Hydrofluoric acid

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Hydrofluoric acid
Ball-and-stick model of hydrogen fluoride
Ball-and-stick model of water
Ball-and-stick model of the fluoride anion
Ball-and-stick model of the hydronium cation
White plastic bottle With safety cap, labeled "QP Panreac" above smaller text "Hydrofluoric Acid 40% QP" with 6 translations. In a bright orange region along the side, warning symbols are visible.
Names
IUPAC name
Fluorane[1]
Other names
Fluorhydric acid
Hydronium fluoride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number 231-634-8
RTECS number MW7875000
UNII
UN number 1790
Properties
HF (aq)
Molar mass not applicable
(see hydrogen fluoride)
Appearance Colorless solution
Density 1.15 g/mL (for 48% soln.)
Melting point Not applicable
(see hydrogen fluoride)
Boiling point Not applicable
(see hydrogen fluoride)
Miscible.
Acidity (pKa) 3.17[2]
Hazards[3]
Safety data sheet Seastar Chemicals MSDS
GHS pictograms CorrosiveAcute Toxicity
GHS signal word DANGER
H280, H300, H310, H314, H318, H330
P260, P262, P264, P270, P271, P280, P284, P301+310, P301+330+331, P302+350, P303+361+353, P304+340, P305+351+338, P310, P320, P321, P322, P330, P361, P363, P403+233, P405, P410+403, P501
NFPA 704
Flammability code 0: Will not burn. E.g. waterHealth code 4: Very short exposure could cause death or major residual injury. E.g. VX 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
4
0
Flash point Non-flammable
Related compounds
Other anions
Hydrochloric acid
Hydrobromic acid
Hydroiodic acid
Related compounds
Hydrogen fluoride
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☒N ?)
Infobox references

Hydrofluoric acid is a solution of hydrogen fluoride (HF) in water. It is a precursor to almost all fluorine compounds, including pharmaceuticals such as fluoxetine (Prozac), diverse materials such as PTFE (Teflon), and elemental fluorine itself. It is a colourless solution that is highly corrosive, capable of dissolving many materials, especially oxides. Its ability to dissolve glass has been known since the 17th century, even before Carl Wilhelm Scheele prepared it in large quantities in 1771.[4] Because of its high reactivity toward glass and moderate reactivity toward many metals, hydrofluoric acid is usually stored in plastic containers (although PTFE is slightly permeable to it).[5]

Hydrogen fluoride gas is an acute poison that may immediately and permanently damage lungs and the corneas of the eyes. Aqueous hydrofluoric acid is a contact-poison with the potential for deep, initially painless burns and ensuing tissue death. By interfering with body calcium metabolism, the concentrated acid may also cause systemic toxicity and eventual cardiac arrest and fatality.

Acidity[edit]

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution.[6] This is in part a result of the strength of the hydrogen–fluorine bond, but also of other factors such as the tendency of HF, H
2
O
, and F
anions to form clusters.[7] At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF
2
) and protons, thus greatly increasing the acidity.[8] This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions.[9] Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated.[8]

The acidity of hydrofluoric acid solutions varies with concentration owing to hydrogen-bond interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization constant Ka = 6.6×104 (or pKa = 3.18),[10] in contrast to corresponding solutions of the other hydrogen halides, which are strong acids (pKa < 0). Concentrated solutions of hydrogen fluoride are much more strongly acidic than implied by this value, as shown by measurements of the Hammett acidity function H0[11](or "effective pH"). The H0 for 100% HF is estimated to be between −10.2 and −11, comparable to the value −12 for sulfuric acid.[12][13]

In thermodynamic terms, HF solutions are highly non-ideal, with the activity of HF increasing much more rapidly than its concentration. The weak acidity in dilute solution is sometimes attributed to the high H—F bond strength, which combines with the high dissolution enthalpy of HF to outweigh the more negative enthalpy of hydration of the fluoride ion.[14] However, Paul Giguère and Sylvia Turrell[15][16] have shown by infrared spectroscopy that the predominant solute species is the hydrogen-bonded ion pair [H
3
O+
·F], which suggests that the ionization can be described as a pair of successive equilibria:

H
2
O
+ HF
[H
3
O+
·F]

 

 

 

 

(1)

[H
3
O+
·F]
H
3
O+
+ F

 

 

 

 

(2)

The first equilibrium lies well to the right (K ≫ 1) and the second to the left (K ≪ 1), meaning that HF is extensively dissociated, but that the tight ion pairs reduce the thermodynamic activity coefficient of H3O+, so that the solution is effectively less acidic.[17]

In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion.[15][17]

[H
3
O+
⋅F] + HF ⇌ H
3
O+
+ HF
2

The increase in free H3O+ due to this reaction accounts for the rapid increase in acidity, while fluoride ions are stabilized (and become less basic) by strong hydrogen bonding to HF to form HF
2
. This interaction between the acid and its own conjugate base is an example of homoassociation (homoconjugation).

Production[edit]

Hydrofluoric acid is produced by treatment of the mineral fluorite (CaF2) with concentrated sulfuric acid. When combined at 265 °C, these two substances react to produce hydrogen fluoride and calcium sulfate according to the following chemical equation:

CaF2 + H2SO4 → 2 HF + CaSO4

Although bulk fluorite is a suitable precursor and a major source of world HF production, HF is also produced as a by-product of the production of phosphoric acid, which is derived from the mineral apatite. Apatite sources typically contain a few percent of fluoroapatite, acid digestion of which releases a gaseous stream consisting of sulfur dioxide (from the H2SO4), water, and HF, as well as particulates. After separation from the solids, the gases are treated with sulfuric acid and oleum to afford anhydrous HF. Owing to the corrosive nature of HF, its production is accompanied by the dissolution of silicate minerals, and, in this way, significant amounts of fluorosilicic acid are generated.[5]

Uses[edit]

Hydrofluoric acid has a variety of uses in industry and research. It is used as a starting material or intermediate in industrial chemistry, mining, refining, glass finishing, silicon chip manufacturing, and in cleaning.[18]

Oil refining[edit]

In a standard oil refinery process known as alkylation, isobutane is alkylated with low-molecular-weight alkenes (primarily a mixture of propylene and butylene) in the presence of the strong acid catalyst derived from hydrofluoric acid. The catalyst protonates the alkenes (propylene, butylene) to produce reactive carbocations, which alkylate isobutane. The reaction is carried out at mild temperatures (0 and 30 °C) in a two-phase reaction.

Production of organofluorine compounds[edit]

The principal use of hydrofluoric acid is in organofluorine chemistry. Many organofluorine compounds are prepared using HF as the fluorine source, including Teflon, fluoropolymers, fluorocarbons, and refrigerants such as freon.[5]

Production of fluorides[edit]

Most high-volume inorganic fluoride compounds are prepared from hydrofluoric acid. Foremost are Na3AlF6, cryolite, and AlF3, aluminium trifluoride. A molten mixture of these solids serves as a high-temperature solvent for the production of metallic aluminium. Given concerns about fluorides in the environment, alternative technologies are being sought. Other inorganic fluorides prepared from hydrofluoric acid include sodium fluoride and uranium hexafluoride.[5]

Etchant and cleaning agent[edit]

Wet etching tanks

In metalworking, hydrofluoric acid is used as a pickling agent to remove oxides and other impurities from stainless and carbon steels because of its limited ability to dissolve steel.[citation needed] It is used in the semiconductor industry as a major component of Wright Etch and buffered oxide etch, which are used to clean silicon wafers. In a similar manner it is also used to etch glass by reacting with silicon dioxide to form gaseous or water-soluble silicon fluorides. It can also be used to polish and frost glass.[18]

SiO2 + 4 HF → SiF4(g) + 2 H2O
SiO2 + 6 HF → H2SiF6 + 2 H2O

A 5% to 9% hydrofluoric acid gel is also commonly used to etch all ceramic dental restorations to improve bonding.[19] For similar reasons, dilute hydrofluoric acid is a component of household rust stain remover, in car washes in "wheel cleaner" compounds, in ceramic and fabric rust inhibitors, and in water spot removers.[18][20] Because of its ability to dissolve iron oxides as well as silica-based contaminants, hydrofluoric acid is used in pre-commissioning boilers that produce high-pressure steam.

Niche applications[edit]

Because of its ability to dissolve (most) oxides and silicates, hydrofluoric acid is useful for dissolving rock samples (usually powdered) prior to analysis. In similar manner, this acid is used in acid macerations to extract organic fossils from silicate rocks. Fossiliferous rock may be immersed directly into the acid, or a cellulose nitrate film may be applied (dissolved in amyl acetate), which adheres to the organic component and allows the rock to be dissolved around it.[21]

Diluted hydrofluoric acid (1 to 3 %wt.) is used in the petroleum industry in a mixture with other acids (HCl or organic acids) in order to stimulate the production of water, oil, and gas wells specifically where sandstone is involved.[citation needed]

Hydrofluoric acid is also used by some collectors of antique glass bottles to remove so-called 'sickness' from the glass, caused by acids (usually from the soil in which the bottle was buried) attacking the soda content of the glass.[citation needed]

Offset printing companies use hydrofluoric acid to remove unwanted images from printing plates. Felt-tip markers called "deletion pens" are available to make the process safer for the worker.[citation needed]

Health and safety[edit]

A hydrofluoric acid burn of the hand
left and right hands, two views, burned index fingers
HF burns, not evident until a day after

In addition to being a highly corrosive liquid, hydrofluoric acid is also a powerful contact poison. Because of the ability of hydrofluoric acid to penetrate tissue, poisoning can occur readily through exposure of skin or eyes, or when inhaled or swallowed. Symptoms of exposure to hydrofluoric acid may not be immediately evident, and this can provide false reassurance to victims, causing them to delay medical treatment.[22] Despite having an irritating odor, HF may reach dangerous levels without an obvious odor.[18] HF interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury.[22] Symptoms of HF exposure include irritation of the eyes, skin, nose, and throat, eye and skin burns, rhinitis, bronchitis, pulmonary edema (fluid buildup in the lungs), and bone damage.[23]

Once absorbed into blood through the skin, it reacts with blood calcium and may cause cardiac arrest. Burns with areas larger than 160 cm2 (25 square inches) have the potential to cause serious systemic toxicity from interference with blood and tissue calcium levels.[24] In the body, hydrofluoric acid reacts with the ubiquitous biologically important ions Ca2+ and Mg2+. Formation of insoluble calcium fluoride is proposed as the etiology for both precipitous fall in serum calcium and the severe pain associated with tissue toxicity.[25] In some cases, exposures can lead to hypocalcemia. Thus, hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca2+ that sequesters the fluoride ions. HF chemical burns can be treated with a water wash and 2.5% calcium gluconate gel[26][27][28] or special rinsing solutions.[29][30] However, because it is absorbed, medical treatment is necessary;[24] rinsing off is not enough. Intra-arterial infusions of calcium chloride have also shown great effectiveness in treating burns.[31]

Hydrogen fluoride is generated upon combustion of many fluorine-containing compounds such as products containing Viton and polytetrafluoroethylene (Teflon) parts.[32] Hydrofluorocarbons in automatic fire suppression systems can release hydrogen fluoride at high temperatures, and this has led to deaths from acute respiratory failure in military personnel when a rocket-propelled grenade hit the fire suppression system in their vehicle.[33] Hydrofluoric acid can be released from volcanoes, sea salt aerosol, and from welding or manufacturing processes.[18]

See also[edit]

References[edit]

  1. ^ Henri A. Favre; Warren H. Powell, eds. (2014). Nomenclature of Organic Chemistry: IUPAC Recommendations and Preferred Names 2013. Cambridge: The Royal Society of Chemistry. p. 131.
  2. ^ Harris, Daniel C. (2010). Quantitative Chemical Analysis (8th international ed.). New York: W. H. Freeman. pp. AP14. ISBN 1429263091.
  3. ^ "Hydrofluoric Acid". PubChem. National Institute of Health. Retrieved October 12, 2017.
  4. ^ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 921. ISBN 978-0-08-022057-4.
  5. ^ a b c d Aigueperse, J. et al. (2005) "Fluorine Compounds, Inorganic" in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, doi:10.1002/14356007.a11_307
  6. ^ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. San Diego: Academic Press. p. 425. ISBN 978-0-12-352651-9.
  7. ^ Clark, Jim (2002). "The acidity of the hydrogen halides". Retrieved 4 September 2011.
  8. ^ a b Chambers, C.; Holliday, A. K. (1975). Modern inorganic chemistry (An intermediate text) (PDF). The Butterworth Group. pp. 328–329. Archived from the original (PDF) on 2013-03-23.
  9. ^ Hannan, Henry J. (2010). Course in chemistry for IIT-JEE 2011. Tata McGraw Hill Education Private Limited. pp. 15–22. ISBN 9780070703360.
  10. ^ Ralph H. Petrucci; William S. Harwood; Jeffry D. Madura (2007). General chemistry: principles and modern applications. Pearson/Prentice Hall. p. 691. ISBN 978-0-13-149330-8. Retrieved 22 August 2011.
  11. ^ Hyman H. H., Kilpatrick M., Katz J. J. (1957). "The Hammett Acidity Function H0 for Hydrofluoric Acid Solutions". Journal of the American Chemical Society. 79 (14): 3668–3671. doi:10.1021/ja01571a016. ISSN 0002-7863.CS1 maint: uses authors parameter (link)
  12. ^ W. L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill 1984), p. 203. ISBN 0-07-032768-8.
  13. ^ F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (5th ed.) John Wiley and Sons: New York, 1988. ISBN 0-471-84997-9. p. 109.
  14. ^ C. E. Housecroft and A. G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170.
  15. ^ a b Giguère, Paul A.; Turrell, Sylvia (1980). "The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H
    3
    O+
    ...F". J. Am. Chem. Soc. 102 (17): 5473. doi:10.1021/ja00537a008.
  16. ^ Radu Iftimie; Vibin Thomas; Sylvain Plessis; Patrick Marchand; Patrick Ayotte (2008). "Spectral Signatures and Molecular Origin of Acid Dissociation Intermediates". J. Am. Chem. Soc. 130 (18): 5901–7. doi:10.1021/ja077846o. PMID 18386892.
  17. ^ a b F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, p. 104.
  18. ^ a b c d e "CDC – The Emergency Response Safety and Health Database: Systemic Agent: HYDROGEN FLUORIDE/ HYDROFLUORIC ACID – NIOSH". www.cdc.gov. Retrieved 2015-12-04.
  19. ^ Powers, John M. and Sakaguchi, Ronald L. (2006) Craig's Restorative Dental Materials, 12th ed., Mosby, ISBN 0323036066
  20. ^ Strachan, John (January 1999). "A deadly rinse: The dangers of hydrofluoric acid". Professional Carwashing & Detailing. 23 (1). Archived from the original on April 25, 2012.
  21. ^ Edwards, D. (1982). "Fragmentary non-vascular plant microfossils from the late Silurian of Wales". Botanical Journal of the Linnean Society. 84 (3): 223–256. doi:10.1111/j.1095-8339.1982.tb00536.x.
  22. ^ a b Yamashita M, Yamashita M, Suzuki M, Hirai H, Kajigaya H (2001). "Ionophoretic delivery of calcium for experimental hydrofluoric acid burns". Crit. Care Med. 29 (8): 1575–8. doi:10.1097/00003246-200108000-00013. PMID 11505130.
  23. ^ "CDC – NIOSH Pocket Guide to Chemical Hazards – Hydrogen fluoride". www.cdc.gov. Retrieved 2015-11-28.
  24. ^ a b "Recommended Medical Treatment for Hydrofluoric Acid Exposure" (PDF). Honeywell Specialty Materials. Archived from the original (PDF) on March 25, 2009. Retrieved 2009-05-06.
  25. ^ Hoffman, Robert S. et al. (2007) Goldfrank's Manual of Toxicologic Emergencies. New York: McGraw-Hill Professional, p. 1333, ISBN 0071509577.
  26. ^ el Saadi MS, Hall AH, Hall PK, Riggs BS, Augenstein WL, Rumack BH (1989). "Hydrofluoric acid dermal exposure". Vet Hum Toxicol. 31 (3): 243–7. PMID 2741315.
  27. ^ Roblin I, Urban M, Flicoteau D, Martin C, Pradeau D (2006). "Topical treatment of experimental hydrofluoric acid skin burns by 2.5% calcium gluconate". J Burn Care Res. 27 (6): 889–94. doi:10.1097/01.BCR.0000245767.54278.09. PMID 17091088.
  28. ^ "Calcium Gluconate Gel as an Antidote to HF Acid Burns". Northwestern University. Archived from the original on April 8, 2009. Retrieved 2012-10-01.
  29. ^ Hultén P, Höjer J, Ludwigs U, Janson A (2004). "Hexafluorine vs. standard decontamination to reduce systemic toxicity after dermal exposure to hydrofluoric acid". J. Toxicol. Clin. Toxicol. 42 (4): 355–61. doi:10.1081/CLT-120039541. PMID 15461243.
  30. ^ "News & Views". Chemical Health and Safety. 12 (5): 35–37. September–October 2005. doi:10.1016/j.chs.2005.07.007.
  31. ^ Siegel DC, Heard JM (March 1992). "Intra-arterial calcium infusion for hydrofluoric acid burns". Aviat Space Environ Med. 63 (3): 206–11. PMID 1567323.
  32. ^ Koch, Ernst-Christian (2002). "Metal-Fluorocarbon-Pyrolants IV: Thermochemical and Combustion Behaviour of Magnesium/Teflon/Viton (MTV)". Propellants, Explosives, Pyrotechnics. 27 (6): 340–351. doi:10.1002/prep.200290004.
  33. ^ Chauviere, Matt; Zierold, Dustin (2011-09-17). "Hydrogen Fluoride Inhalation Injury from a Fire Suppression System". NATO. Retrieved 2013-08-22.

External links[edit]