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Borax crystals
Systematic name Sodium tetraborate


Molecular formula Na2B4O7•10H2O
Molar mass 381.37 g/mol
Appearance white solid
CAS number 1303-96-4
Density and phase 1.73 g/cm³, solid
Solubility in water 5.1 g/100 ml (20 °C)
Melting point 75 °C
Boiling point 320 °C
Basicity (pKb) see text
Crystal structure Monoclinic
Thermodynamic data
Std enthalpy of
-3276.75 kJ/mol
Standard molar
189.53 J·K−1·mol−1
MSDS External MSDS
EU classification not listed
NFPA 704

NFPA 704.svg

Flash point non-flammable
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Related compounds
Other anions Sodium aluminate

Sodium gallate

Other cations Potassium tetraborate
Related compounds Boric acid

Sodium perborate

Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)

Borax, also called sodium borate, sodium tetraborate, or disodium tetraborate, is an important boron compound, a mineral, and a salt of boric acid. It is usually a white powder consisting of soft colorless crystals that dissolve easily in water.

Borax has a wide variety of uses. It is a component of many detergents, cosmetics, and enamel glazes. It is also used to make buffer solutions in biochemistry, as a fire retardant, an anti-fungal compound for fiberglass insulation, an insecticide, a flux in metallurgy, and a precursor for other boron compounds.

The term borax is used for a number of closely related minerals or chemical compounds that differ in their crystal water content, but usually refers to the decahydrate. Commercially sold borax is usually partially dehydrated.



The origin of the name is traceable to the Medieval Latin borax, which comes from the Arabic buraq, which comes from either the Persian burah[1] or the Middle Persian burak.[2]


Borax "cottonball"

Borax occurs naturally in evaporite deposits produced by the repeated evaporation of seasonal lakes. The most commercially important deposits are found in Turkey, Tibet, the Atacama desert in Chile, and near Boron, California, and other locations in the Southwestern United States. Borax can also be produced synthetically from other boron compounds.

Notable characteristics

The structure of the anion [B4O5(OH)4]2− in borax

The term borax is often used for a number of closely related minerals or chemical compounds that differ in their crystal water content:

  • Anhydrous borax (Na2B4O7)
  • Borax pentahydrate (Na2B4O7•5H2O)
  • Borax decahydrate (Na2B4O7•10H2O)

Borax is generally described as Na2B4O7•10H2O. However, it is better formulated as Na2[B4O5(OH)4]•8H2O, since borax contains the [B4O5(OH)4]2− ion. In this structure, there are two four-coordinate boron atoms (two BO4 tetrahedra) and two three-coordinate boron atoms (two BO3 triangles).

Borax is also easily converted to boric acid and other borates, which have many applications. If left exposed to dry air, it slowly loses its water of hydration and becomes the white and chalky mineral tincalconite (Na2B4O7•5H2O).

When borax is burned, it produces a bright orange-colored flame. Because of this, it is sometimes used for homemade pyrotechnics.



Sodium borate is used in biochemical and chemical laboratories to make buffer solutions, e.g. for gel electrophoresis of DNA. It has a lower conductivity, produces sharper bands, and can be run at higher speeds than can gels made from TBE Buffer or TAE Buffer (five to 35 V/cm as compared to five to ten V/cm). At a given voltage, the heat generation and thus the gel temperature is much lower than with TBE or TAE buffers, therefore the voltage can be increased to speed up electrophoresis so that a gel run takes only a fraction of the usual time. Downstream applications, such as isolation of DNA from a gel slice or Southern blot analysis, work as expected with sodium borate gels. Borate buffers (usually at pH 8) are also used as preferential equilibration solution in DMP-based crosslinking reactions.

Lithium borate is similar to sodium borate and has all of its advantages, but permits use of even higher voltages due to the lower conductivity of lithium ions as compared to sodium ions.[3] However, lithium borate is much more expensive.


A mixture of borax and ammonium chloride is used as a flux when welding iron and steel. It lowers the melting point of the unwanted iron oxide (scale), allowing it to run off. Borax is also used mixed with water as a flux when soldering jewelry metals such as gold or silver. It allows the molten solder to flow evenly over the joint in question. Borax is also a good flux for 'pre-tinning' tungsten with zinc, making the tungsten soft-solderable.[4]

Food additive

Borax is used as a food additive in some countries with the E number E285, but is banned in the United States. Its use is similar to salt, and it appears in French and Iranian caviar.

Other uses

  • component of detergents
  • component of cosmetics
  • ingredient in enamel glazes
  • component of glass, pottery, and ceramics
  • fire retardant
  • anti-fungal compound for fiberglass and cellulose insulation
  • component of Slime
  • insecticide to kill ants and fleas
  • precursor for sodium perborate monohydrate that is used in detergents, as well as for boric acid and other borates
  • treatment for thrush in horse's hoofs
  • used to make indelible ink for dip pens by dissolving shellac into heated borax


Boric acid, sodium borate, and sodium perborate are estimated to have a fatal dose from 0.1 to 0.5g/kg.[5] These substances are toxic to all cells, and have a slow excretion rate through the kidneys. Kidney toxicity is the greatest, with liver fatty degeneration, cerebral edema, and gastroenteritis. Boric acid solutions used as an eye wash or on abraded skin are known to be especially toxic to infants, especially after repeated use due to its slow elimination rate.[6]

See also


  1. Retrieved September 3, 2007.
  2. Retrieved September 3, 2007.
  3. Digital object identifier (DOI): 10.1016/j.ab.2004.05.054 Analytical Biochemistry 2004; 333: 1-13. Retrieved September 3, 2007.
  4. Digital object identifier (DOI): 10.1119/1.1972398 Am. J. Phys. 34, xvi, 1966. Retrieved September 3, 2007.
  5. Handbook of Poisoning. Robert H. Dreisback. eighth edition, p.314.
  6. Goodman and Gillman's: The Pharmacological Basis of Therapeutics, 6th edition


  • Farndon, John. The Practical Encyclopedia of Rocks & Minerals: How to Find, Identify, Collect and Maintain the World's best Specimens, with over 1000 Photographs and Artworks. London: Lorenz Books, 2006. ISBN 0754815412
  • Klein, Cornelis, and Barbara Dutrow. Manual of Mineral Science. 23rd ed. New York: John Wiley, 2007. ISBN 978-0471721574
  • Pellant, Chris. Rocks and Minerals. Smithsonian Handbooks. New York: Dorling Kindersley, 2002. ISBN 0789491060
  • Shaffer, Paul R., Herbert S. Zim, and Raymond Perlman. Rocks, Gems and Minerals. Rev. ed. New York: St. Martin's Press, 2001. ISBN 1582381321
  • Spears, John Randolph, and Douglas W. Steeples. Illustrated Sketches of Death Valley and Other Borax Deserts of the Pacific Coast. American Land Classics. Baltimore: Johns Hopkins University Press, 2001. ISBN 0801865077
  • Travis, Norman J., and E. J. Cocks. The Tincal Trail: A History of Borax. London: Harrap, 1984. ISBN 0245537988

External links

All links retrieved February 18, 2013.


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