Aluminium fluoride

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Aluminium fluoride
Aluminium-trifluoride-3D-polyhedra.png
Crystal structure
Names
Other names
Aluminium(III) fluoride
Aluminum trifluoride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.029.137
RTECS number BD0725000
Properties
AlF3
Molar mass 83.977 g/mol (anhydrous)
101.992 g/mol (monohydrate)
138.023 (trihydrate)[1]
Appearance white, crystalline solid
odorless
Density 3.10 g/cm3 (anhydrous)
2.17 g/cm3 (monohydrate)
1.914 g/cm3 (trihydrate)[1]
Melting point 1,290 °C (2,350 °F; 1,560 K)[4] (anhydrous) (sublimes)
5.6 g/L (0 °C)
6.7 g/L (20 °C)
17.2 g/L (100 °C)
-13.4·10−6 cm3/mol[2]
1.3767 (visible range)[3]
Structure
Rhombohedral, hR24
R3c, No. 167[5]
a = 0.49254 nm, c = 1.24477 nm
0.261519
6
Thermochemistry
75.1 J/mol·K[6]
66.5 J/mol·K[6]
−1510.4 kJ/mol[6]
-1431.1 kJ/mol[6]
Hazards[7][8][9]
Safety data sheet InChem MSDS
GHS pictograms CorrosiveAcute toxicityIrritantReproductive toxicity, target organ toxicity, aspiration hazard
GHS signal word DANGER
H301, H302, H314, H315, H319, H335, H361, H372
P260, P261, P264, P270, P271, P280, P301+310, P301+312, P301+330+331, P302+352, P303+361+353, P304+340, P305+351+338, P310, P312, P321, P330, P332+313, P337+313, P362, P363, P403+233, P405, P501
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gasReactivity 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
3
0
US health exposure limits (NIOSH):
PEL (Permissible)
none
REL (Recommended)
2 mg/m3
IDLH (Immediate danger)
N.D.
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☑Y verify (what is ☑Y‹See TfM›☒N ?)
Infobox references

Aluminium fluoride (AlF3) is an inorganic compound used primarily in the production of aluminium. This colorless solid can be prepared synthetically but also occurs in nature as minerals rosenbergite and oskarssonite.

Production and occurrence[edit]

The majority of aluminium fluoride is produced by treating alumina with hydrogen fluoride gas at 700 °C:[4] Fluorosilicic acid may also be used make aluminum fluoride.[10]

H2SiF6 + Al2O3 .3 H2O → 2 AlF3 + 3 SiO2 + 4 H2O

Alternatively, it is manufactured by thermal decomposition of ammonium hexafluoroaluminate.[11] For small scale laboratory preparations, AlF3 can also be prepared by treating aluminium hydroxide or aluminium metal with hydrogen fluoride.

Aluminium fluoride trihydrate is found in nature as the rare mineral rosenbergite. The non-hydrated form appears as the mineral oskarssonite.[12]

Structure[edit]

Its structure adopts the rhenium trioxide motif, featuring distorted AlF6 octahedra. Each fluoride is connected to two Al centers. Because of its 3-dimensional polymeric structure, AlF3 has a high melting point. The other trihalides of aluminium in the solid state differ, AlCl3 has a layer structure and AlBr3 and AlI3, are molecular dimers.[13] Also they have low melting points and evaporate readily to give dimers.[14] In the gas phase aluminium fluoride exists as trigonal molecules of D3h symmetry. The Al-F bond lengths of this gaseous molecule are 163 pm.

Like most gaseous metal trifluorides, AlF3 adopts a planar structure upon evaporation.

Applications[edit]

Aluminium fluoride is an important additive for the production of aluminium by electrolysis.[4] Together with cryolite, it lowers the melting point to below 1000 °C and increases the conductivity of the solution. It is into this molten salt that aluminium oxide is dissolved and then electrolyzed to give bulk Al metal.[11]

Aluminum fluoride complexes are used to study the mechanistic aspects of phosphoryl transfer reactions in biology, which are of fundamental importance to cells, as phosphoric acid anhydrides such as ATP and GTP control most of the reactions involved in metabolism, growth and differentiation.[15] The observation that aluminum fluoride can bind to and activate heterotrimeric G proteins has proven to be useful for the study of G protein activation in vivo, for the elucidation of three-dimensional structures of several GTPases, and for understanding the biochemical mechanism of GTP hydrolysis, including the role of GTPase-activating proteins.[16]

Niche uses[edit]

Together with zirconium fluoride, aluminium fluoride is an ingredient for the production of fluoroaluminate glasses.

It is also used to inhibit fermentation.

Like magnesium fluoride it used as a low-index optical thin film, particularly when far UV transparency is required. Its deposition by physical vapor deposition, particularly by evaporation, is favorable.

Safety[edit]

Aluminum fluoride reported oral animal lethal dose (LD50) is 0.1 g/kg.[17] Repeated or prolonged inhalation exposure may cause asthma, and may have effects on the bone and nervous system, resulting in bone alterations (fluorosis), and nervous system impairment.[18]

Many of the neurotoxic effects of fluoride are due to the formation of aluminum fluoride complexes, which mimic the chemical structure of a phosphate and influence the activity of ATP phosphohydrolases and phospholipase D. Only micromolar concentrations of aluminum are needed to form aluminum fluoride.[19]

Human exposure to aluminum fluoride can occur in an industrial setting, such as emissions from an aluminum reduction processes,[20] or when a person ingests both a fluoride source (e.g., fluoride in drinking water or residue of fluoride-based pesticides) and an aluminum source; sources of human exposure to aluminum include drinking water, tea, food residues, infant formula, aluminum-containing antacids or medications, deodorants, cosmetics, and glassware.[19] Fluoridation chemicals may also contain aluminum fluoride.[21] Data on the potential neurotoxic effects of chronic exposure to the aluminum species existing in water is limited.[22]

References[edit]

  1. ^ a b Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.45. ISBN 1439855110.
  2. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.131. ISBN 1439855110.
  3. ^ Lide, David R. (2003-06-19). CRC Handbook of Chemistry and Physics, 84th Edition. CRC Handbook. CRC Press. ISBN 9780849304842.
  4. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 233. ISBN 0-08-037941-9.
  5. ^ Hoppe, R.; Kissel, D. (1984). "Zur kenntnis von AlF3 und InF3 [1]". Journal of Fluorine Chemistry. 24 (3): 327. doi:10.1016/S0022-1139(00)81321-4.
  6. ^ a b c d Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 5.5. ISBN 1439855110.
  7. ^ Pohanish, Richard P. (2005-03-04). HazMat Data: For First Response, Transportation, Storage, and Security. John Wiley & Sons. ISBN 9780471726104.
  8. ^ "Aluminum Fluoride". PubChem. National Institute of Health. Retrieved October 12, 2017.
  9. ^ "NIOSH Pocket Guide to Chemical Hazards #0024". National Institute for Occupational Safety and Health (NIOSH).
  10. ^ Dreveton, Alain (2012-01-01). "Manufacture of Aluminium Fluoride of High Density and Anhydrous Hydrofluoric Acid from Fluosilicic Acid". Procedia Engineering. SYMPHOS 2011 - 1st International Symposium on Innovation andTechnology in the Phosphate Industry. 46 (Supplement C): 255–265. doi:10.1016/j.proeng.2012.09.471.
  11. ^ a b Aigueperse, J.; Mollard, P.; Devilliers, D.; Chemla, M.; Faron, R.; Romano, R.; Cuer, J. P. (2005) "Fluorine Compounds, Inorganic" in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim. doi:10.1002/14356007.a11_307
  12. ^ Oskarssonite. Mindat.
  13. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
  14. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  15. ^ Wittinghofer, Alfred (1997-11-01). "Signaling mechanistics: Aluminum fluoride for molecule of the year". Current Biology. 7 (11): R682–R685. doi:10.1016/S0960-9822(06)00355-1.
  16. ^ Vincent, Sylvie; Brouns, Madeleine; Hart, Matthew J.; Settleman, Jeffrey (1998-03-03). "Evidence for distinct mechanisms of transition state stabilization of GTPases by fluoride". Proceedings of the National Academy of Sciences. 95 (5): 2210–2215. doi:10.1073/pnas.95.5.2210. ISSN 0027-8424. PMC 19296. PMID 9482864.
  17. ^ "ALUMINUM FLUORIDE, CASRN: 7784-18-1". National Library of Medicine HSDB Database. CDC.gov. June 24, 2005. Retrieved October 12, 2017.
  18. ^ "ALUMINIUM FLUORIDE (ANHYDROUS) International Chemical Safety Cards (ICSC)". CDC.gov National Institute for Occupational Safety and Health (NIOSH). July 22, 2015. Retrieved July 17, 2017.
  19. ^ a b Fluoride in Drinking Water: A Scientific Review of EPA's Standards. https://www.nap.edu/read/11571: The National Academies Press. 2006. pp. 51–52, 219. doi:10.17226/11571.
  20. ^ TOXICOLOGICAL PROFILE FOR FLUORIDES, HYDROGEN FLUORIDE, AND FLUORINE. https://www.atsdr.cdc.gov/toxprofiles/tp11.pdf: U.S. DEPARTMENT OF HEALTH AND HUMAN SERVICES Public Health Service Agency for Toxic Substances and Disease Registry. 2003. p. 211.
  21. ^ Mullenix, Phyllis J (2014). "A new perspective on metals and other contaminants in fluoridation chemicals". International Journal of Occupational and Environmental health. 20 (2): 157–166. doi:10.1179/2049396714Y.0000000062. ISSN 1077-3525. PMC 4090869. PMID 24999851.
  22. ^ Aluminum Compounds Review of Toxicological Literature Abridged Final Report. Prepared for National Institute of Environmental Health Sciences. NTP.gov Nomination Summary for Aluminum contaminants of drinking water (N20025). October 2001

External links[edit]