Hydrofluoric acid


Free Web Hosting with Website Builder
Hydrofluoric acid
Hydrofluoric acid
Hydrogen fluoride molecule
Other names fluoric acid; fluorhydric acid
Identifiers
CAS number 7664-39-3
RTECS number MW7875000
Properties
Molecular formula HF
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)
Solubility in water Miscible.
Acidity (pKa) 3.15 (in water)
Hazards
MSDS External MSDS
MSDS duPont MSDS
Main hazards Very Toxic, Corrosive.
NFPA 704
0
4
1
COR
R-phrases R26/27/28, R35
S-phrases (S1/2), S7/9, S26, S36/37, S45
Flash point nonflammable
Related compounds
Other anions Hydrochloric acid
Hydrobromic acid
Hydroiodic acid
Related compounds Hydrogen fluoride
fluorosilicic acid
Supplementary data page
Structure and
properties		 n, εr, etc.
Thermodynamic
data		 Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)
Infobox references

Hydrofluoric acid is a solution of hydrogen fluoride in water. It is a weak acid. Hydrogen fluoride, often in the aqueous form as hydrofluoric acid, is a valued source of fluorine, being the precursor to numerous pharmaceuticals (e.g., Prozac), diverse polymers (e.g., Teflon), and most other synthetic materials that contain fluorine. Hydrofluoric acid is best known to the public for its ability to dissolve glass by reacting with SiO2, the major component of most glasses. This property has been known since the 17th century, even before hydrofluoric acid had been prepared in large quantities by Scheele in 1771.[1] This dissolution process can be described as follows:

SiO2(s) + 4 HF(aq) → SiF4(g) + 2 H2O(l)
SiO2(s) + 6HF(aq) → H2[SiF6](aq) + 2H2O(l)

Because of its high reactivity toward glass, hydrofluoric acid must be stored (for small quantities) in polyethylene or Teflon containers. It is also unique in its ability to dissolve many metal and semimetal oxides. It is corrosive, as explained below.

Contents

Acidity

Hydrogen fluoride ionizes in aqueous solution in a similar fashion to other common acids:

HF + H2O → H3O+ + F−

When the concentration of HF approaches 100%, the acidity increases dramatically due to the following equilibrium:

2HF → H+ + FHF−

The FHF− anion is stabilized by the very strong hydrogen - fluorine hydrogen bond. Hydrofluoric acid is the only hydrohalic acid that is not considered a strong acid due to its low degree of ionization in aqueous solution.

Production

Main article: hydrogen fluoride

Industrially, hydrofluoric acid is produced by treatment of the mineral fluorite (CaF2) with concentrated sulfuric acid. When combined at 250 °C, these two substances react to produce hydrogen fluoride according to the following chemical equation:

CaF2 + H2SO4 → 2HF + CaSO4

Hydrogen fluoride is generated upon combustion of many fluorine-containing compounds such as products containing Viton and Teflon parts. Hydrogen fluoride converts immediately to hydrofluoric acid upon contact with liquid water.

Uses

Because of its ability to dissolve metal oxides, hydrofluoric acid is used in the purification of both aluminium and uranium. It is also used to etch glass, to remove surface oxides from silicon in the semiconductor industry, as a catalyst for the alkylation of isobutane and butene (olefinic C4) in oil refineries, and to remove oxide impurities from stainless steel in a process called pickling. Dilute hydrofluoric acid is sold as a household rust stain remover. Recently it has even been used in car washes in "wheel cleaner" compounds.[2] Due to its ability to dissolve silicate compounds, hydrofluoric acid is often used to dissolve rock samples (usually powdered) prior to analysis. Similarly hydrofluoric acid attacks many metal oxides, forming the corresponding fluoro derivatives.

Hydrofluoric acid is also used in the synthesis of many fluorine-containing organic compounds, including Teflon, fluoropolymers, perfluorocarbons, and refrigerants such as freon. Additionally, hydrofluoric acid is commonly used in refinery alkylation processes to produce a high-octane gasoline blending component called alkylate from FCCU C3 and C4 olefins and isobutane.

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.

HF is also 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.[3]

Safety

Hydrofluoric acid is corrosive and a contact poison. It should be handled with extreme care, beyond that accorded to other mineral acids, in part because of its low dissociation constant, which allows HF to penetrate tissue more quickly. Symptoms of exposure to hydrofluoric acid may not be immediately evident. HF interferes with nerve function and burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury.[4] HF is known to etch bone, and since it penetrates the skin it can weaken bones without destroying the skin.[5] More seriously, it can absorb into blood through skin and react with blood calcium, causing cardiac arrest.

In the body, hydrofluoric acid reacts with the ubiquitous biologically important ions Ca2+ and Mg2+. 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[6][7][8] or special rinsing solutions.[9][10] However, because it is absorbed, medical treatment is necessary — rinsing off is not enough. In some cases, amputation may be required.

References

  1. ^ Greenwood, Norman N.; Earnshaw, A. (1984). Chemistry of the Elements. Oxford: Pergamon, 921. ISBN 0-08-022057-6. 
  2. ^ Strachan, John (January, 1999). "A deadly rinse: The dangers of hydrofluoric acid". Professional Carwashing & Detailing, http://www.carwash.com/article.asp?IndexID=4230101. Retrieved on 30 August 2006. 
  3. ^ 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 
  4. ^ Yamashita M, Yamashita M, Suzuki M, Hirai H, Kajigaya H (2001). "Iontophoretic delivery of calcium for experimental hydrofluoric acid burns". Crit. Care Med. 29 (8): 1575–8. doi:10.1097/00003246-200108000-00013. PMID 11505130. 
  5. ^ emedicine.com topic804
  6. ^ 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. 
  7. ^ 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 (inactive 20 June 2008). PMID 17091088. 
  8. ^ "Calcium Gluconate Gel as an Antidote to HF Acid Burns". Northwestern University. Retrieved on 2008-01-03.
  9. ^ 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. PMID 15461243. 
  10. ^ "News & Views". Chemical Health and Safety 12 (5): 35–37. September-October 2005. doi:10.1016/j.chs.2005.07.007. 

External links







Why are we here?
All text is available under the terms of the GNU Free Documentation License
This page is cache of Wikipedia. History