Hydrogen fluoride

Hydrogen fluoride is a chemical compound with the chemical formula HF. This colorless gas or liquid is the principal industrial source of fluorine, often as an aqueous solution called hydrofluoric acid. It is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers (e.g. Teflon). HF is widely used in the petrochemical industry as a component of superacids. Hydrogen fluoride boils near room temperature, much higher than other hydrogen halides.

Hydrogen fluoride is a highly dangerous gas, forming corrosive and penetrating hydrofluoric acid upon contact with moisture. The gas can also cause blindness by rapid destruction of the corneas.

French chemist Edmond Frémy (1814–1894) is credited with discovering anhydrous hydrogen fluoride while trying to isolate fluorine. Although Carl Wilhelm Scheele prepared hydrofluoric acid in large quantities in 1771, this acid was known in the glass industry before then.

Hydrogen fluoride
Hydrogen fluoride
Hydrogen-fluoride-2D-dimensions
Hydrogen-fluoride-3D-vdW
Names
Other names
Fluorane
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.759
KEGG
RTECS number MW7875000
UNII
UN number 1052
Properties
HF
Molar mass 20.006 g·mol−1
Appearance colourless gas or colourless liquid (below 19.5 °C)
Density 1.15 g/L, gas (25 °C)
0.99 g/mL, liquid (19.5 °C)
1.663 g/mL, solid (–125 °C)
Melting point −83.6 °C (−118.5 °F; 189.6 K)
Boiling point 19.5 °C (67.1 °F; 292.6 K)
completely miscible (liquid)
Vapor pressure 783 mmHg (20 °C)[1]
Acidity (pKa) 3.17[2][3]
Conjugate acid Fluoronium
Conjugate base Fluoride
1.00001
Structure
Linear
1.86 D
Thermochemistry
8.687 J/g K (gas)
−13.66 kJ/g (gas)
−14.99 kJ/g (liquid)
Hazards
GHS pictograms The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The skull-and-crossbones pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
GHS signal word Danger
H300, H310, H314, 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, 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 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calciumSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
4
1
Lethal dose or concentration (LD, LC):
1276 ppm (rat, 1 hr)
1774 ppm (monkey, 1 hr)
4327 ppm (guinea pig, 15 min)[4]
313 ppm (rabbit, 7 hr)[4]
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 3 ppm[1]
REL (Recommended)
TWA 3 ppm (2.5 mg/m3) C 6 ppm (5 mg/m3) [15-minute][1]
IDLH (Immediate danger)
30 ppm[1]
Related compounds
Other anions
Hydrogen chloride
Hydrogen bromide
Hydrogen iodide
Hydrogen astatide
Other cations
Sodium fluoride
Potassium fluoride
Rubidium fluoride
Caesium fluoride
Related compounds
Water
Ammonia
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Structure

The structure of chains of HF in crystalline hydrogen fluoride.

Although a diatomic molecule, HF forms relatively strong intermolecular hydrogen bonds. Solid HF consists of zigzag chains of HF molecules. The HF molecules, with a short H–F bond of 95 pm, are linked to neighboring molecules by intermolecular H–F distances of 155 pm.[5] Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.[6]

Comparison with other hydrogen halides

Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides which boil between −85 °C (−120 °F) and −35 °C (−30 °F).[7][8][9] This hydrogen bonding between HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase.

Hydrogen fluoride is miscible with water (dissolves in any proportion), whereas the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride and water also form several compounds in the solid state, most notably a 1:1 compound that does not melt until −40 °C (−40 °F), which is 44 °C (79 °F) above the melting point of pure HF.[10]

HF and H2O similarities
Boiling-points Chalcogen-Halogen HF-H2O Phase-Diagram
Boiling points of the hydrogen halides (blue) and hydrogen chalcogenides (red): HF and H2O break trends. Freezing point of HF/ H2O mixtures: arrows indicate compounds in the solid state.

Acidity

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution.[11] 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.[12] At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF
2
) and protons, thus greatly increasing the acidity.[13] This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions.[14] Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated.[13]

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),[15] 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[16](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.[17][18]

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.[19] However, Paul Giguère and Sylvia Turrell[20][21] 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.[22]

In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion.[20][22]

[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). At the limit of 100% liquid HF, there is self-ionization[23][24]

3 HF ⇌ H2F+ + HF
2

which forms an extremely acidic solution (H0 = −11).

The acidity of anhydrous HF can be increased even further by the addition of Lewis acids such as SbF5, which can reduce H0 to −21.[17][18]

Solvent

Dry hydrogen fluoride readily dissolves low-valent metal fluorides, as well as several molecular fluorides. Many proteins and carbohydrates can be dissolved in dry HF and recovered from it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving.[25]

Production and uses

Hydrogen fluoride is produced by the action of sulfuric acid on pure grades of the mineral fluorite and also as a side-product of the extraction of the fertilizer precursor phosphoric acid from various minerals. See also hydrofluoric acid.

The anhydrous compound hydrogen fluoride is more commonly used than its aqueous solution, hydrofluoric acid. HF serves as a catalyst in alkylation processes in oil refineries. A component of high-octane petrol (gasoline) called "alkylate" is generated in alkylation units that combine C3 and C4 olefins and iso-butane to generate petrol (gasoline).[26]

HF is a reactive solvent in the electrochemical fluorination of organic compounds. In this approach, HF is oxidized in the presence of a hydrocarbon and the fluorine replaces C–H bonds with C–F bonds. Perfluorinated carboxylic acids and sulfonic acids are produced in this way.[27]

Hydrogen fluoride is an important catalyst used in the majority of the installed linear alkyl benzene production in the world. The process involves dehydrogenation of n-paraffins to olefins, and subsequent reaction with benzene using HF as catalyst.

Elemental fluorine, F2, is prepared by electrolysis of a solution of HF and potassium bifluoride. The potassium bifluoride is needed because anhydrous hydrogen fluoride does not conduct electricity. Several million kilograms of F2 are produced annually.[28]

Acyl chlorides or acid anhydrides react with hydrogen fluoride to give acyl fluorides.[29]

HF is often used in palynology to remove silicate minerals, for extraction of dinoflagellate cysts, acritarchs and chitinozoans.

1,1-Difluoroethane is produced by the mercury-catalyzed addition of hydrogen fluoride to acetylene:[30]

HC≡CH + 2 HF → CH3CHF2

The intermediate in this process is vinyl fluoride or fluoroethylene, the monomeric precursor to polyvinyl fluoride.

Health effects

HF burned hands
HF burns, not evident until a day after

Upon contact with moisture, including tissue, hydrogen fluoride immediately converts to hydrofluoric acid, which is highly corrosive and toxic, and requires immediate medical attention upon exposure.[31] Breathing in hydrogen fluoride at high levels or in combination with skin contact can cause death from an irregular heartbeat or from fluid buildup in the lungs.[31]

References

  1. ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0334". National Institute for Occupational Safety and Health (NIOSH).
  2. ^ "pKa's of Inorganic and Oxo-Acids" (PDF). Harvard. Retrieved 9 September 2013.
  3. ^ Bruckenstein, S.; Kolthoff, I.M., in Kolthoff, I.M.; Elving, P.J. Treatise on Analytical Chemistry, Vol. 1, pt. 1; Wiley, NY, 1959, pp. 432-433.
  4. ^ a b "Hydrogen fluoride". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  5. ^ Johnson, M. W.; Sándor, E.; Arzi, E. (1975). "The Crystal Structure of Deuterium Fluoride". Acta Crystallographica. B31 (8): 1998–2003. doi:10.1107/S0567740875006711.
  6. ^ Mclain, Sylvia E.; Benmore, CJ; Siewenie, JE; Urquidi, J; Turner, JF (2004). "On the Structure of Liquid Hydrogen Fluoride". Angewandte Chemie International Edition. 43 (15): 1952–55. doi:10.1002/anie.200353289. PMID 15065271.
  7. ^ Pauling, Linus A. (1960). The nature of the chemical bond and the structure of molecules and crystals: An introduction to modern structural chemistry. Cornell University Press. pp. 454–464. ISBN 978-0-8014-0333-0.
  8. ^ Atkins, Peter; Jones, Loretta (2008). Chemical principles: The quest for insight. W. H. Freeman & Co. pp. 184–185. ISBN 978-1-4292-0965-6.
  9. ^ Emsley, John (1981). "The hidden strength of hydrogen". New Scientist. 91 (1264): 291–292. Retrieved 25 December 2012.
  10. ^ Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (2nd ed.). Oxford: Butterworth Heinemann. pp. 812–816. ISBN 0-7506-3365-4.
  11. ^ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. San Diego: Academic Press. p. 425. ISBN 978-0-12-352651-9.
  12. ^ Clark, Jim (2002). "The acidity of the hydrogen halides". Retrieved 4 September 2011.
  13. ^ 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.
  14. ^ Hannan, Henry J. (2010). Course in chemistry for IIT-JEE 2011. Tata McGraw Hill Education Private Limited. pp. 15–22. ISBN 9780070703360.
  15. ^ 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.
  16. ^ 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)
  17. ^ a b W. L. Jolly “Modern Inorganic Chemistry” (McGraw-Hill 1984), p. 203. ISBN 0-07-032768-8.
  18. ^ a b 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.
  19. ^ C. E. Housecroft and A. G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170.
  20. ^ 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.
  21. ^ 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.
  22. ^ a b F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, p. 104.
  23. ^ C. E. Housecroft and A. G. Sharpe Inorganic Chemistry, p. 221.
  24. ^ F. A. Cotton and G. Wilkinson Advanced Inorganic Chemistry, p. 111.
  25. ^ Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.
  26. ^ Aigueperse J, Mollard P, Devilliers D, Chemla M, Faron R, Romano R, Cuer JP (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a11_307. ISBN 3527306730.
  27. ^ G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick "Fluorine Compounds, Organic" in "Ullmann’s Encyclopedia of Industrial Chemistry" 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a11_349
  28. ^ M. Jaccaud, R. Faron, D. Devilliers, R. Romano "Fluorine" in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005 doi:10.1002/14356007.a11_293.
  29. ^ Olah G, Kuhn S (1961). "Preparation of Acyl Fluorides with Anhydrous Hydrogen Fluoride. The General Use of the Method of Colson and Fredenhagen". J. Org. Chem. 26: 237–238. doi:10.1021/jo01060a600.
  30. ^ Siegemund, Günter; Schwertfeger, Werner; Feiring, Andrew; Smart, Bruce; Behr, Fred; Vogel, Herward; McKusick, Blaine (2010). "Fluorine Compounds, Organic". In Bohnet, Matthias; Bellussi, Giuseppe; Bus, James; et al. Ullmann's Encyclopedia of Industrial Chemistry. John Wiley & Sons. doi:10.1002/14356007.a11_349.
  31. ^ a b Facts About Hydrogen Fluoride (Hydrofluoric Acid)

External links

Ammonium bifluoride

Ammonium hydrogen fluoride is the inorganic compound with the formula NH4HF2 or NH4F·HF. It is produced from ammonia and hydrogen fluoride. This colourless salt is a glass-etchant and an intermediate in a once-contemplated route to hydrofluoric acid.

Fluoride

Fluoride () is an inorganic, monatomic anion with the chemical formula F− (also written [F]−), whose salts are typically white or colorless. Fluoride salts typically have distinctive bitter tastes, and are odorless. Its salts and minerals are important chemical reagents and industrial chemicals, mainly used in the production of hydrogen fluoride for fluorocarbons. Fluoride is classified as a weak base since it only partially associates in solution, but concentrated fluoride is corrosive and can attack the skin.

Fluoride is the simplest fluorine anion. In terms of charge and size, the fluoride ion resembles the hydroxide ion. Fluoride ions occur on earth in several minerals, particularly fluorite, but are present only in trace quantities in bodies of water in nature.

Fluoroantimonic acid

Fluoroantimonic acid is an inorganic compound with the chemical formula H2FSbF6 (also written H2F[SbF6], 2HF·SbF5, or simply HF-SbF5). It is an extremely strong acid, easily qualifying as a superacid. The Hammett acidity function, H0, has been measured for different ratios of HF:SbF5. While the H0 of pure HF is −15, addition of just 1 mol % of SbF5 lowers it to around −20. However, further addition of SbF5 results in rapidly diminishing returns, with the H0 reaching −21 at 10 mol %. The use of an extremely weak base as indicator shows that the lowest attainable H0, even with > 50 mol % SbF5, is somewhere between −21 and −23.The "canonical" composition of fluoroantimonic acid is prepared by treating hydrogen fluoride (HF) with antimony pentafluoride (SbF5) in a stoichiometric ratio of 2:1. It is the strongest superacid based on measured H0 value. Only the carborane acids, whose H0 could not be directly determined due to their high melting points, may be stronger acids than fluoroantimonic acid. It has been shown to protonate even hydrocarbons to afford pentacoordinate carbocations (carbonium ions).The reaction to produce fluoroantimonic acid results in formation of the fluoronium ion:

SbF5 + 2 HF → SbF−6 + H2F+The acid is often said to contain "naked protons", but the "free" protons are, in fact, always bonded to hydrogen fluoride molecules. It is the fluoronium ion that accounts for fluoroantimonic acid's extreme acidity. The protons easily migrate through the solution, moving from H2F+ to HF, when present, by the Grotthuss mechanism:

H2F+ + HF ⇌ HF + H2F+Fluoroantimonic acid thermally decomposes at higher temperatures, generating free hydrogen fluoride gas. It is exceptionally corrosive and can only be stored in containers lined with Teflon.

Hydrofluoric acid

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 seventeenth century, even before Carl Wilhelm Scheele prepared it in large quantities in 1771. 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).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.

Hydrogen fluoride laser

The hydrogen fluoride laser is an infrared chemical laser. It is capable of delivering continuous output power in the megawatt range.

Hydrogen fluoride lasers operate at the wavelength of 2.7-2.9 µm. This wavelength is absorbed by the atmosphere, effectively attenuating the beam and reducing its reach, unless used in a vacuum environment. However, when deuterium is used instead of hydrogen, the deuterium fluoride lases at the wavelength of about 3.8 µm. This makes the deuterium fluoride laser usable for terrestrial operations.

The deuterium fluoride laser constructionally resembles a rocket engine. In the combustion chamber, ethylene is burned in nitrogen trifluoride. This reaction produces free excited fluorine radicals. Just after the nozzle, the mixture of helium and hydrogen or deuterium gas is injected to the exhaust stream; the hydrogen or deuterium reacts with the fluorine radicals, producing excited molecules of deuterium fluoride or hydrogen fluoride. The excited molecules then undergo stimulated emission in the optical resonator region of the laser.

Deuterium fluoride lasers have found military applications: the MIRACL laser, the Pulsed Energy Projectile, and the Tactical High Energy Laser are of the deuterium fluoride type.

An Argentine-American physicist and accused spy, Leonardo Mascheroni, has proposed the idea of using hydrogen fluoride lasers to produce nuclear fusion.

Hydrogen halide

Hydrogen halides are diatomic inorganic compounds with the formula HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, or astatine. Hydrogen halides are gases that dissolve in water to give acids which are commonly known as hydrohalic acids.

Indium(III) fluoride

Indium(III) fluoride or indium trifluoride is the chemical compound composed of indium and fluorine with the formula InF3. It has a rhombohedral crystal structure identical to that of rhodium(III) fluoride

.

It is formed by the reaction of indium(III) oxide with hydrogen fluoride or hydrofluoric acid.

Indium(III) fluoride is used in the synthesis of non-oxide glasses. It is also used as a catalyst for the addition of trimethylsilyl cyanide (TMSCN) to aldehydes to form cyanohydrins.

Inorganic nonaqueous solvent

An inorganic nonaqueous solvent is a solvent other than water, that is not an organic compound. Common examples are liquid ammonia, liquid sulfur dioxide, sulfuryl chloride and sulfuryl chloride fluoride, phosphoryl chloride, dinitrogen tetroxide, antimony trichloride, bromine pentafluoride, hydrogen fluoride, pure sulfuric acid and other inorganic acids. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment.

Non-bonding orbital

A non-bonding orbital, also known as non-bonding molecular orbital (NBMO), is a molecular orbital whose occupation by electrons neither increases nor decreases the bond order between the involved atoms. Non-bonding orbitals are often designated by the letter n in molecular orbital diagrams and electron transition notations. Non-bonding orbitals are the equivalent in molecular orbital theory of the lone pairs in Lewis structures. The energy level of a non-bonding orbital is typically in between the lower energy of a valence shell bonding orbital and the higher energy of a corresponding antibonding orbital. As such, a non-bonding orbital with electrons would commonly be a HOMO (highest occupied molecular orbital).

According to molecular orbital theory, molecular orbitals are often modeled by the linear combination of atomic orbitals. In a simple diatomic molecule such as hydrogen fluoride (chemical formula: HF), one atom may have many more electrons than the other. A sigma bonding orbital is created between the atomic orbitals with like symmetry. Some orbitals (e.g. px and py orbitals from the fluorine in HF) may not have any other orbitals to combine with and become non-bonding molecular orbitals. In the HF example, the px and py orbitals remain px and py orbitals in shape but when viewed as molecular orbitals are thought of as non-bonding. The energy of the orbital does not depend on the length of any bond within the molecule. Its occupation neither increases nor decreases the stability of the molecule, relative to the atoms, since its energy is the same in the molecule as in one of the atoms. For example, there are two rigorously non-bonding orbitals that are occupied in the ground state of the hydrogen fluoride diatomic molecule; these molecular orbitals are localized on the fluorine atom and are composed of p-type atomic orbitals whose orientation is perpendicular to the internuclear axis. They are therefore unable to overlap and interact with the s-type valence orbital on the hydrogen atom.

Although non-bonding orbitals are often similar to the atomic orbitals of their constituent atom, they do not need to be similar. An example of a non-similar one is the non-bonding orbital of the allyl anion, whose electron density is concentrated on the first and third carbon atoms.In fully delocalized canonical molecular orbital theory, it is often the case that none of the molecular orbitals of a molecular are strictly non-bonding in nature. However, in the context of localized molecular orbitals, the concept of a filled, non-bonding orbital tends to correspond to electrons described in Lewis structure terms as "lone pairs."

There are several symbols used to represent unoccupied non-bonding orbitals. Occasionally, n* is used, in analogy to σ* and π*, but this usage is rare. Often, the atomic orbital symbol is used, most often p for p orbital; others have used the letter a for a generic atomic orbital. (By Bent's rule, unoccupied orbitals for a main-group element are almost always of p character, since s character is stabilizing and will be used for bonding orbitals. As an exception, the LUMO of phenyl cation is an spx (x ≈ 2) atomic orbital, due to the geometric constraint of the benzene ring.) Finally, woodward and Hoffmann used the letter ω for non-bonding orbitals (occupied or unoccupied) in their monograph Conservation of Orbital Symmetry.

Olah reagent

The Olah reagent is a nucleophilic fluorinating agent. It consists of a mixture of 70% hydrogen fluoride and 30% pyridine; alcohols react with this reagent to give alkyl fluorides:

It acts as a stabilized, less volatile form of hydrogen fluoride. It is used in the fluorination of steroids and in deprotection of peptides. Instead of hydrogen fluoride, several other fluorinating agents can be used, such as diethylaminosulfur trifluoride (DAST).

Perfluorotripentylamine

Perfluorotripentylamine is a perfluorocarbon. It is used as an electronics coolant, and has a high boiling point. It is colorless, odorless, and insoluble in water. Unlike ordinary amines, perfluoroamines are of low basicity. Perfluorinated amines are components of fluorofluids, used as immersive coolants for supercomputers.It is prepared by electrofluorination of the amine using hydrogen fluoride as solvent and source of fluorine:

N(C5H11)3 + 33 HF → N(C5F11)3 + 33 H2

Potassium bifluoride

Potassium bifluoride is the inorganic compound with the formula KHF2. This colourless salt consists of the potassium cation and the bifluoride (HF2−) anion. The salt is used in etchant for glass. Sodium bifluoride is related and is also of commercial use as an etchant as well as in cleaning products.

Potassium fluoride

Potassium fluoride is the chemical compound with the formula KF. After hydrogen fluoride, KF is the primary source of the fluoride ion for applications in manufacturing and in chemistry. It is an alkali halide and occurs naturally as the rare mineral carobbiite. Solutions of KF will etch glass due to the formation of soluble fluorosilicates, although HF is more effective.

Potassium hexafluoronickelate(IV)

Potassium hexafluoronickelate(IV) is an inorganic compound with the chemical formula K
2
NiF
6
. It can be produced through the reaction of potassium fluoride, nickel dichloride, and fluorine.

It reacts violently with water, releasing oxygen. It dissolves in anhydrous hydrogen fluoride to produce a light-red solution. Potassium hexafluoronickelate(IV) decomposes at 350 °C, forming potassium hexafluoronickelate(III), nickel(II) fluoride, and fluorine:[better source needed]

Potassium hexafluoronickelate is a strong oxidant. It can turn chlorine pentafluoride and bromine pentafluoride into ClF+
6
and BrF+
6
, respectively:

( X = Cl or Br , -60 °C , aHF = anhydrous hydrogen fluoride).

It adopts the structure seen for K2PtCl6.

Potassium hexafluorophosphate

Potassium hexafluorophosphate is the chemical compound with the formula KPF6. This colourless salt consists of potassium cations and hexafluorophosphate anions. It is prepared by the reaction:

PCl5 + KCl + 6 HF → KPF6 + 6 HClThis exothermic reaction is conducted in liquid hydrogen fluoride. The salt is stable in hot alkaline aqueous solution, from which it can be recrystallized. The sodium and ammonium salts are more soluble in water whereas the rubidium and caesium salts are less so.

KPF6 is a common laboratory source of the hexafluorophosphate anion, a non-coordinating anion that confers lipophilicity to its salts. These salts are often less soluble than the closely related tetrafluoroborates.

Protic solvent

In chemistry, a protic solvent is a solvent that has a hydrogen atom bound to an oxygen (as in a hydroxyl group), a nitrogen (as in an amine group) or a fluorine (as in hydrogen fluoride). In general terms, any solvent that contains a labile H+ is called a protic solvent. The molecules of such solvents readily donate protons (H+) to reagents. Conversely, aprotic solvents cannot donate hydrogen.

SathyabamaSat

SathyabamaSat is a micro experimental satellite developed by students and faculty of Sathyabama University, Chennai to collect data on greenhouse gases (water vapor, carbon monoxide, carbon dioxide, methane and hydrogen fluoride). It was launched along with the Cartosat-2C satellite atop PSLV-C34. It was launched June 22, 2016.

Stishovite

Stishovite is an extremely hard, dense tetragonal form (polymorph) of silicon dioxide. It is very rare on the Earth's surface, however, it may be a predominant form of silicon dioxide in the Earth, especially in the lower mantle.Stishovite was named after Sergey M. Stishov, a Russian high-pressure physicist who first synthesized the mineral in 1961. It was discovered in Meteor Crater in 1962 by Edward C. T. Chao.Unlike other silica polymorphs, the crystal structure of stishovite resembles that of rutile (TiO2). The silicon in stishovite adopts an octahedral coordination geometry, being bound to six oxides. Similarly, the oxides are three-connected, unlike low-pressure forms of SiO2. In most silicates, silicon is tetrahedral, being bound to four oxides. It was long considered the hardest known oxide (~30 GPa Vickers); however, boron suboxide has been discovered in 2002 to be much harder. At normal temperature and pressure, stishovite is metastable.

Stishovite can be separated from quartz by applying hydrogen fluoride (HF); unlike quartz, stishovite will not react.

Uranium hexachloride

Uranium hexachloride is an inorganic chemical compound of uranium in the +6 oxidation state. The chemical compound Uranium hexachloride (UCl6) is a metal halide composed of Uranium and Chlorine. It is a multi-luminescent dark green crystalline solid with a vapor pressure between 1-3 mmHg at 373.15K UCl6 is stable in a vacuum, dry air, nitrogen and helium at room temperature. It is soluble in carbon tetrachloride (CCl4). Compared to the other uranium halides, little is known about UCl6.

Hydrogen compounds
Molecules
Deuterated
molecules
Unconfirmed
Related
Other
Alkali metal hydrides
Alkaline earth hydrides
Group 13 hydrides
Group 14 hydrides
Pnictogen hydrides
Hydrogen chalcogenides
Hydrogen halides
Transition metal hydrides
Lanthanide hydrides
Actinide hydrides

This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.