Chemistry:Barium carbonate
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| Identifiers | |
3D model (JSmol)
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PubChem CID
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| UN number | 1564 |
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| Properties | |
| BaCO3 | |
| Molar mass | 197.34 g/mol |
| Appearance | white crystals |
| Odor | odorless |
| Density | 4.286 g/cm3 |
| Melting point | 811 °C (1,492 °F; 1,084 K) polymorphic transformation |
| Boiling point | 1,450 °C (2,640 °F; 1,720 K) decomposes[1] from 1360 °C |
| 16 mg/L (8.8°C) 22 mg/L (18 °C) 24 mg/L (20 °C) 24 mg/L (24.2 °C)[1] | |
Solubility product (Ksp)
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2.58·10−9 |
| Solubility | decomposes in acid insoluble in methanol |
| -58.9·10−6 cm3/mol | |
Refractive index (nD)
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1.676 |
| Structure | |
| orthorhombic | |
| Thermochemistry | |
Heat capacity (C)
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85.35 J/mol·K[1] |
Std molar
entropy (S |
112 J/mol·K[2] |
Std enthalpy of
formation (ΔfH⦵298) |
-1219 kJ/mol[2] |
Gibbs free energy (ΔfG˚)
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-1139 kJ/mol[1] |
| Hazards | |
| Safety data sheet | ICSC 0777 |
| GHS pictograms |  [3]
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| GHS Signal word | Warning |
| H302[3] | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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418 mg/kg, oral (rat) |
| Related compounds | |
Other cations
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Beryllium carbonate Magnesium carbonate Calcium carbonate Strontium carbonate Radium carbonate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
 verify (what is   ?)
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| Infobox references | |
Barium carbonate is the inorganic compound with the formula BaCO3. Like most alkaline earth metal carbonates, it is a white salt that is poorly soluble in water. It occurs as the mineral known as witherite. In a commercial sense, it is one of the most important barium compounds.[5]
Preparation
Barium carbonate is made commercially from barium sulfide by treatment with sodium carbonate at 60 to 70 °C (soda ash method) or, more commonly carbon dioxide at 40 to 90 °C:
In the soda ash process, an aqueous solution of barium sulfide is treated with sodium carbonate:[5]
- BaS + H2O + CO2 → BaCO3 + H2S
Reactions
Barium carbonate reacts with acids such as hydrochloric acid to form soluble barium salts, such as barium chloride:
- BaCO3 + 2 HCl → BaCl2 + CO2 + H2O
Pyrolysis of barium carbonate gives barium oxide.[6]
Uses
It is mainly used to remove sulfate impurities from feedstock of the chlor-alkali process. Otherwise it is a common precursor to barium-containing compounds such as ferrites.[5]
Other uses
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 it can leach from glazes into food and drink. To reduce toxicity concerns, it is often substituted with strontium carbonate, which behaves in a similar way in glazes but is of lower toxicity.
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
- ↑ 1.0 1.1 1.2 1.3 "Barium carbonate". http://chemister.ru/Database/properties-en.php?dbid=1&id=377.
- ↑ 2.0 2.1 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed.. Houghton Mifflin Company. ISBN 978-0-618-94690-7.
- ↑ 3.0 3.1 Sigma-Aldrich Co., Barium carbonate. Retrieved on 2014-05-06.
- ↑ Sciences labs MSDS
- ↑ 5.0 5.1 5.2 Kresse, Robert; Baudis, Ulrich; Jäger, Paul; Riechers, H. Hermann; Wagner, Heinz; Winkler, Jochen; Wolf, Hans Uwe (2007). "Ullmann's Encyclopedia of Industrial Chemistry". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a03_325.pub2.
- ↑ P. Ehrlich (1963). "Barium Oxide". in G. Brauer. Handbook of Preparative Inorganic Chemistry, 2nd Ed.. 1. NY, NY: Academic Press. pp. 933–944.
External links
| H2CO3 | He | ||||||||||||||||
| Li2CO3, LiHCO3 |
BeCO3 | B | C | (NH4)2CO3, NH4HCO3 |
O | F | Ne | ||||||||||
| Na2CO3, NaHCO3, Na3H(CO3)2 |
MgCO3, Mg(HCO3)2 |
Al2(CO3)3 | Si | P | S | Cl | Ar | ||||||||||
| K2CO3, KHCO3 |
CaCO3, Ca(HCO3)2 |
Sc | Ti | V | Cr | MnCO3 | FeCO3 | CoCO3 | NiCO3 | CuCO3 | ZnCO3 | Ga | Ge | As | Se | Br | Kr |
| Rb2CO3 | SrCO3 | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag2CO3 | CdCO3 | In | Sn | Sb | Te | I | Xe |
| Cs2CO3, CsHCO3 |
BaCO3 | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl2CO3 | PbCO3 | (BiO)2CO3 | Po | At | Rn | |
| Fr | Ra | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | |
| ↓ | |||||||||||||||||
| La2(CO3)3 | Ce2(CO3)3 | Pr | Nd | Pm | Sm | Eu | Gd | Tb | Dy | Ho | Er | Tm | Yb | Lu | |||
| Ac | Th | Pa | UO2CO3 | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No | Lr | |||
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