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Aluminium fluoride

(Redirected from Aluminium trifluoride)

Aluminium fluoride is an inorganic compound with the formula AlF3. It forms hydrates AlF3·xH2O. Anhydrous AlF3 and its hydrates are all colorless solids. Anhydrous AlF3 is used in the production of aluminium. Several occur as minerals.

Aluminium fluoride

Anhydrous AlF3
Names
Other names
Aluminium(III) fluoride
Aluminum trifluoride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.029.137 Edit this at Wikidata
RTECS number
  • BD0725000
UNII
  • InChI=1S/Al.3FH/h;3*1H/q+3;;;/p-3 checkY
    Key: KLZUFWVZNOTSEM-UHFFFAOYSA-K checkY
  • InChI=1/Al.3FH/h;3*1H/q+3;;;/p-3
    Key: KLZUFWVZNOTSEM-DFZHHIFOAC
  • monomer: F[Al](F)F
  • crystal form: F[Al](F[Al]0(F)(F)(F)F)(F[Al]1(F)(F)(F)F)(F[Al]2(F)(F)(F)F)(F[Al]3(F)(F)(F)F)F[Al](F[Al](F[Al]4(F)(F)(F)F)(F[Al]5(F)(F)(F)F)(F[Al]6(F)(F)(F)F)(F0)F)(F[Al](F[Al]7(F)(F)(F)F)(F[Al]8(F)(F)(F)F)(F1)(F4)F)(F[Al](F[Al]9(F)(F)(F)F)(F[Al]0(F)(F)(F)F)(F5)(F7)F)(F[Al](F[Al]1(F)(F)(F)F)(F2)(F8)(F9)F)F[Al](F3)(F6)(F0)(F1)F
Properties
AlF3
Molar mass
  • 83.977 g/mol (anhydrous)
  • 101.992 g/mol (monohydrate)
  • 138.023 (trihydrate)
[1]
Appearance Colorless to white crystalline solid
Odor 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]
GHS labelling:
Corrosive Acute toxicity Irritant Reproductive toxicity, target organ toxicity, aspiration hazard
Danger
H301, H302, H314, H315, H319, H335, H361, H372
P260, P261, P264, P270, P271, P280, P301+P310, P301+P312, P301+P330+P331, P302+P352, P303+P361+P353, P304+P340, P305+P351+P338, P310, P312, P321, P330, P332+P313, P337+P313, P362, P363, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
0
0
NIOSH (US health exposure limits):
PEL (Permissible)
 none
REL (Recommended)
 2 mg/m3
IDLH (Immediate danger)
 N.D.
Safety data sheet (SDS) InChem MSDS
Related compounds
Other anions
Other cations
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Occurrence and production

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Aside from anhydrous AlF3, several hydrates are known. With the formula AlF3·xH2O, these compounds include monohydrate (x = 1), two polymorphs of the trihydrate (x = 3), a hexahydrate (x = 6), and a nonahydrate (x = 9).[10]

The majority of aluminium fluoride is produced by treating alumina with hydrogen fluoride at 700 °C:[4] Hexafluorosilicic acid may also be used make aluminium fluoride.[11]

H2[SiF6] + Al2O3 + 3 H2O → 2 AlF3 + SiO2 + 4 H2O

Alternatively, it is manufactured by thermal decomposition of ammonium hexafluoroaluminate.[12] For small scale laboratory preparations, AlF3 can also be prepared by treating aluminium hydroxide or aluminium with hydrogen fluoride.

Aluminium fluoride trihydrate is found in nature as the rare mineral rosenbergite.

The anhydrous form appears as the relatively recently (as of 2020) recognized mineral óskarssonite.[13][14] A related, exceedingly rare mineral, is zharchikhite, Al(OH)2F.[15][14]

Structure

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According to X-ray crystallography, anhydrous AlF3 adopts the rhenium trioxide motif, featuring distorted AlF6 octahedra. Each fluoride is connected to two Al centers. Because of its three-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.[16][page needed] Also they have low melting points and evaporate readily to give dimers.[17][page needed] 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

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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.[12]

Aluminium 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 adenosine triphosphate and guanosine triphosphate control most of the reactions involved in metabolism, growth and differentiation.[18] The observation that aluminium 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.[19]

Niche uses

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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 is 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

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The reported oral animal lethal dose (LD50) of aluminium fluoride is 100 mg/kg.[20] 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.[21]

Many of the neurotoxic effects of fluoride are due to the formation of aluminium fluoride complexes, which mimic the chemical structure of a phosphate and influence the activity of ATP phosphohydrolases and phospholipase D. Only micromolar concentrations of aluminium are needed to form aluminium fluoride.[22]

Human exposure to aluminium fluoride can occur in an industrial setting, such as emissions from aluminium reduction processes,[23] or when a person ingests both a fluoride source (e.g., fluoride in drinking water or residue of fluoride-based pesticides) and an aluminium source; sources of human exposure to aluminium include drinking water, tea, food residues, infant formula, aluminium-containing antacids or medications, deodorants, cosmetics, and glassware.[22] Fluoridation chemicals may also contain aluminium fluoride.[24] Data on the potential neurotoxic effects of chronic exposure to the aluminium species existing in water are limited.[25]

See also

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References

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  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 1-4398-5511-0.
  2. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.131. ISBN 1-4398-5511-0.
  3. ^ Lide, David R. (2003-06-19). CRC Handbook of Chemistry and Physics. CRC Handbook (84th ed.). 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 978-0-08-037941-8.
  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 1-4398-5511-0.
  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. ^ Guangmei Wang; Anja-Verena Mudring (2016). "The missing Hydrate AlF3·6H2O [Al(H2O)6]F3: Ionothermal Synthesis, Crystal Structure and Characterization of Aluminum Fluoride Hexahydrate". Solid State Sciences. 61: 61. doi:10.1016/j.solidstatesciences.2016.09.007.
  11. ^ 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 and Technology in the Phosphate Industry. 46 (Supplement C): 255–265. doi:10.1016/j.proeng.2012.09.471.
  12. ^ a b Aigueperse, J.; Mollard, P.; Devilliers, D.; Chemla, M.; Faron, R.; Romano, R.; Cuer, J. P. (2005). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_307. ISBN 978-3527306732.
  13. ^ "Óskarssonite".
  14. ^ a b "List of Minerals". 21 March 2011.
  15. ^ "Zharchikhite".
  16. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  17. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego, CA: Academic Press. ISBN 0-12-352651-5..
  18. ^ 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. PMID 9382787. S2CID 17666164.
  19. ^ 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. Bibcode:1998PNAS...95.2210V. doi:10.1073/pnas.95.5.2210. ISSN 0027-8424. PMC 19296. PMID 9482864.
  20. ^ "ALUMINUM FLUORIDE, CASRN: 7784-18-1". National Library of Medicine HSDB Database. CDC.gov. June 24, 2005. Retrieved October 12, 2017.
  21. ^ "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.
  22. ^ a b Fluoride in Drinking Water: A Scientific Review of EPA's Standards. The National Academies Press. 2006. pp. 51–52, 219. doi:10.17226/11571. ISBN 978-0-309-10128-8.
  23. ^ TOXICOLOGICAL PROFILE FOR FLUORIDES, HYDROGEN FLUORIDE, AND FLUORINE (PDF). U.S. DEPARTMENT OF HEALTH AND HUMAN SERVICES Public Health Service Agency for Toxic Substances and Disease Registry. 2003. p. 211.
  24. ^ 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.
  25. ^ 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
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