfate and magnesium sulfate) that cause efflorescence.
References [edit]
- ^ Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 0-618-94690-X.
- ^ Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 0-618-94690-X.
- ^ Withering, William (1784). "Experiments and Observations on Terra Poderosa". Philosophical Transactions of the Royal Society of London 74: 293–311.
- ^ Whitelaw, G.P. (2003-10-25). "Standard Chrome Bath Control". finishing.com. Archived from the original on 13 December 2006. Retrieved 2006-11-29.
- ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
This article incorporates text from a publication now in the public domain: Chisholm, Hugh, ed. (1911). Encyclopædia Britannica (11th ed.). Cambridge University Press.
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