Wikipedia:WikiProject Chemicals/Chembox validation/VerifiedDataSandbox and Hexafluorosilicic acid: Difference between pages

{{short description|Octahedric silicon compound}}

{{redirect|Fluorosilicate|Fluorosilicate glass and glass-ceramics|Fluorosilicate glass}}

| Watchedfields = changed

| verifiedrevid = 477003251

| ImageFile=Hexafluorosilicic acid molecular structure.png

| ImageFile2=

| OtherNames = Fluorosilicic acid, fluosilic acid, hydrofluorosilicic acid, silicofluoride, silicofluoric acid, oxonium hexafluorosilanediuide, oxonium hexafluoridosilicate(2−)

|Section1={{Chembox Identifiers

| CASNo = 16961-83-4

| CASNo_Ref = {{cascite|correct|CAS}}

| UNII_Ref = {{fdacite|correct|FDA}}

| UNII = 53V4OQG6U1

| PubChem = 21863527

| ChemSpiderID = 17215660

| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}

| EINECS = 241-034-8

| UNNumber = 1778

| RTECS = VV8225000

| SMILES = [H+].[H+].F[Si-2](F)(F)(F)(F)F

| SMILES1 = [H+].[H+].F[Si--](F)(F)(F)(F)F

| StdInChI_Ref = {{stdinchicite|correct|chemspider}}

| InChI = 1/F6Si/c1-7(2,3,4,5)6/q-2/p+2

| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}

| InChIKey = OHORFAFFMDIQRR-SKRXCDHZAM}}

|Section2={{Chembox Properties

| H=2 | Si=1 | F=6

| Appearance = transparent, colorless, fuming liquid

| Odor = sour, pungent

| Density = 1.22 g/cm³ (25% soln.)<br/>1.38 g/cm³ (35% soln.)<br/>1.46 g/cm³ (61% soln.)

| MeltingPt = {{circa}} {{convert|19|C|F K}} {{nowrap|(60–70% solution)}}<br/><&nbsp;{{convert|-30|C|F K}} {{nowrap|(35% solution)}}

| BoilingPtC = 108.5

| BoilingPt_notes = (decomposes)

| Solubility = miscible

| RefractIndex = 1.3465

| pKa = 1.92<ref name=P82db>{{cite book|title=Ionisation Constants of Inorganic Acids and Bases in Aqueous Solution|editor-first=D.&nbsp;D.|editor-last=Perrin|edition=2^nd|series=[[IUPAC]] Chemical Data|issue=29|publisher=Pergamon|location=Oxford|year=1982|publication-date=1984|orig-date=1969|lccn=82-16524|isbn=0-08-029214-3|at=Entry 91}}</ref>

}}

|Section3={{Chembox Structure

| MolShape = Octahedral SiF₆^{2<nowiki>&minus;</nowiki>}

|Section7={{Chembox Hazards

| NFPA-H = 3

| NFPA-F = 0

| NFPA-R = 0

| NFPA-S =

| ExternalSDS = [https://www.sigmaaldrich.com/US/en/sds/sial/01302 External MSDS]

| GHSPictograms = {{GHS05}}

| GHSSignalWord = Danger

| HPhrases = {{H-phrases|314}}

| PPhrases = {{P-phrases|260|264|280|301+330+331|303+361+353|304+340|305+351+338|310|321|363|405|501}}

| FlashPt = Non-flammable

| LD50 = 430 mg/kg (oral, rat)

|Section8={{Chembox Related

| OtherCations = [[Ammonium hexafluorosilicate]]<br />

| OtherAnions = [[Hexafluorotitanic acid]]<br />[[Hexafluorozirconic acid]]

| OtherCompounds = [[Hexafluorophosphoric acid]]<br/>[[Fluoroboric acid]]}}

'''Hexafluorosilicic acid''' is an [[inorganic compound]] with the [[chemical formula]] {{Chem|H|2|SiF|6}}. Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.

Hexafluorosilicic acid is produced naturally on a large scale in volcanoes.<ref name="Palache">Palache, C., Berman, H., and Frondel, C. (1951) Dana’s System of Mineralogy, Volume II: Halides, Nitrates, Borates, Carbonates, Sulfates, Phosphates, Arsenates, Tungstates, Molybdates, etc. John Wiley and Sons, Inc., New York, 7th edition.</ref><ref name="Anthony">Anthony, J.W., Bideaux, R.A., Bladh, K.W., and Nichols, M.C. (1997) Handbook of Mineralogy, Volume III: Halides, Hydroxides, Oxides. Mineral Data Publishing, Tucson.

*[https://www.handbookofmineralogy.org/pdfs/bararite.pdf Bararite]

*[https://www.handbookofmineralogy.org/pdfs/cryptohalite.pdf Cryptohalite]</ref> It is manufactured as a coproduct in the production of [[phosphate fertilizer]]s. The resulting hexafluorosilicic acid is almost exclusively consumed as a precursor to [[aluminum trifluoride]] and [[sodium hexafluoroaluminate|synthetic cryolite]], which are used in aluminium processing. Salts derived from hexafluorosilicic acid are called '''hexafluorosilicates'''.

== Structure ==

[[File:(H5O2)2SiF6, ICSDcode 40388full.png|thumb|left|Structure of (H₅O₂)₂SiF₆. The hydrogen bonding between the fluoride and protons are indicated by dashed lines. Color code: green = F, orange = Si, red = O, gray = H.<ref name=ZAAC>{{cite journal |doi=10.1002/zaac.19885590103|title=The Crystalline Hydrates of Hexafluorosilicic Acid: A Combined Phase-Analytical and Structural Study |year=1988 |last1=Mootz |first1=D. |last2=Oellers |first2=E.-J. |journal=Zeitschrift für anorganische und allgemeine Chemie |volume=559 |pages=27–39 }}</ref>]]

Hexafluorosilicic acid has been crystallized as various hydrates. These include ([[Hydronium#Zundel cation|H₅O₂]])₂SiF₆, the more complicated (H₅O₂)₂SiF₆·2H₂O, and (H₅O₂)(H₇O₃)SiF₆·4.5H₂O. In all of these salts, the octahedral hexafluorosilicate anion is [[hydrogen bond]]ed to the cations.<ref name=ZAAC/>

Aqueous solutions of hexafluorosilicic acid are often described as {{Chem|H|2|SiF|6}}.

==Production and principal reactions==

Hexafluorosilicic acid is produced commercially from fluoride-containing minerals that also contain silicates. Specifically, [[apatite]] and [[fluorapatite]] are treated with [[sulfuric acid]] to give [[phosphoric acid]], a precursor to several water-soluble fertilizers. This is called the [[Phosphoric acid#Wet process|wet phosphoric acid process]].<ref name=USGS>USGS. [http://minerals.usgs.gov/minerals/pubs/commodity/fluorspar/fluormcs07.pdf Fluorspar].</ref> As a by-product, approximately 50&nbsp;kg of hexafluorosilicic acid is produced per tonne of HF owing to reactions involving silica-containing mineral impurities.<ref name= NTP >{{Cite web| url=https://ntp.niehs.nih.gov/ntp/htdocs/chem_background/exsumpdf/fluorosilicates_508.pdf| title=Sodium Hexafluorosilicate [CASRN 16893-85-9] and Fluorosilicic Acid [CASRN 16961-83-4] Review of Toxicological Literature| website=National Toxicology Program (U.S.)| access-date=13 July 2017| url-status=live| archive-url=https://web.archive.org/web/20121022152457/http://ntp.niehs.nih.gov/ntp/htdocs/Chem_Background/ExSumPDF/Fluorosilicates.pdf| archive-date=22 October 2012}}</ref>{{rp|p=3}}

Some of the [[hydrogen fluoride]] (HF) produced during this process in turn reacts with [[silicon dioxide]] (SiO₂) impurities, which are unavoidable constituents of the mineral feedstock, to give [[silicon tetrafluoride]]. Thus formed, the silicon tetrafluoride reacts further with HF.{{Citation needed|date=May 2022}} The net process can be described as:<ref name="Ullmann">{{Ullmann|first1=J. |last1=Aigueperse |first2=P. |last2=Mollard |first3=D. |last3=Devilliers |first4=M. |last4=Chemla |first5=R. |last5=Faron |first6=R. |last6=Romano |first7=J. P. |last7=Cuer |title=Fluorine Compounds, Inorganic |date=2005 |doi=10.1002/14356007.a11_307}}</ref>{{Page needed|date=May 2022}}

:{{chem2|6 HF + SiO2 → SiF6(2-) + 2 H3O+}}

Hexafluorosilicic acid can also be produced by treating silicon tetrafluoride with hydrofluoric acid.<ref name=Ullmann/>

== Reactions ==

Hexafluorosilic acid is only stable in [[hydrogen fluoride]] or acidic aqueous solutions. In any other circumstance, it acts as a source of [[hydrofluoric acid]]. Thus, for example, hexafluorosilicic acid pure or in [[oleum]] solution evolves [[silicon tetrafluoride]] until the residual hydrogen fluoride re-establishes equilibrium:<ref name=Ullmann/>

:H₂SiF₆&nbsp;{{eqm}} 2&nbsp;HF(''l'') + SiF₄(''g'')

In alkaline-to-neutral aqueous solutions, hexafluorosilicic acid readily hydrolyzes to fluoride anions and amorphous, [[hydrated silica]] ("SiO₂"). Strong bases give fluorosilicate salts at first, but any [[limiting reactant|stoichiometric excess]] begins hydrolysis.<ref name=Ullmann/> At the concentrations usually used for water fluoridation, 99% hydrolysis occurs:{{r|NTP}}<ref>{{cite journal|last2=Wilson|first2=Erin|last3=Callender|first3=Andrew|last4=Morris|first4=Michael D.|last5=Beck|first5=Larry W.|year=2006|title=Reexamination of Hexafluorosilicate Hydrolysis by ¹⁹F NMR and pH Measurement|journal=Environ. Sci. Technol.|volume=40|issue=8|pages=2572–2577|doi=10.1021/es052295s|last1=Finney|first1=William F.|pmid=16683594|bibcode=2006EnST...40.2572F}}</ref>

:{{Chem|SiF|6|2-}} + 2&nbsp;H₂O → 6&nbsp;F⁻ + SiO₂ + 4&nbsp;H⁺

=== Alkali and alkaline earth salts ===

Neutralization of solutions of hexafluorosilicic acid with [[alkali metal]] bases produces the corresponding alkali metal fluorosilicate salts:

:H₂SiF₆ + 2&nbsp;NaOH → Na₂SiF₆ + 2 H₂O

The resulting salt Na₂SiF₆ is mainly used in water fluoridation. Related ammonium and barium salts are produced similarly for other applications. At room temperature 15-30% concentrated hexafluorosilicic acid undergoes similar reactions with [[chloride]]s, [[hydroxide]]s, and [[carbonate]]s of alkali and [[alkaline earth metal]]s.<ref name=":2">{{cite book |title=Silicon Tetrafluoride |year=1953 |series=Inorganic Syntheses |volume=4 |pages=147–8 |doi=10.1002/9780470132357.ch47 |vauthors=Hoffman CJ, Gutowsky HS, Schumb WC, Breck DW}}</ref>

[[Sodium hexafluorosilicate]] for instance may be produced by treating [[sodium chloride]] ({{chem2|NaCl}}) by hexafluorosilicic acid:{{r|NTP|p=3}}<ref name=":1">{{Cite patent|country=Us|number=A345458|title=Patent Silicon tetrafluoride generation|status=Granted|pubdate=January 3, 1982|gdate=1982|invent1=Keith|invent2=L. Yaws|inventor1-first=C. Hansen|inventor2-first=Carl}}</ref>{{rp|7}}

: {{chem2|2NaCl + H2SiF6}} {{overset|27 °C|→}} {{chem2|Na2SiF6↓ + 2 HCl}}

: {{chem2|BaCl2 + H2SiF6}} {{overset|27 °C|→}} {{chem2|BaSiF6↓ + 2 HCl}}

Heating [[sodium hexafluorosilicate]] gives [[silicon tetrafluoride]]:{{R|name=:1|page=8}}

:{{chem2|Na2SiF6}} {{overset|>400 °C|→}} {{chem2|SiF4 + 2 NaF}}

== Uses ==

The majority of the hexafluorosilicic acid is converted to [[aluminium fluoride]] and [[Sodium hexafluoroaluminate|synthetic cryolite]]. These materials are central to the conversion of aluminium ore into [[aluminium]] metal. The conversion to aluminium trifluoride is described as:<ref name=Ullmann/>

:H₂SiF₆ + Al₂O₃ → 2&nbsp;AlF₃ + SiO₂ + H₂O

Hexafluorosilicic acid is also converted to a variety of useful hexafluorosilicate salts. The potassium salt, [[Potassium fluorosilicate]], is used in the production of porcelains, the magnesium salt for hardened concretes and as an insecticide, and the barium salts for phosphors.

Hexafluorosilicic acid and the salts are used as [[wood preservation]] agents.<ref>{{cite journal |author=Carsten Mai, [[Holger Militz]] |year=2004 |title=Modification of wood with silicon compounds. inorganic silicon compounds and sol-gel systems: a review |journal=Wood Science and Technology |volume=37 |issue=5 |page=339 |doi=10.1007/s00226-003-0205-5 |s2cid=9672269}}</ref>

=== Lead refining ===

Hexafluorosilicic acid is also used as an electrolyte in the [[Betts electrolytic process]] for refining lead.

=== Rust removers ===

Hexafluorosilicic acid (identified as hydrofluorosilicic acid on the label) along with [[oxalic acid]] are the active ingredients used in ''Iron Out'' rust-removing cleaning products, which are essentially varieties of [[laundry sour]].

===Niche applications===

H₂SiF₆ is a specialized [[reagent]] in [[organic synthesis]] for cleaving Si–O bonds of [[silyl ether]]s. It is more reactive for this purpose than HF. It reacts faster with ''t''-[[Silyl ether|butyldimethysilyl]] ([[TBDMS]]) ethers than [[Silyl ether|triisopropylsilyl]] ([[Silyl ether|TIPS]]) ethers.<ref>{{cite book|last1=Pilcher|first1= A. S. |last2=DeShong |first2=P. |chapter=Fluorosilicic Acid |title=Encyclopedia of Reagents for Organic Synthesis|date=2001 |publisher=John Wiley & Sons |doi=10.1002/047084289X.rf013|isbn= 0471936235 }}</ref>

=== Treating concrete ===

The application of hexafluorosilica acid to a calcium rich surface such as concrete will give that surface some resistance to acid attack.<ref>Properties of Concrete by A M Neville</ref>

:CaCO₃ + H₂O → &nbsp;Ca²⁺ + 2&nbsp;OH⁻ + CO₂

:H₂SiF₆ → 2&nbsp;H⁺ + {{Chem|SiF|6|2-}}

:{{Chem|SiF|6|2-}} + 2&nbsp;H₂O → 6&nbsp;F⁻ + SiO₂ + 4&nbsp;H⁺

: &nbsp;Ca²⁺ + 2&nbsp;F⁻ → CaF₂

[[Calcium fluoride]] (CaF₂) is an insoluble solid that is acid resistant.

==Natural salts==

Some rare minerals, encountered either within volcanic or coal-fire fumaroles, are salts of the hexafluorosilicic acid. Examples include [[Ammonium fluorosilicate|ammonium hexafluorosilicate]] that naturally occurs as two polymorphs: cryptohalite and [[bararite]].<ref name="Cryptohalite, Mindat">{{Cite web|url=https://www.mindat.org/min-1163.html|title=Cryptohalite}}</ref><ref name="Bararite, Mindat">{{Cite web|url=https://www.mindat.org/min-511.html|title = Bararite}}</ref><ref name="Kuszewski et al. 2020">{{cite journal | doi=10.1016/j.scitotenv.2019.134274 | title=Carbon‑nitrogen compounds, alcohols, mercaptans, monoterpenes, acetates, aldehydes, ketones, SF6, PH3, and other fire gases in coal-mining waste heaps of Upper Silesian Coal Basin (Poland) – a re-investigation by means of in situ FTIR external database approach | year=2020 | last1=Kruszewski | first1=Łukasz | last2=Fabiańska | first2=Monika J. | last3=Segit | first3=Tomasz | last4=Kusy | first4=Danuta | last5=Motyliński | first5=Rafał | last6=Ciesielczuk | first6=Justyna | last7=Deput | first7=Ewa | journal=Science of the Total Environment | volume=698 | page=134274 | pmid=31509784 | bibcode=2020ScTEn.698m4274K | s2cid=202563638 }}</ref>

==Safety==

Hexafluorosilicic acid can release [[hydrogen fluoride]] (HF) when evaporated, so it has similar risks. Inhalation of the vapors may cause [[lung edema]]. Like hydrogen fluoride, it attacks glass and [[stoneware]].<ref>{{cite web |url=http://niosh.dnacih.com/nioshdbs/ipcsneng/neng1233.htm |title=Fluorosilicic Acid – International Chemical Safety Cards |website=NIOSH |access-date=2015-03-10}}</ref> The [[Median lethal dose|LD₅₀]] value of hexafluorosilicic acid is 430&nbsp;mg/kg.{{r|NTP}}

==See also==

* [[Ammonium fluorosilicate]]

* [[Sodium fluorosilicate]]

* [[Potassium fluorosilicate]]

==References==

{{Reflist|30em}}

{{Hydrogen compounds}}

[[Category:Mineral acids]]

[[Category:Hydrogen compounds]]

[[Category:Nonmetal halides]]

[[Category:Hexafluorosilicates]]

[[Category:Fluoro complexes]]