Compounds of Uranium and Fluorine (Chemical Compounds)

Author: Christopher Estep

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First Edition, 2012

ISBN 978-81-323-3416-3

© All rights reserved.

Published by: Research World 4735/22 Prakashdeep Bldg, Ansari Road, Darya Ganj, Delhi - 110002 Email: [email protected]

Table of Contents Chapter 1 - Uranium Compounds Chapter 2 - Uranium Hexafluoride and Uranium Hydride Chapter 3 - Uranium Dioxide Chapter 4 - Uranium Trioxide Chapter 5 - Fluorocarbon Chapter 6 - Chlorofluorocarbon Chapter 7 - Perfluorocarbon Chapter 8 - Tetrafluoroborates Chapter 9 - Fluoride Compounds

Chapter-1

Uranium Compounds

Uranium boride Uranium boride

Identifiers CAS number

12007-36-2 Properties

Molecular formula Molar mass Density Melting point Solubility in other solvents

UB2 259.651 g/mol 12.7 g/cm3 2430 °C x.xx g/l

Uranium boride (UB2), a compound of uranium and boron, is a very stable glassy boride material that is insoluble in water. It is being explored as a method of immobilising uranium based radioactive waste, and rendering it safe for long term storage. Some applications in endocurietherapy, a method of radiation therapy where in radioactive microspheres are implanted directly into the treatment site and allowed to remain for an extended period of time, may also use this class of material as it would not be attacked while in situ.

Uranium borohydride Uranium borohydride Identifiers ChemSpider

15385433 Properties

Molecular formula Molar mass Solubility in other solvents

U(BH4)4 297.27 g/mol Decomposes

Uranium borohydride U(BH4)4 is a volatile uranium complex with borohydride. This green-coloured compound is polymeric in the solid but evaporates to a tetrahedral monomer.

Preparation This compound was first prepared by reacting uranium tetrafluoride with aluminium borohydride: UF4 + 2 Al(BH4)3 → U(BH4)4 + 2 Al(BH4)F2 It may also be prepared by the solid state reaction of uranium tetrachloride with lithium borohydride in vacuo: UCl4 + 4 LiBH4 → U(BH4)4 + 4 LiCl Although U(BH4)4 is polymeric in the solid state, U(BH3CH3)4 is monomeric and hence more volatile

History During the Manhattan Project the need arose to find volatile compounds of uranium suitable for use in the diffusion separation of uranium isotopes. Uranium borohydride is, after uranium hexafluoride, the most volatile compound of uranium known with a vapor pressure of 4 mmHg (530 Pa) at 60 °C. Uranium borohydride was discovered by Hermann Irving Schlesinger and Herbert C. Brown, who also discovered sodium borohydride. Uranium hexafluoride is very corrosive, which presented serious handling difficulties, leading to serious consideration of the borohydride. However by the time the synthesis method was finalized, the uranium hexafluoride related problems were already solved. Borohydrides are nonideal ligands for isotope separations since boron occurs naturally with two abundant isotopes, 10B (20%) and 11B (80%).

Enrico Fermi's purported comment when he observed the neutron cross section for boron—"My God! It's as big as the side of a barn!"—not only gave a name to the unit of cross section (the barn), it also put an end to using uranium borohydride in the diffusion process.

Sodium uranate Sodium uranate Molecular formula Molar mass

Sodium uranate Properties Na2O7U2 634.03 g mol−1

Sodium uranate or Yellow uranium oxide, a uranium compound with the chemical formula Na2O (UO3)2·6H2O is a yellow orange powder once used in pottery to produce ivory to yellow shades in glazes. It was also used in porcelain dentures to give them a fluorescence similar to that of natural teeth. It was added as a mix with cerium oxide. The uranium composed from 0.008 to 0.1% by weight uranium with an average of about 0.02%. The practice appears to have stopped in the late 1980s. The alkaline process of milling uranium ores involves precipitating sodium uranate from the pregnant leaching solution to produce the semi-refined product referred to as yellowcake.

Uranium(III) chloride Uranium(III) chloride

CAS number PubChem

IUPAC name Uranium(III) chloride Other names Uranium chloride, Uranium trichloride Identifiers 10025-93-1 167444

ChemSpider Molecular formula Molar mass Appearance Density Melting point Boiling point Solubility in water Hybridisation Flash point Autoignition temperature

146484 Properties Cl3U 344.39 g mol−1 Green crystalline solid 5.500 g/cm³, liquid 837 °C, 1110 K, 1539 °F 1657 °C, 1930 K, 3015 °F Soluble Structure Tricapped trigonal prismatic Hazards Non-flammable Non-flammable

Related compounds Uranium(IV) chloride, Related compounds Uranium(V) chloride, Uranium(VI) chloride

Uranium(III) chloride , UCl3, is a chemical compound that contains the earth metal uranium and chlorine. UCl3 is used mostly to reprocess spent nuclear fuel. Uranium(III) chloride is synthesized various ways from uranium(IV) chloride; however, UCl3 is less stable than UCl4.

Preparation There are two ways to synthesize uranium(III) chloride. The following processes describe how to produce uranium(III) chloride. (1) In a mixture of NaCl-KCl at 670-710 °C, add uranium tetrachloride with uranium metal. 3UCl4 + U → 4UCl3 (2) Heat uranium(IV) chloride in hydrogen gas. 2UCl4 + H2 → 2UCl3 + 2HCl

Properties In solid uranium(III) chloride each uranium atom has nine chlorine atoms as near neighbours, at approximately the same distance, in a tricapped trigonal prismatic configuration.

Uranium(III) chloride is a green crystalline solid at room temperature. UCl3 melts at 837 °C and boils at 1657 °C. Uranium(III) chloride has a density of 5500 kg/m³ or 5.500 g/cm³. Its composition by weight: Chlorine: 30.84% Uranium: 69.16% Its formal oxidative states: Chlorine: −1 Uranium: +3 Uranium(III) chloride is very soluble in water and is also very hygroscopic. UCl3 is more stable in a solution of hydrochloric acid.

Uses Reagent Uranium(III) chloride is used in reactions with tetrahydrofuran (THF) and sodium methylcyclopentadiene to prepare various uranium metallocene complexes.

Catalyst Uranium(III) chloride is used as a catalyst during reactions between lithium aluminium hydride (LiAlH4) and olefins to produce alkyl aluminate compounds.

Molten form The molten form of uranium(III) chloride is a typical compound in pyrochemical processes as it is important in the reprocessing of spent nuclear fuels.UCl3 is usually the form that uranium takes as spent fuel in electrorefining processes.,

Hydrates There are three hydrates of uranium(III) chloride: 1. UCl3.2H2O.2CH3CN 2. UCl3.6H2O 3. UCl3.7H2O Each are synthesized by the reduction of uranium(IV) chloride in methylcyanide (acetonitrile), with specific amounts of water and propionic acid.

Precautions While there are no long-term data on the toxic effects thas UCl3, it is important to minimize exposure to this compound when possible. Similar to other uranium compounds that are soluble, UCl3 is likely absorbed into the blood through the alveolar pockets of the lungs within days of exposure. Exposure to uranium(III) chloride leads to toxicity of the renal system.

Uranium tetrachloride Uranium tetrachloride

CAS number ChemSpider Molecular formula Molar mass Density Melting point Boiling point

IUPAC name Tetrachlorouranium Other names Uranium (IV) Chloride Identifiers 10026-10-5 19969614 Properties UCl4

379.84 g/mol 4.87 g/cm3 590°C 791°C Structure Crystal structure Octahedral Related compounds Related uranium trichloride, uranium pentachloride, compounds uranium hexachloride

Uranium tetrachloride (UCl4) is compound of uranium in oxidation state +4. It was used in the electromagnetic isotope separation (EMIS) process of uranium enrichment.

Chemical properties Uranium tetrachloride is a hygroscopic, dark green solid, which sublimes in a high vacuum at ca. 500°C. The crystal structure shows the uranium to be surrounded by eight chlorine atoms, four at 264 pm and the other four at 287pm. The molecule UCl4 is a Lewis acid and dissolves in solvents that can act as non-protic Lewis bases. Dissolution in protic solvents is more complicated. When UCl4 is added to water the uranium aqua ion is formed. UCl4 + xH2O → [U(H2O)x]4+ + 4Cl-

The aqua ion [U(H2O)x]4+, (x is 8 or 9) is strongly hydrolyzed. [U(H2O)x]4+

[U(H2O)x-1(OH)]3+ + H+

The pKa for this reaction is ca. 1.6, so hydrolysis is absent only in solutions of acid strength 1 mol dm-3 or stronger (pH < 0). Further hydrolysis occurs at pH > 3. Weak chloro complexes of the aqua ion may be formed. Published estimates of the log K value for the formation of [UCl]3+(aq) vary from -0.5 to +3 because of difficulty in dealing with simultaneous hydrolysis. With alcohols, partial solvolysis may occur. UCl4 + xROH

UCl4-x(OR)x + xH+

Uranium tetralchloride dissolves in non-protic solvents such as tetrahydrofuran, acetonitrile, dimethyl formamide etc. that can act as Lewis bases. Solvates of formula UCl4Lx are formed which may be isolated. The solvent must be completely free of dissolved water, or hydrolysis will occur, with the solvent, S, picking up the released proton. UCl4 +H2O + S

UCl3(OH) + SH+ +Cl-

The solvent molecules may be replaced by other ligand in a reaction such as UCl4 + 2Cl- → [UCl6]2-. The solvent is not shown, just as when complexes of other metal ions are formed in aqueous solution. Solutions of UCl4 are susceptible to oxidation by air, resulting in the production of complexes of the uranyl ion.

Applications Uranium tetrachloride is produced commercially by the reaction of carbon tetrachloride with pure uranium dioxide UO2 at 370°C. It has been used as feed in the electromagnetic isotope separation (EMIS) process of uranium enrichment. Beginning in 1944, the Oak Ridge Y-12 Plant converted UO3 to UCl4 feed for the for Ernest O. Lawrence’s Alpha Calutrons. Its major benefit being the uranium tetrachloride used in the calutrons is not as corrosive as the uranium hexafluoride used in most other enrichment technologies This process was abandoned in the 1950s. In the 1980s, however, Iraq unexpectedly revived this option as part of its nuclear weapons program. In the enrichment process, uranium tetrachloride is ionized into a uranium plasma.The uranium ions are then accelerated and passed through a strong magnetic field. After traveling along half of a circle the beam is split into a region nearer the outside wall which is depleted and a region nearer the inside wall which is enriched in U-235. The large amounts of energy required in maintaining the

strong magnetic fields as well as the low recovery rates of the uranium feed material and slower more inconvenient facility operation make this an unlikely choice for large scale enrichment plants. Work is being done in the use of molten uranium chloride-alkali chloride mixtures as reactor fuels in molten salt reactors. Uranium tetrachloride melts dissolved in a lithium chloride-potassium chloride eutectic have also been explored as a means to recover actinides from irradiated nuclear fuels through pyrochemical nuclear reprocessing.

Uranium tetrafluoride Uranium tetrafluoride

IUPAC name Uranium(IV) fluoride Uranium tetrafluoride Identifiers CAS number 10049-14-6 ChemSpider 14676181 Properties Molecular formula UF4 Molar mass 314.02 g/mol Appearance Green crystalline solid. Density 6.70 g/cm3, solid. Melting point 1036 °C Boiling point 1417 °C Solubility in water Insoluble. Structure Crystal structure Monoclinic, mS60 Space group C2/c, No. 15 Hazards MSDS External MSDS EU Index 092-002-00-3 Very toxic (T+) EU classification Dangerous for the environment (N) R-phrases R26/28, R33, R51/53

S-phrases Flash point

(S1/2), S20/21, S45, S61 Non-flammable Related compounds Uranium(IV) chloride Other anions Uranium(IV) bromide Uranium(IV) iodide Thorium(IV) fluoride Protactinium(IV) fluoride Other cations Neptunium(IV) fluoride Plutonium(IV) fluoride Uranium hexafluoride Related compounds Uranium dioxide

Uranium tetrafluoride (UF4) is a green crystalline solid compound of uranium with an insignificant vapor pressure and very slight solubility in water. Uranium in its tetravalent (uranous) state is very important in different technological processes. In the uranium refining industry it is known as green salt. UF4 is generally an intermediate in the conversion of uranium hexafluoride (UF6) to either uranium oxides (U3O8 or UO2) or uranium metal. It is formed by the reaction of UF6 with hydrogen gas in a vertical tube-type reactor or by the action of hydrogen fluoride(HF) on uranium dioxide. UF4 is less stable than the uranium oxides and reacts slowly with moisture at ambient temperature, forming UO2 and HF, which are very corrosive; it is thus a less favorable form for long-term disposal. The bulk density of UF4 varies from about 2.0 g/cm3 to about 4.5 g/cm3 depending on the production process and the properties of the starting uranium compounds. A molten salt reactor design, a type of nuclear reactor where the working fluid is a molten salt, would use UF4 as the core material. UF4 is generally chosen over other salts because of the usefulness of the elements without isotope separation, better neutron economy and moderating efficiency, lower vapor pressure and better chemical stability. Like all uranium salts UF4 is toxic and thus harmful by inhalation, ingestion and through skin contact. Being radioactive it may also cause cancer, probably through exposure to the breakdown product radon gas and its daughters.

Ammonium diuranate Ammonium diuranate

Abbreviations CAS number ChemSpider Molecular formula Molar mass

Identifiers ADU 7783-22-4 170692 Properties H8N2O7U2 624.13 g mol−1

Ammonium diuranate or (ADU) ((NH4)2U2O7), is one of the intermediate chemical forms of uranium produced during yellowcake production. The name 'yellowcake' originally given to this bright yellow substance, now applies to mixtures of uranium oxides which are actually hardly ever yellow. It also is an intermediate in mixed-oxide (MOX) fuel fabrication. It is precipitated during production by adding aqueous ammonium hydroxide after uranium extraction by tertiary amines in an organic kerosene solvent. This precipitate is then thickened and centrifuged before being calcined to uranium oxide. Canadian practice favours the production of uranium oxide from ammonium diuranate, rather than from uranyl nitrate as is the case elsewhere. Ammonium diuranate was once used to produce colored glazes in ceramics.

Uranyl peroxide Uranyl peroxide Molecular formula Molar mass

Properties UO4·nH2O 302.03 g/mol (as UO4)

Uranyl peroxide or uranium peroxide hydrate (UO4·nH2O) is a pale-yellow, soluble peroxide of uranium. It is found present at one stage of the enriched uranium fuel cycle and in yellowcake prepared via the in situ leaching and resin ion exchange system. This compound, also expressed as: UO3·(H2O2)·(H2O), is very similar to uranium trioxide hydrate UO3·nH2O. The dissolution behaviour of both compounds are very sensitive to the hydration state (n can vary between 0 and 4). One main characteristic of uranium peroxide is that it consists of small needles with an average AMAD of about 1.1 µm.

The uranyl minerals Studtite, UO4·4H2O, and metastudtite, UO4·2H2O, are the only minerals discovered to date found to contain peroxide.

Triuranium octoxide Triuranium octoxide

Other names pitchblende Identifiers CAS number 1317-99-3 Properties Molecular formula U3O8 Molar mass 842.1 g/mol Melting point 1150°C Boiling point decomposes to UO2 at 1300 °C Insoluble in water; Solubility in other solvents Soluble in nitric and sulfuric acids.

Triuranium octoxide (U3O8) is a compound of uranium. It is present as an olive green to black, odorless solid. In spite of its color, it is one of the more popular forms of yellowcake and is shipped between mills and refineries in this form. Triuranium octoxide occurs naturally as the olive-green-colored mineral pitchblende. U3O8 is readily produced from UF6 and has potential long-term stability in a geologic environment. In the presence of oxygen (O2), uranium dioxide (UO2) is oxidized to U3O8, whereas uranium trioxide (UO3) loses oxygen at temperatures above 500°C and is reduced to U3O8. The compound can be produced by any one of three primary chemical conversion processes, involving either uranium tetrafluoride (UF4) or uranyl fluoride (UO2F2) as intermediates. It is generally considered to be the more attractive form for disposal purposes because, under normal environmental conditions, U3O8 is one of the most kinetically and thermodynamically stable forms of uranium and also because it is the form of uranium found in nature. Its particle density is 8.3 g cm−3.

Solid state structure The solid is a layered structure where the layers are bridged by oxygen atoms, each layer contains uranium atoms which are in different coordination environments in the above diagram these are shown in plum and green.

Bond valence study Using a 6Å x 6Å x 6Å box with the uranium atom in the centre the bond valence calculation was performed for both U1 and U2 in solid. It was found using the parameters for U(VI) that the calculated oxidation states for U1 and U2 are 5.11 and 5.10. Using the parameters for U(IV) the calculated oxidation states are 5.78 and 5.77 respectively for U1 and U2. These study suggests that all the uranium atoms have the same oxidation state, so that the oxidation states are disordered through the lattice.

Uranyl acetate Uranyl acetate

IUPAC name Uranium bis(acetato)-O)dioxo-dihydrate Other names Uranyl ethanoate Identifiers 541-09-3 , (anhydrous) CAS number [6159-44-0] (dihydrate) Properties UO2(CH3COO)2 (anhydrous) Molecular formula UO2(CH3COO)2·2H2O (dihydrate) Molar mass 424.146 g/mol (dihydrate) Appearance yellow crystals (dihydrate) Density 2.89 g/cm3 (dihydrate) Melting point decomposes at 80°C (dihydrate) Solubility slightly soluble in ethanol

MSDS

Hazards External MSDS

Uranyl acetate (UO2(CH3COO)2·2H2O) is the acetate salt of uranium and is a yellow crystalline solid made up of yellow rhombic crystals and has a slight acetic odor. Uranyl acetate is slightly radioactive, the precise radioactivity depends on the isotopes of uranium present. This compound is a nuclear fuel derivative, and its use and possession are sanctioned by international law.

Production Commercial preparations of uranyl acetate are usually made from depleted uranium and are prepared by reacting metallic uranium with acetic acid.

Uses Uranyl acetate is extensively used as a negative stain in electron microscopy, most procedures in electron microscopy for biology require the use of uranyl acetate. Negative staining protocols typically treat the sample with 1% to 5% aqueous solution. Uranyl acetate staining is simple and quick to perform and one can examine the sample within a few minutes after staining. Some biological samples are not amenable to uranyl acetate staining and, in these cases, alternative staining techniques and or low-voltage electron microscopy technique may be more suitable. 1% and 2% uranyl acetate solutions are used as an indicator, and a titrant in stronger concentrations in analytical chemistry, as it forms an insoluble salt with sodium (the vast majority of sodium salts are water-soluble). Uranyl acetate solutions show evidence of being sensitive to light, especially UV and will precipitate if exposed. Uranyl acetate is also used in a standard test—American Association of State Highway and Transportation Officials (AASHTO) Designation T 299—for alkali-silica reactivity in aggregates (crushed stone or gravel) being considered for use in cement concrete.

Safety Uranyl acetate is both radioactive and toxic. Normal commercial stocks prepared from depleted uranium have a typical radioactivity of 0.37 - 0.51 µCi/g. This is a very mild level of radioactivity and is not sufficient to be harmful while the material remains external to the body. Uranyl acetate is very toxic if ingested, inhaled as dust or by skin contact if skin is cut or abraded. The toxicity is due to the combined effect of chemical toxicity and mild radioactivity and there is a danger of cumulative effects from long term exposure.

Uranium carbonate Uranyl carbonate

ChemSpider Molecular formula Molar mass

IUPAC name Uranium carbonate Other names Uranium Carbonate Identifiers 14272213 Properties UO2(CO3) 330 g/mol

Uranyl carbonate, UO2(CO3), is a carbonate of uranium that forms the backbone of several uranyl mineral species such as Andersonite, McKelveyite and Wyartite and most importantly Rutherfordine. It is also found in both the mineral and organic fractions of coal and its fly ash and is the main component of uranium in mine tailing seepage water. Uranium like other actinides readily forms a dioxide uranyl core (UO2). In the environment, this uranyl core readily complexes with carbonate to form charged complexes. Although uranium forms insoluble solids or adsorbs to mineral surfaces at alkaline pH it is these soluble carbonate complexes that increase its solubility, availability, and mobility with low affinities to soil. Uranium(VI) generally forms a pHdependent suite of uranyl-carbonate complexes in ground water solutions:    

UO2(OH)2+1 UO2(CO3)2−2 UO2(CO3)3−4 UO2(CO3)(OH)3−1

A common method for concentrating uranium from a solution uses solutions of uranyl carbonates which are passed through a resin bed where the complex ions are transferred to the resin by ion exchange with a negative ion like chloride. After build-up of the uranium complex on the resin, the uranium is eluted with a salt solution and the uranium is precipitated in another process.

The uranyl carbonate minerals Uranyl-carbonate complexes form a large class of mineral species. Several have been described in literature. These include:                      

Andersonite (Hydrated Sodium Calcium Uranyl Carbonate) Astrocyanite-(Ce) (Hydrated Copper Cerium Neodymium Lanthanum Praseodymium Samarium Calcium Yttrium Uranyl Carbonate Hydroxide) Bayleyite (Hydrated Magnesium Uranyl Carbonate) Bijvoetite-(Y) (Hydrated Yttrium Dysprosium Uranyl Carbonate Hydroxide) Fontanite (Hydrated Calcium Uranyl Carbonate) Grimselite (Hydrated Potassium Sodium Uranyl Carbonate) Joliotite (Hydrated Uranyl Carbonate) Liebigite (Hydrated Calcium Uranyl Carbonate) McKelveyite (Hydrated Barium Sodium Calcium Uranium Yttrium Carbonate) Metazellerite (Hydrated Calcium Uranyl Carbonate) Rabbittite (Hydrated Calcium Magnesium Uranyl Carbonate Hydroxide) Roubaultite (Copper Uranyl Carbonate Oxide Hydroxide) Rutherfordine (Uranyl Carbonate) Schrokingerite (Hydrated Sodium Calcium Uranyl Sulfate Carbonate Fluoride) Shabaite (Hydrated Copper Cerium Neodymium Lanthanum Praseodymium Samarium Calcium Yttrium Uranyl Carbonate Hydroxide) Sharpite (Hydrated Calcium Uranyl Carbonate Hydroxide) Swartzite (Hydrated Calcium Magnesium Uranyl Carbonate) Voglite (Hydrated Calcium Copper Uranyl Carbonate) Wyartite (Hydrated Calcium Uranyl Carbonate Hydroxide) Widenmannite (Lead Uranyl Carbonate) Zellerite (Hydrated Calcium Uranyl Carbonate) Znucalite (Hydrated Calcium Zinc Uranyl Carbonate Hydroxide)

Ammonium uranyl carbonate Ammonium uranyl carbonate

IUPAC name uranium(VI)dioxide di-ammonium carbonate Other names uranyl ammonium carbonate Properties Molecular formula UO2CO3·2(NH4)2CO3 Molar mass 522.199 g/mol Melting point Decomposes between 165°C and 185°C Solubility in water Insoluble

Ammonium uranyl carbonate(UO2CO3·2(NH4)2CO3) is known in the uranium processing industry as AUC and is also called uranyl ammonium carbonate. This compound is important as a component in the conversion process of uranium hexafluoride (UF6) to uranium dioxide (UO2). The ammonium uranyl carbonate is combined with steam and hydrogen at 500-600°C to yield UO2. In another process aqueous uranyl nitrate, known as uranyl nitrate liquor (UNL) is treated with ammonium bicarbonate to form ammonium uranyl carbonate as a solid precipitate. This is separated from the solution, dried with methanol and then calcinated with hydrogen directly to UO2 to obtain a sinterable grade powder. The ex-AUC uranium dioxide powder is freeflowing, relatively coarse (10 µ) and porous with specific surface area in the range of 5m2/g and suitable for direct pelletisation, avoiding the granulation step. Conversion to UO2 is often performed as the first stage of nuclear fuel fabrication. The AUC process is followed in South Korea and Argentina. In the AUC route, calcination, reduction and stabilization are simultaneously carried out in a vertical fluidized bed reactor. In most countries, sinterable grade UO2 powder for nuclear fuel is obtained by the ammonium diuranate (ADU) process, which requires several more steps. Ammonium uranyl carbonate is also one of the many forms called yellowcake; in this case it is the product obtained by the heap leach process.

Uranium pentafluoride Uranium pentafluoride

CAS number ChemSpider Molecular formula Molar mass

Identifiers 10049-14-6 14676182 Properties UF5 333.02 g/mol

Uranium pentafluoride is a coordination polymer which consists of UF5 units linked by bridging fluorides forming linear chains. The pentafluoride of uranium has been characterised by C.J. Howard, J.C Taylor and A.B. Waugh.

Uranyl chloride Uranyl chloride IUPAC name Dichlorodioxouranium Other names Uranium(VI), dichlorodioxy Properties Molecular formula UO2Cl2 Molar mass 340.90

Melting point Boiling point Solubility in other solvents Hazards MSDS

Decomposes Decomposes 320 @ 18C External MSDS

Uranyl chloride, UO2Cl2 is an unstable, bright yellow coloured chemical compound of uranium. It forms large sand-like crystals which are highly soluble in water, alcohols and ethers. Uranyl chloride, and its two hydrates (UO2Cl2·H2O and UO2Cl2·3H2O) decomposes in the presence of light, a fact discovered by Adolph Gehlen in 1804, This photosensitivity periodically attracted scientific curiosity and various unsuccessful attempts to develop photographic applications using the salts. As with most other uranic species this compound also exhibits fluorescence. Uranyl chloride is formed when chlorine gas is passed over uranium dioxide at a red heat. However it is more usually obtained by dissolving uranium oxide in hydrochloric acid and evaporating.

Industrial importance The company Indian Rare Earths Limited (IREL) has developed a process to extract uranium from the Western and Eastern coastal dune sands of India. After pre-processing with high intensity magnetic separators and fine grinding, the mineral sands (known as monazite), are digested with caustic soda at about 120C and water. The hydroxide concentrate is further digested with concentrated hydrochloric acid to solubilise all hydroxides to form a feed solution composed of chlorides of uranium and other rare earth elements including thorium. The solution is subjected to solvent extraction with dual solvent systems to produce uranyl chloride and thorium oxalate. The crude uranyl chloride solution is subsequently refined to nuclear grade ammonium diuranate by a purification process involving precipitation and solvent extraction in a nitrate media.

Health and environmental Uranyl chloride is spectacularly toxic by inhalation and if swallowed. There is also a danger of cumulative effects. The target organs are the liver and kidneys. It is toxic to aquatic organisms, and may cause long-term adverse effects in the aquatic environment. As with all compounds of uranium it is radioactive to a degree dependent on its isotopic ratios.

Uranyl fluoride Uranyl fluoride

IUPAC name Uranium fluoride oxide Other names Uranium oxyfluoride Identifiers CAS number 13536-84-0 Properties Molecular formula UO2F2 Molar mass 308.02 g/mol Melting point Decomposes @ 300°C Boiling point Sublimes Solubility in other solvents VS

Uranyl fluoride (UO2F2), a compound of uranium, is an intermediate in the conversion of uranium hexafluoride UF6 to an uranium oxide or metal form and is a direct product of the reaction of UF6 with moisture in the air. It is very soluble in water. Uranyl fluoride also is hygroscopic and changes in color from brilliant orange to yellow after reacting with water. Uranyl fluoride is reported to be stable in air to 300°C, above which slow decomposition to U3O8 occurs. When heated to decomposition, UO2F2 emits toxic fluoride fumes. In accidental releases of UF6, UO2F2, as a solid particulate compound, may deposit on the ground. The overall chemical reaction of this event can be represented as: UF6+ 2H2O → UO2F2+ 4HF. These reactions can take place whether the uranium hexafluoride is a solid or a gas, but will take place almost instantaneously when the UF6 is in a gaseous state. The resulting hydrofluoric acid and the presence of additional water results in formation of solids (primarily Hydrofluoric adducts of hydrated uranyl fluoride (UO2F2-nH2O).

Toxicology Chemical hazards are far more significant than radioactive hazards, though there is a radioactivity concern if prepared with enriched uranium. Material is corrosive, and

harmful by inhalation, ingestion or skin absorption. Ingestion or inhalation may be fatal. Effects of exposure may be delayed.

Uranyl nitrate Uranyl nitrate

IUPAC name (T-4)-bis(nitrato-κO)dioxouranium Other names Uranium nitrate Identifiers CAS number 10102-06-4 ChemSpider 22177973 Properties Molecular formula UO2(NO3)2 Molar mass 394.04 g/mol yellow-green solid Appearance hygroscopic Density 2.81 g/cm3 Melting point 60 °C Boiling point 118 °C decomp. Solubility in water ~66 g/100 mL Solubility in tributyl soluble phosphate Hazards MSDS External MSDS EU Index 092-002-00-3 EU classification Very toxic (T+)

R-phrases S-phrases Flash point Other anions

Dangerous for the environment (N) R26/28, R33, R51/53 (S1/2), S20/21, S45, S61 Non-flammable Related compounds Uranyl chloride Uranyl sulfate

Uranyl nitrate (UO2(NO3)2) is a water soluble yellow uranium salt. The yellow-green crystals of uranium nitrate hexahydrate are triboluminescent. Uranyl nitrate can be prepared by reaction of uranium salts with nitric acid. It is soluble in water, ethanol, acetone, and ether, but not in benzene, toluene, or chloroform.

Uses During the first half of the 19th century, many photosensitive metal salts had been identified as candidates for photographic processes, among them uranyl nitrate. The prints thus produced were alternately referred to as uranium prints, urbanities, or more commonly uranotypes. The first uranium printing processes were invented by a Scotsman, J. Charles Burnett, between 1855 and 1857, and used this compound as the sensitive salt. Burnett, authored an 1858 article comparing "Printing by the Salts of the Uranic and Ferric Oxides" The basis for the process lies in the ability of the uranyl ion to pick up two electrons and reduce to the lower oxidation state of uranium(IV) under ultraviolet light. Uranotypes can vary from print to print from a more neutral, brown russet to strong Bartolozzi red, with a very long tone grade. Surviving prints are slightly radioactive, a property which serves as a means of non-destructively identifying them. Several other more elaborate photographic processes employing the compound sprung up and vanished throughout the second half of the century with names like Wothlytype, Mercuro-Uranotype and the Auro-Uranium process. Uranium papers were manufactured commercially at least until the end of the 19th century, vanishing in the face of the superior sensitivity and practical advantages of the silver halides. Nevertheless between the 1930s through the 1950s Kodak Books still described a uranium toner (Kodak T-9) using uranium nitrate hexahydrate. Some alternative process photographers including artists Blake Ferris & Robert Schramm continue to make uranotype prints today. Along with uranyl acetate it is used as a negative stain for viruses in electron microscopy; in tissue samples it stabilizes nucleic acids and cell membranes. Uranyl nitrate was used to fuel Aqueous Homogeneous Reactors in the 1950s. However it proved too corrosive in this application, and the experiments were abandoned. Uranyl nitrate is important for nuclear reprocessing; it is the compound of uranium that results from dissolving the decladded spent nuclear fuel rods or yellowcake in nitric acid, for further separation and preparation of uranium hexafluoride for isotope separation for preparing of enriched uranium.

Health and environmental issues Uranyl nitrate is an oxidizing and highly toxic compound and should not be ingested; it causes severe renal insufficiency and acute tubular necrosis and is a lymphocyte mitogen. Target organs include the kidneys, liver, lungs and brain. It also represents a severe fire and explosion risk when heated or subjected to shock in contact with oxidizable substances.

Uranyl hydroxide Uranyl hydroxide

Molecular formula Molar mass Density

Properties H2O4U 304.04 g mol−1 5.73 - 6.73 g/cc

Uranyl hydroxide is a hydroxide of uranium with the chemical formula UO2(OH)2 in the monomeric form and (UO2)2(OH)4 in the dimeric; both forms may exist in normal aqueous media. Uranyl hydroxide hydrate is precipitated as a colloidal yellowcake from oxidized uranium liquors near neutral pH. Uranyl hydroxide was once used in glassmaking and ceramics in the colouring of the vitreous phases and the preparation of pigments for high temperature firing. The introduction of alkaline diuranates into glasses leads to yellow by transmission, green by reflection; moreover these glasses become dichroic and fluorescent under ultraviolet rays. Uranyl hydroxide is teratogenic and radioactive, and should be handled with the appropriate care.

Uranyl sulfate Uranyl sulfate

Properties UO2SO4 366.09 g/mol 3,28 g/cm3 @ 20 °C 27,5 g/100 mL in water at 25 °C Related compounds Uranyl chloride Other anions Uranyl nitrate Uranyl carbonate Related compounds Uranium dioxide

Molecular formula Molar mass Density Solubility in water

Uranyl sulfate (UO2SO4) a sulfate of uranium is an odorless lemon-yellow sand-like solid in its pure crystalline form. It has found use as a negative stain in microscopy and tracer in biology. The Aqueous Homogeneous Reactor experiment, constructed in 1951, circulated a fuel composed of 565 grams of U-235 enriched to 14.7% in the form of uranyl sulfate. The acid process of milling uranium ores involves precipitating uranyl sulfate from the pregnant leaching solution to produce the semi-refined product referred to as yellowcake. Radioactivity was discovered using uranyl double sulfate, K2UO2(SO4)2.

Uranium sulfate Uranium sulfate ChemSpider Molecular formula Molar mass

Identifiers 11383849 Properties U(SO4)2 430.15 g/mol

Uranium sulfate (U(SO4)2) is a water soluble salt of uranium. It is a very toxic compound and should not be ingested. Uranium sulfate minerals commonly are

widespread around uranium bearing mine sites, where they usually form during the evaporation of acid sulfate-rich mine tailings which have been leached by oxygenbearing waters. Uranium sulfate is a transitional compound in the production of Uranium hexafluoride. It was also used to fuel aqueous Homogeneous Reactors

Preparation Uranyl sulfate in solution is readily photochemically reduced to uranium(IV) sulfate. The photoreduction can be carried out in the sunshine and this requires the addition of ethanol as a reducing agent. Uranium(IV) crystallizes or is precipitated by ethanol in excess. It can be obtained with different degrees of hydration.

Uranyl zinc acetate Uranyl zinc acetate IUPAC name zinc bis(acetato-O)dioxouranate Other names zinc uranyl acetate Identifiers CAS number 10138-94-0 Properties Molecular formula ZnUO2(CH3COO)4 Molar mass 571.59 g/mol

Uranyl zinc acetate (ZnUO2(CH3COO)4) is a compound of uranium. Uranyl zinc acetate is used as a laboratory reagent in the determination of sodium concentrations of solutions using a method of quantitatively precipitating sodium with uranyl zinc acetate and gravimetrically determining the sodium as uranyl zinc sodium acetate, (UO2)2ZnNa(CH3COO)-6H2O. This method was important to determine Na in urine for diagnostic purposes. Zinc uranyl acetate is sometimes called "sodium reagent" since pale yellow NaZn(UO2)3(C2H3O2)9 is one of the very few insoluble sodium compounds. The process for catalytic synthesis of toluene-2,4-diisocyanate (TDI) from dimethyl carbonate (DMC) consists of two steps. Starting from the catalytic reaction between toluene-2,4-diamine (TDA) and DMC, dimethyl toluene-2,4-dicarbamate (TDC) is formed, and then decomposed to TDI. For the first step, the yield of TDC is 53.5% at a temperature of 250 °C, over Zn(OAc)2/alpha–Al2O3 catalyst. For the second step, the yield of TDI is 92.6% at temperatures of 250–270 °C and under pressure of 2.7 kPa, over uranyl zinc acetate catalyst, when di-n-octyl sebacate(DOS) is used as heat-carrier, and a mixture of tetrahydrofuran (THF) and nitrobenzene is used as solvent.

Chapter-2

Uranium Hexafluoride and Uranium Hydride

Uranium hexafluoride Uranium fluoride

IUPAC name Uranium hexafluoride Uranium(VI) fluoride Identifiers CAS number 7783-81-5 PubChem 24560 ChemSpider 22966 2978 (<1% 235U) UN number 2977 (>1% 235U) RTECS number YR4720000 Properties Molecular formula UF6 Molar mass 352.02 g/mol Appearance colorless solid Density 5.09 g/cm3, solid Melting point 64.05 °C (triple point) Boiling point 56.5 °C (sublimes) Solubility in water reacts soluble in chloroform, CCl4, liquid Solubility chlorine and bromine

Crystal structure Space group Coordination geometry Dipole moment Std enthalpy of formation ΔfHo298 Standard molar entropy So298 MSDS EU Index EU classification R-phrases S-phrases Flash point Other anions Other cations Related uranium fluorides

dissolves in nitrobenzene Structure Orthorhombic, oP28 Pnma, No. 62 octahedral (Oh) 0 Thermochemistry – solid: −(2197,7 ± 1,8) kJ·mol−1 – gaseous: −(2148,1 ± 1,8) kJ·mol−1 – solid: −430,4 ± 1,5 J·K−1·mol−1 – gaseous: −280,4 ± 1,5 J·K−1·mol−1 Hazards ICSC 1250 092-002-00-3 Very toxic (T+) Dangerous for the environment (N) R26/28, R33, R51/53 (S1/2), S20/21, S45, S61 Non-flammable Related compounds Uranium hexachloride Neptunium hexafluoride Plutonium hexafluoride Uranium(III) fluoride Uranium(IV) fluoride Uranium(V) fluoride

Uranium hexafluoride (UF6), referred to as "hex" in the nuclear industry, is a compound used in the uranium enrichment process that produces fuel for nuclear reactors and nuclear weapons. It forms solid grey crystals at standard temperature and pressure (STP), is highly toxic, reacts violently with water and is corrosive to most metals. It reacts mildly with aluminium, forming a thin surface layer of AlF3 that resists further reaction.

Preparation Milled uranium ore—U3O8 or "yellowcake"—is dissolved in nitric acid, yielding a solution of uranyl nitrate UO2(NO3)2. Pure uranyl nitrate is obtained by solvent extraction, then treated with ammonia to produce ammonium diuranate ("ADU", (NH4)2U2O7). Reduction with hydrogen gives UO2, which is converted with hydrofluoric acid (HF) to uranium tetrafluoride, UF4. Oxidation with fluorine yields UF6.

Properties

UF6 in a glass ampoule.

Physical Properties At room pressure, it sublimes at 56.5 °C. The triple point is at 64.05 °C and 1.5 bar. The solid state structure was determined by neutron diffraction at 77 K and 293 K.

This is a simple mononuclear molecule

The crystal structure of uranium hexafluoride

Chemical Properties It has been shown that uranium hexafluoride is an oxidant and a Lewis acid which is able to bind to fluoride, for instance the reaction of copper fluoride with uranium hexafluoride in acetonitrile is reported to form copper(II) heptafluorouranate(VI), F7Cu3U. Polymeric uranium(VI) fluorides containing organic cations have been isolated and characterised by X-ray diffraction.

Application in the nuclear fuel cycle

Phase diagram of UF6. UF6 is used in both of the main uranium enrichment methods, gaseous diffusion and the gas centrifuge method, because it has a triple point at 64.05 °C (147 °F, 337 K) and slightly higher than normal atmospheric pressure. Fluorine has only a single stable naturally occurring isotope, so isotopologues of UF6 differ in their molecular weight based solely on the uranium isotope present. All the other uranium fluorides are involatile solids which are coordination polymers. Gaseous diffusion requires about 60 times as much energy as the gas centrifuge process; even so, this is just 4% of the energy that can be produced by the resulting enriched uranium. In addition to its use in enrichment, uranium hexafluoride has been used in an advanced reprocessing method (fluoride volatility) which was developed in the Czech Republic. In this process, used oxide nuclear fuel is treated with fluorine gas to form a mixture of fluorides. This is then distilled to separate the different classes of material.

Storage in gas cylinders

UF6-cylinder

Ruptured 14-ton UF6 shipping cylinder. 1 fatality, dozens injured. ~29500 lbs of material released. 1986

About 95% of the depleted uranium produced to date is stored as uranium hexafluoride, DUF6, in steel cylinders in open air yards close to enrichment plants. Each cylinder contains up to 12.7 tonnes (or 14 US tons) of solid UF6. In the U.S. alone, 560,000 tonnes of depleted UF6 had accumulated by 1993. In 2005, 686,500 tonnes in 57,122 storage cylinders were located near Portsmouth, Ohio, Oak Ridge, Tennessee, and Paducah, Kentucky. The long-term storage of DUF6 presents environmental, health, and safety risks because of its chemical instability. When UF6 is exposed to moist air, it reacts with the water in the air to produce UO2F2 (uranyl fluoride) and HF (hydrogen fluoride) both of which are highly soluble and toxic. Storage cylinders must be regularly inspected for signs of corrosion and leaks. The estimated lifetime of the steel cylinders is measured in decades. There have been several accidents involving uranium hexafluoride in the United States. The U.S. government has been converting DUF6 to solid uranium oxides for disposal. Such disposal of the entire DUF6 inventory could cost anywhere from $15 million to $450 million.

Uranium hydride Uranium hydride Other names Uranium trihydride Uranium(III) hydride Identifiers CAS number 7440-61-1 , 13598-56-6 Properties Molecular formula UH3 Molar mass 241.05 Density 10.95 g/cm3 Melting point °C Solubility in nitric acid decomposes Structure Crystal structure Cubic, cP32 Space group Pm3n, No. 223 Lattice constant a = 664.3 pm Hazards MSDS MSDS Flash point Pyrophoric

Uranium hydride, also called uranium trihydride (UH3) is an inorganic compound, a hydride of uranium. It is a highly toxic, brownish gray to brownish black pyrophoric powder or brittle solid. It is electrically conductive. Its specific gravity at 20 °C is 10.95, much lower than that of uranium. It is slightly soluble in hydrochloric acid and decomposes in nitric acid.

Uranium metal heated to 250 to 300 °C (482 to 572 °F) reacts with hydrogen to form uranium hydride. Even higher temperatures will reversibly remove the hydrogen. This property makes uranium hydrides convenient starting materials to create reactive uranium powder along with various uranium carbide, nitride, and halide compounds. Two crystal modifications of uranium hydride exist: an α form that is obtained at low temperatures and a β form that is created when the formation temperature is above 250 °C. Uranium hydride expands considerably during formation and is therefore not an interstitial compound. In its lattice, each uranium atom is surrounded by 6 other uranium atoms and 12 atoms of hydrogen; each hydrogen atom occupies a large tetrahedral hole in the lattice. The density of hydrogen in uranium hydride is approximately the same as in liquid water or in liquid hydrogen. The U-H-U linkage through a hydrogen atom is present in the structure. Hydrogen, deuterium, and tritium can be purified by reacting with uranium, then thermally decomposing the resulting hydride/deuteride/tritide. Extremely pure hydrogen has been prepared from beds of uranium hydride for decades. Heating uranium hydride is a convenient way to introduce hydrogen into a vacuum system. The swelling and pulverization at uranium hydride synthesis can be used for preparation of very fine uranium metal if the powdered hydride is thermally decomposed. Uranium hydride can be used for isotope separation of hydrogen, preparing uranium metal powder, and as a reducing agent. Uranium hydride forms when uranium metal in e.g. Magnox fuel with corroded cladding gets exposed to water; the reaction proceeds as follows: 7 U + 6 H2O → 3 UO2 + 4 UH3 The formed uranium hydride is pyrophoric; when the metal (e.g. a damaged fuel rod) gets exposed to air afterwards, a lot of heat may be generated and the bulk uranium metal itself can be ignited. Hydride-contaminated uranium can be passivated by exposition to a gaseous mixture of 98% helium with 2% oxygen. Condensed moisture on uranium metal promotes formation of hydrogen and uranium hydride; pyrophoric surface may be formed in absence of oxygen. This poses a problem with underwater storage of spent nuclear fuel in spent fuel ponds. Depending on the size and distribution on the hydride particles, selfignition occurs after indeterminate length of exposure to air. Such exposure poses risk of self-ignition of fuel debris in radioactive waste storage vaults. Uranium metal exposed to steam produces a mixture of uranium hydride and uranium dioxide. Exposition of uranium metal to hydrogen leads to hydrogen embrittlement. Hydrogen diffuses through metal and forms a network of brittle hydride over the grain boundaries. Hydrogen can be removed and ductility renewed by annealing in vacuum.

Uranium hydride exposed to water evolves hydrogen. In contact with strong oxidizers this may cause fire and explosions. Contact with halocarbons may cause a violent reaction. Polystyrene-impregnated uranium hydride powder is non-pyrophoric and can be pressed, however its hydrogen-carbon ratio is unfavorable. Hydrogenated polystyrene was introduced in 1944 instead. Uranium hydride slugs were used in the Tickling the Dragons Tail series of experiments to determine the critical mass of uranium. Uranium hydride and uranium deuteride were suggested as a fissile material for an uranium hydride bomb. The tests with uranium hydride and uranium deuteride during the Operation Upshot-Knothole were however disappointing. During early phases of Manhattan Project, in 1943, uranium hydride was investigated as a promising bomb material; it was however abandoned by spring 1944 as it turned out such design would be inefficient. Uranium deuteride is said to be usable for design of some types of neutron initiators. Uranium hydride enriched to about 10% uranium-235 is proposed as a combined nuclear fuel/neutron moderator for the Hydrogen Moderated Self-regulating Nuclear Power Module. According to the aforementioned patent application, the reactor design in question begins producing power when hydrogen gas at a sufficient temperature and pressure is admitted to the core (made up of granulated uranium metal) and reacts with the uranium metal to form uranium hydride. Uranium hydride is both a nuclear fuel and a neutron moderator; apparently it, like other neutron moderators, will slow neutrons sufficiently to allow for fission reactions to take place; the U-235 atoms within the hydride also serve as the nuclear fuel. Once the nuclear reaction has started, it will continue until it reaches a certain temperature, approximately 800 °C (1,500 °F), where, due to the chemical properties of uranium hydride, it chemically decomposes and turns into hydrogen gas and uranium metal. The loss of neutron moderation due to the chemical decomposition of the uranium hydride will consequently slow — and eventually halt — the reaction. When temperature returns to an acceptable level, the hydrogen will again combine with the uranium metal, forming uranium hydride, restoring moderation and the nuclear reaction will start again. Uranium zirconium hydride (UZrH), a combination of uranium hydride and zirconium(II) hydride, is used as a fuel/moderator in the TRIGA-class reactors. On heating with diborane, uranium hydride produces uranium boride. With bromine at 300 °C, uranium(IV) bromide is produced. With chlorine at 250 °C, uranium(IV) chloride is produced. Hydrogen fluoride at 20 °C produces uranium(IV) fluoride. Hydrogen chloride at 300 °C produces uranium(III) chloride. Hydrogen bromide at 300 °C produces uranium(III) bromide. Hydrogen iodide at 300 °C produces uranium(III) iodide. Ammonia at 250 °C produces uranium(III) nitride. Hydrogen sulfide at 400 °C produces

uranium(IV) sulfide. Oxygen at 20 °C produces triuranium octoxide. Water at 350 °C produces uranium dioxide. Uranium hydride ion may interfere with some mass spectrometry measurements, appearing as a peak at mass 239, creating false increase of signal for plutonium-239.

Chapter-3

Uranium Dioxide

Uranium dioxide

CAS number RTECS number Molecular formula Molar mass Appearance Density Melting point Solubility in water Crystal structure Space group Coordination geometry

IUPAC name Uranium dioxide Uranium(IV) oxide Other names Urania Uranous oxide Identifiers 1344-57-6 YR4705000 Properties UO2 270.03 g/mol black powder 10.97 g/cm3 2865 °C (3140 K) insoluble Structure Fluorite (cubic), cF12 Fm3m, No. 225 Tetrahedral (O2–); cubic (UIV) Hazards

MSDS EU Index

ICSC 1251 092-002-00-3 Very toxic (T+) EU classification Dangerous for the environment (N) R-phrases R26/28, R33, R51/53 S-phrases (S1/2), S20/21, S45, S61 Flash point Non-flammable Related compounds Triuranium octoxide Related uranium oxides Uranium trioxide

Uranium dioxide or uranium(IV) oxide (UO2), also known as urania or uranous oxide, is an oxide of uranium, and is a black, radioactive, crystalline powder that naturally occurs in the mineral uraninite. It is used in nuclear fuel rods in nuclear reactors. A mixture of uranium and plutonium dioxides is used as MOX fuel. Prior to 1960 it was used as yellow and black color in ceramic glazes and glass.

Production Uranium dioxide is produced by reducing uranium trioxide with hydrogen. UO3 + H2 → UO2 + H2O at 700 °C (970 K) This reaction takes part in the reprocessing of nuclear fuel and enrichment of uranium for nuclear fuel.

Chemistry Structure The solid is isostructural with (has the same structure as) fluorite (calcium fluoride). In addition, the dioxides of plutonium and neptunium have the same structures.

Oxidation Uranium dioxide is oxidized in contact with oxygen to the triuranium octaoxide. 3UO2 + O2 → U3O8 at 700 °C (970 K)

Aqueous electrochemistry The electrochemistry of uranium dioxide has been investigated in detail as the galvanic corrosion of uranium dioxide controls the rate at which used nuclear fuel dissolves.

Oxidation of uranium metal It has been reported that water causes the rate of the oxidation of both plutonium and uranium metal to increase when compared with the situation which exists when water is absent.

Uses

Uranium oxide fuel pellet

Nuclear Fuel UO2 is used mainly as nuclear fuel, specifically as UO2 or as a mixture of UO2 and PuO2 (plutonium dioxide) called a mixed oxide (MOX fuel) for fuel rods in nuclear reactors. Note that the thermal conductivity of uranium dioxide is very low when compared with uranium, uranium nitride, uranium carbide and zirconium cladding material. This low thermal conductivity can result in localised overheating in the centres of fuel pellets. The graph below shows the different temperature gradients in different fuel compounds. For these fuels the thermal power density is the same and the diameter of all the pellets are the same.

The thermal conductivity of zirconium metal and uranium dioxide as a function of temperature

This is a 20 mm diameter fuel pellet, note that the central temperature is very different for the different fuel solids, also for the different pellets it has the lowest centre line temperature, power density is 250 W per cubic meter and rim temperature of 200 °C

Color for ceramics glaze All uranium oxides were used to color glass and ceramics. Uranium oxide-based ceramics become green or black when fired in a reducing atmosphere and yellow to orange when fired with oxygen. Orange-colored Fiestaware is a well-known example of a product with a uranium-based glaze. Uranium oxide has also been used in formulations of enamel, uranium glass, and porcelain. Prior to 1960, uranium oxides were used as colored glazes. It is possible to determine with a Geiger counter if a glaze or glass contains uranium oxides.

Other use Depleted UO2 (DUO2) can be used as a material for radiation shielding. For example, DUCRETE is a "heavy concrete" material where gravel is replaced with uranium dioxide aggregate; this material is investigated for use for casks for radioactive waste. Casks can be also made of DUO2-steel cermet, a composite material made of an aggregate of uranium dioxide serving as radiation shielding, graphite and/or silicon carbide serving as

neutron radiation absorber and moderator, and steel as the matrix, whose high thermal conductivity allows easy removal of decay heat. Depleted uranium dioxide can be also used as a catalyst, e.g. for degradation of volatile organic compounds in gaseous phase, oxidation of methane to methanol, and removal of sulfur from petroleum. It has high efficiency and long-term stability when used to destroy VOCs when compared with some of the commercial catalysts, such as precious metals, TiO2, and Co3O4 catalysts. Much research is being done in this area, DU being favoured for the uranium component due to its low radioactivity. Use of uranium dioxide as a material for rechargeable batteries is investigated. The batteries could have high power density and potential of 4.7V per cell. Another investigated application is in photoelectrochemical cells, for solar-assisted hydrogen production. UO2 is used as a photoanode. In earlier times it was also used as heat conductor for current limitation ( URDOXresistor), which was the first use of its semiconductor properties.

Semiconductor properties Uranium dioxide is a semiconductor material. Its band gap is about 1.3 eV, which lies between the band gap for silicon and gallium arsenide, near the optimum for efficiency vs band gap curve for absorption of solar radiation, suggesting its possible use for very efficient solar cells based on Schottky diode structure; it also absorbs at five different wavelengths, including infrared, further enhancing its efficiency. Its intrinsic conductivity at room temperature is about the same as of single crystal silicon. Its dielectric constant is about 22, which is almost twice as high as of silicon (11.2) and GaAs (14.1), which poses an advantage over Si and GaAs for construction of integrated circuits, as it may allow higher density integration with higher breakdown voltages and with lower susceptibility to the CMOS tunneling breakdown. The Seebeck coefficient of uranium dioxide at room temperature is about 750 µV/K, a value significantly higher than the 270 µV/K of thallium tin telluride (Tl2SnTe5) and thallium germanium telluride (Tl2GeTe5) and of bismuth-tellurium alloys, other materials promising for thermopower applications and Peltier elements. The radioactive decay impact of the235U and238U on its semiconducting properties was not measured as of 2005. Due to the slow decay rate of these isotopes, it should not meaningfully influence the properties of uranium dioxide solar cells and thermoelectric devices, but it may become an important factor for VLSI chips. Use of depleted uranium oxide is necessary for this reason. The capture of alpha particles emitted during radioactive decay as helium atoms in the crystal lattice may also cause gradual long-term changes in its properties.

The stoichiometry of the material dramatically influences its electrical properties. For example, the electrical conductivity of UO1.994 is orders of magnitude lower at higher temperatures than the conductivity of UO2.001. Uranium dioxide, like U3O8, is a ceramic material capable of withstanding high temperatures (about 2300 °C, in comparison with at most 200 °C for silicon or GaAs), making it suitable for high-temperature applications like thermophotovoltaic devices. Uranium dioxide is also resistant to radiation damage, making it useful for rad-hard devices for special military and aerospace applications. A Schottky diode of U3O8 and a p-n-p transistor of UO2 were successfully manufactured in a laboratory.

Toxicity Uranium dioxide is known to be absorbed by phagocytosis in the lungs.

Chapter-4

Uranium Trioxide

Uranium trioxide

CAS number Molecular formula Molar mass Appearance Density Melting point Solubility in water

Space group MSDS EU Index EU classification R-phrases S-phrases Flash point

IUPAC name Uranium trioxide Uranium(VI) oxide Other names Uranyl oxide Uranic oxide Identifiers 1344-58-7 Properties UO3 286.29 g/mol yellow-orange powder 5.5–8.7 g/cm3 ~200–650 °C (decomposes) Partially soluble Structure I41/amd (γ-UO3) Hazards External MSDS 092-002-00-3 Very toxic (T+) Dangerous for the environment (N) R26/28, R33, R51/53 (S1/2), S20/21, S45, S61 Non-flammable

Related compounds Uranium dioxide Related uranium oxides Triuranium octoxide

Uranium trioxide (UO3), also called uranyl oxide, uranium(VI) oxide, and uranic oxide, is the hexavalent oxide of uranium. The solid may be obtained by heating uranyl nitrate to 400 °C. Its most commonly encountered polymorph, γ-UO3, is a yellow-orange powder.

Production and use There are three methods to generate uranium trioxide. As noted below, two are used industrially in the reprocessing of nuclear fuel and uranium enrichment.

1. U3O8 can be oxidized at 500°C with oxygen. Note that above 750 °C even in 5 atm O2 UO3 decomposes into U3O8. 2. Uranyl nitrate, UO2(NO3)2·6H2O can be heated to yield UO3. This occurs during the reprocessing of nuclear fuel. Fuel rods are dissolved in HNO3 to separate uranyl nitrate from plutonium and the fission products (the PUREX method). The pure uranyl nitrate is converted to solid UO3 by heating at 400 °C. After reduction with hydrogen (with other inert gas present) to uranium dioxide, the uranium can be used in new MOX fuel rods. 3. Ammonium diuranate or sodium diuranate (Na2U2O7·6H2O) may be decomposed. Sodium diuranate, also known as yellowcake, is converted to uranium trioxide in the enrichment of uranium. Uranium dioxide and uranium tetrafluoride are intermediates in the process which ends in uranium hexafluoride. Uranium trioxide is shipped between processing facilities in the form of a gel.

Cameco Corporation, which operates at the world's largest uranium refinery at Blind River, Ontario, produces high-purity uranium trioxide. It has been reported that the corrosion of uranium in a silica rich aqueous solution forms both uranium dioxide and uranium trioxide. In pure water, schoepite (UO2)8O2(OH)12·12(H2O) is formed in the first week and then after four months studtite (UO2)O2·4(H2O) was produced. Reports on the corrosion of uranium metal have been published by the Royal Society.

Health and safety hazards Like all hexavalent uranium compounds, UO3 is hazardous by inhalation, ingestion, and through skin contact. It is a poisonous, slightly radioactive substance, which may cause shortness of breath, coughing, acute arterial lesions, and changes in the chromosomes of white blood cells and gonads leading to congenital malformations if inhaled. However, once ingested, uranium is mainly toxic for the kidneys and may severely affect their function.

Structure Solid state structure The only well characterized binary trioxide of any actinide is UO3, of which several polymorphs are known. Solid UO3 loses O2 on heating to give green-colored U3O8: reports of the decomposition temperature in air vary from 200–650 °C. Heating at 700 °C under H2 gives dark brown uranium dioxide (UO2), which is used in MOX nuclear fuel rods.

Alpha Hydrated uranyl peroxide formed by the addition of hydrogen peroxide to an aqueous The α (alpha) solution of uranyl nitrate when heated to 200form: a layered 225 °C forms an amorphous uranium trioxide solid where the which on heating to 400-450 °C will form 2D layers are alpha-uranium trioxide. It has been stated that linked by oxygen the presence of nitrate will lower the atoms (shown in temperature at which the exothermic change red) from the amorphous form to the alpha form occurs.

Beta

β (beta) UO3. This solid has a structure which defeats most attempts to describe it.

This form can be formed by heating ammonium diuranate, while P.C. Debets and B.O. Loopstra, found four solid phases in the UO3-H2O-NH3 system that they could all be considered as being UO2(OH)2.H2O where some of the water has been replaced with ammonia. No matter what the exact stoichiometry or structure, it was found that calcination at 500°C in air forms the beta form of uranium trioxide.

Gamma The most frequently encountered polymorph is γ-UO3, whose x-ray structure has been solved from powder diffraction data. The compound crystallizes in the space group I41/amd with two uranium atoms in the The γ (gamma) asymmetric unit. Both are surrounded form, with the by somewhat distorted octahedra of different uranium oxygen atoms. One uranium atom has environments in two closer and four more distant green and yellow oxygen atoms whereas the other has four close and two more distant oxygen atoms as neighbors. Thus it is not incorrect to describe the structure as [UO2]2+[UO4]2- , that is uranyl uranate.

The environment of the uranium atoms shown as yellow in the gamma form

The chains of U2O2 rings in the gamma form in layers, alternate layers running at 90 degrees to each other. These chains are shown as containing the yellow uranium atoms, in an octahedral environment which are

distorted towards square planar by an elongation of the axial oxygenuranium bonds.

Delta

The delta (δ) form is a cubic solid where the oxygen atoms are arranged between the uranium atoms.

High pressure form There is a high-pressure solid form with U2O2 and U3O3 rings in it.

Hydrates

Hydrous and anhydrous forms of UO3

Several hydrates of uranium trioxide are known, e.g., UO3·6H2O.

Bond valence parameters It is possible by bond valence calculations to estimate how great a contribution a given oxygen atom is making to the assumed valence of uranium. Bond valence calculations use parameters which are estimated after examining a large number of crystal structures of uranium oxides (and related uranium compounds), note that the oxidation states which this method provides are only a guide which assists in the understanding of a crystal structure. The formula to use is

The sum of the s values is equal to the oxidation state of the metal centre. For uranium binding to oxygen the constants RO and B are tabulated in the table below. For each oxidation state use the parameters from the table shown below. Oxidation state RO B U(VI) 2.08 Å 0.35 U(V) 2.10 Å 0.35 U(IV) 2.13 Å 0.35 It is possible to do these calculations on paper or software.

Molecular forms While uranium trioxide is encountered as a polymeric solid under ambient conditions, some work has been done on the molecular form in the gas phase, in matrix isolations studies, and computationally.

Gas phase At elevated temperatures gaseous UO3 is in equilibrium with solid U3O8 and molecular oxygen. 2 U3O8(s) + O2(g)

6 UO3(g)

With increasing temperature the equilibrium is shifted to the right. This system has been studied at temperatures between 900 °C and 2500 °C. The vapor pressure of monomeric UO3 in equilibrium with air and solid U3O8 at ambient pressure, about 10−5 mbar (1 mPa) at 980 °C, rising to 0.1 mbar (10 Pa) at 1400 °C, 0.34 mbar (34 Pa) at 2100 °C, 1.9 mbar (193 Pa) at 2300 °C, and 8.1 mbar (809 Pa) at 2500 °C.

Matrix isolation Infrared spectroscopy of molecular UO3 isolated in an argon matrix indicates a T-shaped structure (point group C2v) for the molecule. This is in contrast to the commonly encountered D3h molecular symmetry exhibited by most trioxides. From the force constants the authors deduct the U-O bond lengths to be between 1.76 and 1.79 Å (176 to 179 pm).

Computational study

The calculated geometry of molecular uranium trioxide has C2v symmetry. Calculations predict that the point group of molecular UO3 is C2v, with an axial bond length of 1.75 Å, an equatorial bond length of 1.83 Å and an angle of 161 ° between the axial oxygens. The more symmetrical D3h species is a saddle point, 49 kJ/mol above the C2v minimum. The authors invoke a second-order Jahn-Teller effect as explanation.

Reactivity Uranium trioxide reacts at 400 °C with freon-12 to form chlorine, phosgene, carbon dioxide and uranium tetrafluoride. The freon-12 can be replaced with freon-11 which forms carbon tetrachloride instead of carbon dioxide. This is a case of a hard perhalogenated freon which is normally considered to be inert being converted chemically at a moderate temperature.

2 CF2Cl2 + UO3 → UF4 + CO2 + COCl2 + Cl2 4 CF2Cl2 + UO3 → UF4 + 3 COCl2 + CCl4 + Cl2 Uranium trioxide can be dissolved in a mixture of tributyl phosphate and thenoyltrifluoroacetone in supercritical carbon dioxide, ultrasound was employed during the dissolution.

Electrochemical modification The reversible insertion of magnesium cations into the lattice of uranium trioxide by cyclic voltammetry using a graphite electrode modified with microscopic particles of the uranium oxide has been investigated. This experiment has also been done for U3O8. This is an example of electrochemistry of a solid modified electrode, the experiment which used for uranium trioxide is related to a carbon paste electrode experiment. It is also possible to reduce uranium trioxide with sodium metal to form sodium uranium oxides. It has been the case that it is possible to insert lithium into the uranium trioxide lattice by electrochemical means, this is similar to the way that some rechargeable lithium ion batteries work. In these rechargeable cells one of the electrodes is a metal oxide which contains a metal such as cobalt which can be reduced, to maintain the electroneutrality for each electron which is added to the electrode material a lithium ion enters the lattice of this oxide electrode.

Amphoterism and reactivity to form related uranium(VI) anions and cations Uranium oxide is amphoteric and reacts as acid and as a base, depending on the conditions. As an acid UO3 + H2O → UO2−4 + 2 H+ Dissolving uranium oxide in a strong base like sodium hydroxide forms the doubly negatively charged uranate anion (UO2−4). Uranates tend to concatenate, forming diuranate, U2O2−7, or other poly-uranates. Important diuranates include ammonium diuranate ((NH4)2U2O7), sodium diuranate (Na2U2O7) and magnesium diuranate (MgU2O7), which forms part of some yellowcakes. It is worth noting that uranates of the form M2UO4 do not contain UO2−4 ions, but rather flattened UO6 octahedra, containing a uranyl group and bridging oxygens. As a base UO3 + H2O → UO2+2 + 2 OH− Dissolving uranium oxide in a strong acid like sulfuric or nitric acid forms the double positive charged uranyl cation. The uranyl nitrate formed (UO2(NO3)2·6H2O) is soluble in ethers, alcohols, ketones and esters; for example, tributylphosphate. This solubility is

used to separate uranium from other elements in nuclear reprocessing, which begins with the dissolution of nuclear fuel rods in nitric acid. The uranyl nitrate is then converted to uranium trioxide by heating. From nitric acid one obtains uranyl nitrate, trans-UO2(NO3)2·2H2O, consisting of eightcoordinated uranium with two bidentate nitrato ligands and two water ligands as well as the familiar O=U=O core.

Uranium oxides in ceramics UO3-based ceramics become green or black when fired in a reducing atmosphere and yellow to orange when fired with oxygen. Orange-coloured Fiestaware is a well-known example of a product with a uranium-based glaze. UO3-has also been used in formulations of enamel, uranium glass, and porcelain. Prior to 1960, UO3 was used as an agent of crystallization in crystalline coloured glazes. It is possible to determine with a Geiger counter if a glaze or glass was made from UO3.

Chapter-5

Fluorocarbon

Perfluorohexane, a stable fluoroalkane liquid Fluorocarbons, sometimes referred to as perfluorocarbons, are organofluorine compounds that contain only carbon and fluorine bonded together in strong carbon– fluorine bonds. Fluoroalkanes that contain only single bonds are more chemically and thermally stable than alkanes. However, fluorocarbons with double bonds (fluoroalkenes) and especially triple bonds (fluoroalkynes) are more reactive than their corresponding hydrocarbons. Fluoroalkanes can serve as oil-repellant/water-repellant fluoropolymers, solvents, liquid breathing research agents, and powerful greenhouse gases. Unsaturated fluorocarbons tend to be used as reactants.

Perfluoroisobutene, a reactive and toxic fluoroalkene gas Many chemical compounds are labeled as fluorocarbons, perfluorinated, or with the prefix perfluoro- despite containing atoms other than carbon or fluorine, such as chlorofluorocarbons and perfluorinated compounds; however, these molecules are fluorocarbon derivatives, and not true fluorocarbons. Fluorocarbon derivatives share many of the properties of fluorocarbons, while also possessing new properties due to the inclusion of new atoms. For example, fluorocarbon derivatives can function as fluoropolymers, refrigerants, solvents, anesthetics, fluorosurfactants, and ozone depletors.

Usage of term The formal IUPAC definition of a fluorocarbon is a molecule consisting wholly of fluorine and carbon. However, other fluorocarbon based molecules that are not technically fluorocarbons are commonly referred to as fluorocarbons, because of similar structures and identical properties. Compounds with atoms other than carbon and fluorine are not true fluorocarbons and they are considered as fluorocarbon derivatives in a separate section below.

Properties

Perfluorodecalin, a dense clear liquid

Physical properties Fluorocarbon liquids are colorless. They have high density, up to over twice that of water, due to their high molecular weight. Low intermolecular forces give the liquids low viscosities when compared to liquids of similar boiling points. Also, low surface tension, heats of vaporization, and refractive indices are notable. They are not miscible with most organic solvents (e.g., ethanol, acetone, ethyl acetate and chloroform), but are miscible with some hydrocarbons (e.g., hexane in some cases). They have very low solubility in water, and water has a very low solubility in them (on the order of 10 ppm). The number of carbon atoms in a fluorocarbon molecule largely determines most physical properties. The greater the number of carbon atoms, the higher the boiling point, density, viscosity, surface tension, critical properties, vapor pressure and refractive index. Gas solubility decreases as carbon atoms increase, while melting point is determined by other factors as well, so is not readily predicted.

Tetrafluoromethane, a fluorocarbon gas

London dispersion force reduction As the high electronegativity of fluorine reduces the polarizability of the atom, fluorocarbons are only weakly susceptible to the fleeting dipoles that form the basis of the London dispersion force. As a result, fluorocarbons have low intramolecular

attractive forces and are lipophobic in addition to being hydrophobic/non-polar. Thus fluorocarbons find applications as oil-, water-, and stain-repellants in products such as Gore-Tex and fluoropolymer carpet coatings. The reduced participation in the London dispersion force makes the solid polytetrafluoroethylene (PTFE) slippery as it has a very low coefficient of friction. Also, the low attractive forces in fluorocarbon liquids make them compressible and gas soluble while smaller fluorocarbons are extremely volatile. There are five fluoroalkane gases; tetrafluoromethane (bp −128 °C), hexafluoroethane (bp −78.2 °C), octafluoropropane (bp −36.5 °C), perfluoro-n-butane (bp −2.2 °C) and perfluoro-iso-butane (bp −1 °C). Nearly all other fluoroalkanes are liquids with the exception of perfluorocyclohexane, which sublimes at 51 °C. As a result of the high gas solubility of fluorocarbon liquids, they have been the subject of medical research as blood carriers because of their oxygen solubility. Fluorocarbons also have low surface energies and high dielectric strengths.

The partial charges in the polarized carbon–fluorine bond

Fluoroalkane stability Fluorocarbons with only single bonds are very stable because of the strength and nature of the carbon–fluorine bond. It is called the strongest bond in organic chemistry. Its strength is a result of the electronegativity of fluorine imparting partial ionic character through partial charges on the carbon and fluorine atoms. The partial charges shorten and strengthen the bond through favorable coulombic interactions. Additionally, multiple carbon–fluorine bonds increase the strength and stability of other nearby carbon–fluorine bonds on the same geminal carbon, as the carbon has a higher positive partial charge. Furthermore, multiple carbon–fluorine bonds also strengthen the "skeletal" carbon– carbon bonds from the inductive effect. Therefore, saturated fluorocarbons are more chemically and thermally stable than their corresponding hydrocarbon counterparts.

However, fluoroalkanes are not inert. They are susceptible to reduction through the Birch reduction.

Tetrafluoroethylene, an important reactant

Fluoroalkene and fluoroalkyne reactivity When fluorocarbons are unsaturated, they are less stable and more reactive than fluoroalkanes, or comparable hydrocarbons, due to the electronegativity of fluorine. The reactivity of the simplest fluoroalkyne, difluoroacetylene, is an example of this instability; difluoroacetylene easily polymerizes. Another example is fluorofullerene, which has weaker and longer carbon–fluorine bonds than saturated fluorocarbons. It is reactive towards nucleophiles and hydrolyzes in solution. Additionally, the polymerization of the fluoroalkene tetrafluoroethylene (which results in PTFE) is more energetically favorable than that of ethylene. Unsaturated fluorocarbons have a driving force towards sp3 hybridization due to the electronegative fluorine atoms seeking a greater share of bonding electrons with reduced s character in orbitals.

One notable exception to this trend is fluorobenzene, which is stabilized by its aromaticity.

Manufacture Prior to World War II, the only known route to fluorocarbons was by direct reaction of fluorine with the hydrocarbon. This highly exothermic process was capable only of synthesising tetrafluoromethane, hexafluoroethane and octafluoropropane; larger hydrocarbons decomposed in the extreme conditions. The Manhattan project saw the need for some very robust chemicals, including a wider range of fluorocarbons, requiring new manufacturing methods. The so-called "catalytic" method involved reacting fluorine and hydrocarbon on a bed of gold-plated copper turnings, the metal removing the heat of the reaction (so not really acting as a catalyst at all), allowing larger hydrocarbons to survive the process. However, it was the Fowler process that allowed the large scale manufacture of fluorocarbons required for the Manhattan project.

The Fowler Process The Fowler process uses cobalt fluoride to moderate the reaction. In the laboratory, this is typically done in two stages, the first stage being fluorination of cobalt difluoride to cobalt trifluoride. 2 CoF2 + F2 → 2 CoF3 During the second stage, in this instance to make perfluorohexane, the hydrocarbon feed is introduced and is fluorinated by the cobalt trifluoride, which is converted back to cobalt difluoride. Both stages are performed at high temperature. C6H14 + 28 CoF3 → C6F14 + 14 HF + 28 CoF2 Industrially, both steps are combined, for example in the manufacture of the Flutec range of fluorocarbons, using a vertical stirred bed reactor, with hydrocarbon introduced at the bottom, and fluorine introduced half way up the reactor. The fluorocarbon vapor is recovered from the top.

Electrochemical Fluorination An alternative technique, electrochemical fluorination (ECF) (also known as the Simons' process) involves electrolysis of a substrate dissolved in hydrogen fluoride. As fluorine is itself manufactured by the electrolysis of hydrogen fluoride, this is a rather more direct route to fluorocarbons. The process is run at low voltage (5 - 6 V) so that free fluorine is not liberated. The choice of substrate is restricted as ideally it should be soluble in hydrogen fluoride. Ethers and tertiary amines are typically employed. To make perfluorohexane, trihexylamine is used, for example: 2 N(C6H13)3 + 90 HF → 6 C6F14 + 2 NF3 + 45 H2

The perfluorinated amine will also be produced: N(C6H13)3 + 42 HF → 2 N(C6F13)3 + 21H2 Both of these products, and others, are manufactured by 3M as part of the Fluorinert range.

Derivatives Fluorocarbon derivatives are highly fluorinated molecules that can be commonly referred to as fluorocarbons. They are economically useful because they share part or nearly all of the properties of fluorocarbons. Some fluorocarbon derivatives have markedly different properties than fluorocarbons. For example, fluorosurfactants powerfully reduce surface tension by concentrating at the liquid-air interface due to the lipophobicity of fluorocarbons, due to the polar functional group added to the fluorocarbon chain. Other groups or atoms for fluorocarbon based compounds the oxygen atom incorporated into an ether group for anesthetics, and the chlorine atom for chlorofluorocarbons (CFCs). In a sharp contrast to true fluorocarbons, the chlorine atom produces a chlorine radical which degrades ozone.

Fluorosurfactants   

Perfluorooctanesulfonic acid (PFOS) Perfluorooctanoic acid (PFOA) Perfluorononanoic acid

Anesthetics     

Methoxyflurane (contains chlorine) Enflurane (contains chlorine) Isoflurane (contains chlorine) Sevoflurane Desflurane

Halogenated derivatives    

Polychlorotrifluoroethylene ([CFClCF2]n) Perfluorooctyl bromide (Perflubron) Dichlorodifluoromethane Chlorodifluoromethane

Hydrofluorocarbons  

Polyvinylidene fluoride ([CH2CF2]n) Tetrafluoroethane

Environmental and Health Concerns Despite the presence of some natural fluorocarbons such as tetrafluoromethane, which has been reported in rocks, man-made fluorocarbons are potent greenhouse gases.

Chapter-6

Chlorofluorocarbon

A chlorofluorocarbon (CFC) is an organic compound that contains carbon, chlorine, and fluorine, produced as a volatile derivative of methane and ethane. A common subclass is the hydrochlorofluorocarbons (HCFCs), which contain hydrogen, as well. They are also commonly known by the DuPont trade name Freon. The most common representative is dichlorodifluoromethane (R-12 or Freon-12). Many CFCs have been widely used as refrigerants, propellants (in aerosol applications), and solvents. The manufacture of such compounds is being phased out by the Montreal Protocol because they contribute to ozone depletion.

Structure, properties, production As in simpler alkanes, carbon in the CFCs and the HCFCs is tetrahedral. Since the fluorine and chlorine atoms differ greatly in size from hydrogen and from each other, the methane derived CFCs deviate from perfect tetrahedral symmetry. The physical properties of the CFCs and HCFCs are tunable by changes in the number and identity of the halogen atoms. In general they are volatile, but less so than parent alkane. The decreased volatility is attributed to the molecular polarity induced by the halides and the polarizability of halides, which induces intermolecular interactions. Thus, methane boils at -161 °C whereas the fluoromethanes boil between -51.7 (CF2H2) and 128 °C (CF4). The CFCs have still higher boiling points because the chloride is even more polarizable than fluoride. Because of their polarity, the CFCs are useful solvents. The CFCs are far less flammable than methane, in part because they contain fewer C-H bonds and in part because, in the case of the chlorides and bromides, the released halides quench the free radicals that sustain flames. The densities of CFCs are invariably higher than the corresponding alkanes. In general the density of these compounds correlates with the number of chlorides.

CFCs and HCFCs are usually produced by halogen exchange starting from chlorinated methanes and ethanes. Illustrative is the synthesis of chlorodifluoromethane from chloroform: HCCl3 + 2 HF → HCF2Cl + 2 HCl The brominated derivatives are generated by free-radical reactions of the chlorofluorocarbons, replacing C-H bonds with C-Br bonds. The production of the anesthetic 2-bromo-2-chloro-1,1,1-trifluoroethane ("halothane") is illustrative: CF3CH2Cl + Br2 → CF3CHBrCl + HBr

Reactions The most important reaction of the CFCs is the photo-induced scission of a C-Cl bond: CCl3F → CCl2F. + Cl. The chlorine atom, written often as Cl., behaves very differently from the chlorine molecule (Cl2). The radical Cl. is long-lived in the upper atmosphere, where it catalyzes the conversion of ozone into O2. Ozone absorbs UV-radiation better than O2 does, so its depletion allows more of this high energy radiation to reach the Earth's surface. Bromine atoms are even more efficient catalysts, hence brominated CFCs are also regulated.

Applications Applications exploit the low toxicity, low reactivity, and low flammability of the CFCs and HCFCs. Every permutation of fluorine, chlorine, and hydrogen based on methane and ethane has been examined and most have been commercialized. Furthermore, many examples are known for higher numbers of carbon as well as related compounds containing bromine. Uses include refrigerants, blowing agents, propellants in medicinal applications, and degreasing solvents. Billions of kilograms of chlorodifluoromethane are produced annually as a precursor to tetrafluoroethylene, the monomer that is converted into Teflon.

Classes of compounds, nomenclature  



Chlorofluorocarbons (CFCs): when derived from methane and ethane these compounds have the formulae CClmF4-m and C2ClmF6-m, where m is nonzero. Hydrochlorofluorocarbons (HCFCs): when derived from methane and ethane these compounds have the formulae CClmFnH4-m-n and C2ClxFyH6-x-y, where m, n, x, and y are nonzero. Bromochlorofluorocarbons and bromofluorocarbons have formulae similar to the CFCs and HCFCs but also bromine.



Hydrofluorocarbons (HFC's): when derived from methane, ethane, propane, and butane, these compounds have the respective formulae CFmH4-m, C2FmH6-m, C3FmH8-m, and C4FmH10-m, where m is nonzero.

Commercial names Freon is DuPont's brand name for CFCs, HCFCs and related compounds. Other commercial names from around the world are Algofrene, Arcton, Asahiflon, Daiflon, Eskimo, FCC, Flon, Flugene, Forane, Fridohna, Frigen, Frigedohn, Genetron, Isceon, Isotron, Kaiser, Kaltron, Khladon, Ledon, Racon, and Ucon.

Numbering system A numbering system is used for fluorinated alkanes, prefixed with Freon-, R-, CFC-, and HCFC-. The rightmost value indicates the number of fluorine atoms, the next value to the left is the number of hydrogen atoms plus 1, and the next value to the left is the number of carbon atoms less one (zeroes are not stated). Remaining atoms are chlorine. Thus, Freon-12 indicates a methane derivative (only two numbers) containing two fluorine atoms (the second 2) and no hydrogen (1-1=0). It is therefore CCl2F2. Another, easier equation that can be applied to get the correct molecular formula of the CFC/R/Freon class compounds is this to take the numbering and add 90 to it. The resulting value will give the number of carbons as the first numeral, the second numeral gives the number of hydrogen atoms, and the third numeral gives the number of fluorine atoms. The rest of the unaccounted carbonbonds are occupied by chlorine atoms. The value of this equation is always a three figure number. An easy example is that of CFC12, which gives: 90+12=102 -> 1 carbon, 0 hydrogens, 2 fluorine atoms, and hence 2 chlorine atoms resulting in CCl2F2. The main advantage of this method of deducing the molecular composition in comparison with the method described in the paragraph above, is that it gives the number of carbon atoms of the molecule. Freons containing bromine is signified by four numbers. Isomers, which are common for ethane and propane derivatives, are indicated by letters following the numbers. Principal CFCs Boiling Systematic name point Chem. formula (°C) Trichlorofluoromethane Freon-11, R-11, CFC-11 23 CCl3F Dichlorodifluoromethane Freon-12, R-12, CFC-12 −29.8 CCl2F2 Chlorotrifluoromethane Freon-13, R-13, CFC-13 -81 CClF3 Chlorodifluoromethane R-22, HCFC-22 -40.8 CHClF2 Dichlorofluoromethane R-21, HCFC-21 8.9 CHCl2F Chlorofluoromethane Freon 31, R-31, HCFC-31 CH2ClF BCF, Halon 1211, H-1211, Bromochlorodifluoromethane CBrClF2 Freon 12B1 Common/Trivial name(s), code

1,1,2-Trichloro-1,2,2trifluoroethane 1,1,1-Trichloro-2,2,2trifluoroethane 1,2-Dichloro-1,1,2,2tetrafluoroethane 1-Chloro-1,1,2,2,2pentafluoroethane 2-Chloro-1,1,1,2tetrafluoroethane 1,1-Dichloro-1-fluoroethane 1-Chloro-1,1-difluoroethane Tetrachloro-1,2difluoroethane Tetrachloro-1,1difluoroethane 1,1,2-Trichlorotrifluoroethane 1-Bromo-2-chloro-1,1,2trifluoroethane 2-Bromo-2-chloro-1,1,1trifluoroethane 1,1-Dichloro-2,2,3,3,3pentafluoropropane 1,3-Dichloro-1,2,2,3,3pentafluoropropane

Freon 113, R-113, CFC113, 1,1,2Trichlorotrifluoroethane Freon 113a, R-113a, CFC113a Freon 114, R-114, CFC114, Dichlorotetrafluoroethane Freon 115, R-115, CFC115, Chloropentafluoroethane

47.7

Cl2FC-CClF2

45.9

Cl3C-CF3

3.8

ClF2C-CClF2

−38

ClF2C-CF3

R-124, HCFC-124

−12

CHFClCF3

R-141b, HCFC-141b R-142b, HCFC-142b

32 −9.2

Cl2FC-CH3 ClF2C-CH3

Freon 112, R-112, CFC-11291.5 Freon 112a, R-112a, CFC91.5 112a Freon 113, R-113, CFC-11348

CCl2FCCl2F CClF2CCl3 CCl2FCClF2

Halon 2311a

51.7

CHClFCBrF2

Halon 2311

50.2

CF3CHBrCl

R-225ca, HCFC-225ca

51

CF3CF2CHCl2

R-225cb, HCFC-225cb

56

CClF2CF2CHClF

History Carbon tetrachloride (CCl4) was used in fire extinguishers and glass "anti-fire grenades" from the late nineteenth century until around the end of World War II. Experimentation with chloroalkanes for fire suppression on military aircraft began at least as early as the 1920s. Freon is a trade name for a group of CFCs which are used primarily as refrigerants, but also have uses in fire-fighting and as propellants in aerosol cans. Bromomethane is widely used as a fumigant. Dichloromethane is a versatile industrial solvent. The Belgian scientist Frédéric Swarts pioneered the synthesis of CFCs in the 1890s. He developed an effective exchange agent to replace chloride in carbon tetrachloride with fluoride to synthesize CFC-11 (CCl3F) and CFC-12 (CCl2F2). In the late 1920s, Thomas Midgley, Jr. improved the process of synthesis and led the effort to use CFC as refrigerant to replace ammonia (NH3), chloromethane (CH3Cl), and sulfur dioxide (SO2), which are toxic but were in common use. In searching for a new

refrigerant, requirements for the compound were: low boiling point, low toxicity, and to be generally non-reactive. In a demonstration for the American Chemical Society, Midgley flamboyantly demonstrated all these properties by inhaling a breath of the gas and using it to blow out a candle in 1930.

Commercial development and use of CFCs and related compounds

During World War II, various chloroalkanes were in standard use in military aircraft, although these early halons suffered from excessive toxicity. Nevertheless, after the war they slowly became more common in civil aviation as well. In the 1960s, fluoroalkanes and bromofluoroalkanes became available and were quickly recognized as being highly effective fire-fighting materials. Much early research with Halon 1301 was conducted under the auspices of the US Armed Forces, while Halon 1211 was, initially, mainly developed in the UK. By the late 1960s they were standard in many applications where water and dry-powder extinguishers posed a threat of damage to the protected property, including computer rooms, telecommunications switches, laboratories, museums and art collections. Beginning with warships, in the 1970s, bromofluoroalkanes also progressively came to be associated with rapid knockdown of severe fires in confined spaces with minimal risk to personnel. By the early 1980s, bromofluoroalkanes were in common use on aircraft, ships, and large vehicles as well as in computer facilities and galleries. However, concern was beginning to be expressed about the impact of chloroalkanes and bromoalkanes on the ozone layer. The Vienna Convention for the Protection of the Ozone Layer did not cover bromofluoroalkanes as it was thought, at the time, that emergency discharge of extinguishing systems was too small in volume to produce a significant impact, and too important to human safety for restriction.

Regulation Since the late 1970s, the use of CFCs has been heavily regulated because of their destructive effects on the ozone layer. After the development of his electron capture

detector, James Lovelock was the first to detect the widespread presence of CFCs in the air, finding a mole fraction of 60 ppt of CFC-11 over Ireland. In a self-funded research expedition ending in 1973, Lovelock went on to measure CFC-11 in both the Arctic and Antarctic, finding the presence of the gas in each of 50 air samples collected, but incorrectly concluding that CFCs are not hazardous to the environment. The experiment did however provide the first useful data on the presence of CFCs in the atmosphere. The damage caused by CFCs was discovered by Sherry Rowland and Mario Molina who, after hearing a lecture on the subject of Lovelock's work, embarked on research resulting in the first publication suggesting the connection in 1974. It turns out that one of CFCs' most attractive features—their low reactivity— is key to their most destructive effects. CFCs' lack of reactivity gives them a lifespan that can exceed 100 years, giving them time to diffuse into the upper stratosphere. Once in the stratosphere, the sun's ultraviolet radiation is strong enough to cause the homolytic cleavage of the C-Cl bond. By 1987, in response to a dramatic seasonal depletion of the ozone layer over Antarctica, diplomats in Montreal forged a treaty, the Montreal Protocol, which called for drastic reductions in the production of CFCs. On March 2, 1989, 12 European Community nations agreed to ban the production of all CFCs by the end of the century. In 1990, diplomats met in London and voted to significantly strengthen the Montreal Protocol by calling for a complete elimination of CFCs by the year 2000. By the year 2010 CFCs should be completely eliminated from developing countries as well.

Ozone-depleting gas trends Because the only CFCs available to countries adhering to the treaty is from recycling, their prices have increased considerably. A worldwide end to production should also terminate the smuggling of this material. By the time of the Montreal Protocol it was realised that deliberate and accidental discharges during system tests and maintenance accounted for substantially larger volumes than emergency discharges, and consequently halons were brought into the treaty, albeit with many exceptions.

Regulatory Gap While the production and consumption of CFCs are regulated under the Montreal Protocol, emissions from pre-existing banks of CFCs are not regulated under the agreement. As of 2002, there were 5,791 kilotons of CFCs in existing products such as refrigerators, air conditioners, aerosol cans and others. Approximately one-third of these CFCs are projected to be emitted over the next decade if action is not taken, posing a threat to both the ozone layer and the climate. A proportion of these CFCs can be safely captured and destroyed.

Regulation and DuPont In 1978 the United States banned the use of CFCs such as Freon in aerosol cans, the beginning of a long series of regulatory actions against their use. The critical DuPont manufacturing patent for Freon ("Process for Fluorinating Halohydrocarbons", U.S. Patent #3258500) was set to expire in 1979. In conjunction with other industrial peers DuPont sponsored efforts such as the "Alliance for Responsible CFC Policy" to question anti-CFC science, but in a turnabout in 1986 DuPont, with new patents in hand, publicly condemned CFCs. DuPont representatives appeared before the Montreal Protocol urging that CFCs be banned worldwide and stated that their new HCFCs would meet the worldwide demand for refrigerants.

Phase out of CFCs Use of certain chloroalkanes as solvents for large scale application, such as dry cleaning, have been phased out, for example, by the IPPC directive on greenhouse gases in 1994 and by the Volatile Organic Compounds (VOC) directive of the EU in 1997. Permitted chlorofluoroalkane uses are medicinal only. Bromofluoroalkanes have been largely phased out and the possession of equipment for their use is prohibited in some countries like the Netherlands and Belgium, from 1 January 2004, based on the Montreal Protocol and guidelines of the European Union. Production of new stocks ceased in most (probably all) countries as of 1994. However many countries still require aircraft to be fitted with halon fire suppression systems because no safe and completely satisfactory alternative has been discovered for this application. There are also a few other, highly specialized uses. These programs recycle halon through "halon banks" coordinated by the Halon Recycling Corporation to ensure that discharge to the atmosphere occurs only in a genuine emergency and to conserve remaining stocks. The interim replacements for CFCs are hydrochlorofluorocarbons (HCFCs), which deplete stratospheric ozone, but to a much lesser extent than CFCs. Ultimately, hydrofluorocarbons (HFCs) will replace HCFCs with essentially no ozone destruction (although all three groups of halocarbons are powerful greenhouse gases). DuPont began producing hydrofluorocarbons as alternatives to Freon in the 1980s. These included Suva

refrigerants and Dymel propellants. Natural refrigerants are climate friendly solutions that are enjoying increasing support from large companies and governments interested in reducing global warming emissions from refrigeration and air conditioning. Hydrofluorocarbons are included in the Kyoto Protocol because of their very high Global Warming Potential and are facing calls to be regulated under the Montreal Protocol due to the recognition of halocarbon contributions to climate change. On September 21, 2007, approximately 200 countries agreed to accelerate the elimination of hydrochlorofluorocarbons entirely by 2020 in a United Nations-sponsored Montreal summit. Developing nations were given until 2030. Many nations, such as the United States and China, who had previously resisted such efforts, agreed with the accelerated phase out schedule; however, presently China and Brazil are nations notable for their large and increasing production of chlorofluorocarbons.

Development of alternatives for CFCs Work on alternatives for chlorofluorocarbons in refrigerants began in the late 1970s after the first warnings of damage to stratospheric ozone were published. The hydrochlorofluorocarbons (HCFCs) are less stable in the lower atmosphere, enabling them to break down before reaching the ozone layer. Nevertheless, a significant fraction of the HCFCs do break down in the stratosphere and they have contributed to more chlorine buildup there than originally predicted. Later alternatives lacking the chlorine, the hydrofluorocarbons (HFCs) have an even shorter lifetimes in the lower atmosphere. One of these compounds, HFC-134a, is now used in place of CFC-12 in automobile air conditioners. Hydrocarbon refrigerants (a propane/isobutane blend) are also used extensively in mobile air conditioning systems in Australia, the USA and many other countries, as they have excellent thermodynamic properties and perform particularly well in high ambient temperatures. One of the natural refrigerants (along with Ammonia and Carbon Dioxide), hydrocarbons have negligible environmental impacts and are also used worldwide in domestic and commercial refrigeration applications, and are becoming available in new split system air conditioners. Various other solvents and methods have replaced the use of CFCs in laboratory analytics. Applications and replacements for CFCs Previously used CFC Replacement CFC-12 (CCl2F2); CFCHFC-23 (CHF3); HFC-134a 11(CCl3F); CFC(CF3CFH2); HFC-507 (a 1:1 13(CClF3); HCFC-22 azeotropic mixture of HFC 125 (CF3 Refrigeration & (CHClF2); CFC-113 CHF2) and HFC-143a (CF3CH3)); air-conditioning (Cl2FCCClF2); CFC-114 HFC 410 (a 1:1 azeotropic mixture (CClF2CClF2); CFC-115 of HFC-32 (CF2H2) and HFC-125 (CF3CClF2); (CF3CF2H)) Propellants in HFC-134a (CF3CFH2); HFC-227ea medicinal CFC-114 (CClF2CClF2) (CF3CHFCF3) aerosols Blowing agents CFC-11 (CCl3F); CFC 113 HFC-245fa (CF3CH2CHF2); HFCApplication

for foams

(Cl2FCCClF2); HCFC-141b 365 mfc (CF3CH2CF2CH3) (CCl2FCH3)

Solvents, degreasing CFC-11 (CCl3F); CFC-113 None agents, cleaning (CCl2FCClF2) agents

Environmental Impacts As previously discussed, CFCs were phased out via the Montreal Protocol due to their part in ozone depletion. However, the atmospheric impacts of CFCs are not limited to its role as an active ozone reducer. This anthropogenic compound is also a greenhouse gas, with a much higher potential to enhance the greenhouse effect than CO2. Infrared bands trap heat from escaping earth's atmosphere. In the case of CFCs, the strongest of these bands are located at the spectral region - referred to as an atmospheric window due to the relative transparency of the atmosphere within this region. The strength of CFC bands and the unique susceptibility of the atmosphere, at which the compound absorbs and emits radiation, are two factors that contribute to CFC's "super" greenhouse effect. Another such factor is the low concentration of the compound. Because CO2 is close to saturation with high concentrations, it takes more of the substance to enhance the greenhouse effect. Conversely, the low concentration of CFCs allow their effects to increase linearly with mass.

Safety According to their Material Safety Data Sheets, CFCs and HCFCs are colourless, volatile, relatively non-toxic liquids and gases with a faintly sweet ethereal odour. Overexposure may cause dizziness, loss of concentration, Central Nervous System depression and/or cardiac arrhythmia. Vapors displace air and can cause asphyxiation in confined spaces. Although non-flammable, their combustion products include hydrofluoric acid, and related species.

Chapter-7

Perfluorocarbon

Perfluorohexane, a fluorocarbon Perfluorocarbons (PFCs) are fluorocarbons, compounds derived from hydrocarbons by replacement of hydrogen atoms by fluorine atoms. PFCs are made up of carbon and fluorine atoms only, such as octafluoropropane, perfluorohexane and perfluorodecalin. A perfluorocarbon can be arranged in a linear, cyclic, or polycyclic shape. Perfluorocarbon derivatives are perfluorocarbons with some functional group attached, for example perfluorooctanesulfonic acid. Perfluorocarbon derivatives can be very different from perfluorocarbons in their properties, applications and toxicity. The term Perfluorinated compounds or perflourochemical (also abbreviated to PFC) may indicate perfluorcarbons, but is often used to include perfluorocarbon derivatives.

Properties Perfluorocarbons have chemical inertness and thermal stability. This is attributed to the strength of the carbon-fluorine bond and the shielding effect of the fluorine atoms. The electronegativity of fluorine reduces the polarizability of the electron clouds. This results in reduced van der Waals forces between fluorocarbons, as these species tend to be volatile and have low cohesive energies in liquids. There are six perfluorocarbon gases; tetrafluoromethane (carbon tetrafluoride) (bp −128 °C), hexafluoroethane (bp −78.2 °C), octafluoropropane (perfluoropropane) (bp −36.5 °C), perfluorocyclobutane (bp −6 °C), perfluoro-n-butane (bp −2.2 °C) and perfluoro-iso-butane (bp −1 °C). Virtually all the other commercially available perfluorocarbons are liquids (the exception being perfluorocyclohexane, which sublimes at 51 °C. Perfluorocarbon liquids are colorless. They have high density, up to over twice that of water, due to their high molecular weight. Very low intermolecular forces gives the liquids low viscosities (compared to liquids of similar boiling points), low surface tension and low heats of vaporization. They have particularly low refractive indices too. They are not miscible with most organic solvents (e.g., ethanol, acetone, ethyl acetate and chloroform), but are miscible with some hydrocarbons (e.g., hexane in some cases). They have very low solubility in water, and water has a very low solubility in them (on the order of 10 ppm). However, they are relatively good solvents for gases, again because of the very low intermolecular forces. The number of carbon atoms in the perfluorocarbon molecule largely defines most physical properties. The greater the number of carbon atoms, the higher the boiling point, density, viscosity, surface tension, critical properties, vapour pressure and refractive index. Gas solubility decreases as carbon atoms increase, while melting point is determined by other factors as well, so is not readily predicted.

Manufacture Prior to World War II, the only known route to perfluorocarbons was by direct reaction of fluorine with the hydrocarbon. This highly exothermic process was capable only of synthesising tetrafluoromethane, hexafluoroethane and octafluoropropane; larger hydrocarbons decomposed in the extreme conditions. The Manhattan project saw the need for some very robust chemicals, including a wider range of perfluorocarbons, requiring new manufacturing methods. The so-called "catalytic" method involved reacting fluorine and hydrocarbon on a bed of gold-plated copper turnings, the metal removing the heat of the reaction (so not really acting as a catalyst at all), allowing larger hydrocarbons to survive the process. However, it was the Fowler process that allowed the large scale manufacture of perfluorocarbons required for the Manhattan project.

The Fowler Process The Fowler process uses cobalt fluoride to moderate the reaction. In the laboratory, this is typically done in two stages, the first stage being fluorination of cobalt difluoride to cobalt trifluoride. 2 CoF2 + F2 → 2 CoF3 During the second stage, in this instance to make perfluorohexane, the hydrocarbon feed is introduced and is fluorinated by the cobalt trifluoride, which is converted back to cobalt difluoride. Both stages are performed at high temperature. C6H14 + 28 CoF3 → C6F14 + 14 HF + 28 CoF2 Industrially, both steps are combined, for example in the manufacture of the Flutec range of perfluorocarbons, using a vertical stirred bed reactor, with hydrocarbon introduced at the bottom, and fluorine introduced half way up the reactor. The perfluorocarbon vapor is recovered from the top.

Electrochemical fluorination An alternative technique, electrochemical fluorination (ECF) (also known as the Simons' process) involves electrolysis of a substrate dissolved in hydrogen fluoride. As fluorine is itself manufactured by the electrolysis of hydrogen fluoride, this is a rather more direct route to perfluorocarbons. The process is run at low voltage (5 - 6 V) so that free fluorine is not liberated. The choice of substrate is restricted as ideally it should be soluble in hydrogen fluoride. Ethers and tertiary amines are typically employed. To make perfluorohexane, trihexylamine is used, for example: 2 N(C6H13)3 + 90 HF → 6 C6F14 + 2 NF3 + 45 H2 The perfluorocarbon amine will also be produced: N(C6H13)3 + 42 HF → 2 N(C6F13)3 + 21H2 Both of these products, and others, are manufactured by 3M as part of the Fluorinert range.

Medical applications Medical applications require high purity perfluorocarbons. Impurities with nitrogen bonds can have high toxicity; hydrogen-containing compounds (which can release hydrogen fluoride) and unsaturated compounds must also be excluded. Infrared spectroscopy, nuclear magnetic resonance and cell cultures can be used to test the perfluorocarbon.

Eye surgery Perfluorocarbons are commonly used in eye surgery as temporary replacements of the vitreous humor in retinal detachment surgery. Retinal tears following a penetrating trauma or retinal detachments associated with proliferative vitreoretinopathy can be corrected with surgery in which the dense perfluorocarbon liquid, typically perfluoro-noctane, is injected into the eye, to push out vitreous liquid trapped behind the retina, and to aid removal of membranes (essentially scar tissue). Perfluoro-1,3-dimethylcyclohexane has been used in the removal of a lens nucleus dislocated into the vitreous cavity, the lens floating on the heavy perfluorocarbon for easy removal . Octafluoropropane can be used almost in a reverse sense. It is injected into the eye diluted in air (typically 12% to 16%). The patient must then lie face down for about an hour. The gas bubble pushes onto the retina to perform the same task as before . The octafluorpropane may remain in the eye for up to three months after surgery before it is completely expelled. Air travel or other environments involving changes in pressure should be avoided. Use of nitrous oxide as an anaesthetic can be disastrous to the possible future optical abilities of the patient ; dissolved nitrous oxide from the blood accumulates in the bubble, increasing intraocular pressure to the point that blood flow to the retina is cut off and the retina dies.

Imaging Perfluorocarbons are also used in contrast-enhanced ultrasound to improve ultrasound signal backscatter. The perfluorocarbons used in the microbubbles are gases at body temperature (though they may be liquids at room temperature). The gas-filled microbubbles oscillate and vibrate when a sonic energy field is applied and characteristically reflect ultrasound waves. This distinguishes the microbubbles from surrounding tissues. Their stability, inertness, low diffusion rate and solubility increase the duration of contrast enhancement as compared to microbubbles containing air. Perfluorocarbons can also be used in magnetic resonance imaging (MRI), though this is not as common. Usually MRI is set up to detect hydrogen nuclei, but it is also possible to use MRI for 19-fluorine nuclei. As there is no fluorine in the human body naturally, it is very easy to determine exactly where the sample has gone. Perfluorocarbons can be introduced into the blood in an emulsion, or neat in the lungs. In radiographic imaging, the perfluorocarbon derivative perfluorooctyl bromide (PFOB) is employed, as this is more opaque to X-rays.

Liquid breathing Perfluorocarbons dissolve relatively high concentrations of gases, for example, 100 ml of perfluorodecalin at 25°C will dissolve 49 ml of oxygen at STP. This led Leland C. Clark in 1966 to experiment with liquid breathing, resulting in the submersion of a mouse for several hours in an oxygenated perfluorocarbon. The mice he used later died due to

trauma to their lungs; however, this may have been due to impurities in the perfluorocarbon. In recent years there has been new interest in liquid breathing for various procedures from lung lavage to treatment of congenital diaphragmatic hernia. Perfluorocarbon liquids (and liquids in general) are much denser and more viscous than air; rates of breathing, and therefore of gas exchange, are limited, and there are challenges still to be overcome such as efficient removal of carbon dioxide.

Artificial blood Clark's experiments also triggered interest in using perfluorocarbons in so-called "artificial blood" as oxygen therapeutics which function as artificial erythrocytes, serving to transport and deliver oxygen. The Green Cross Corporation attempted to commercialize this technology in the 1980s under the Fluosol tradename, without success. Recently, however, there has been renewed interest in this field. In this application, the perfluorocarbon is used as a part of an emulsion, typically using Pluronic F-68 or egg yolk phospholipids (lecithin) as surfactants, in water. For example, Fluosol-DC: Ingredient w/v% Perfluorodecalin 25.0 Yolk phospholids 3.6 Fatty acid (emulsion stabilizer) trace D-Sorbitol (emulsion stabilizer) 3.5 NaCl 0.204 KCl 0.010 0.007 MgCl2 Sodium lactate 0.105 A perfluorocarbon droplet size of 0.1 -0.2μm enables the droplets to be present in the plasma gaps between erythrocytes in the microcirculation structures, this is an advantage when oxygen supply to tissue by red blood cells is low due to acute anemia or hemodilution. Perfluorocarbons are most effective in small capillaries or blood vessels, where the droplets can be present in the plasma between red blood cells, augmenting local oxygen delivery and increasing the oxygen content in the arterial blood.

Side effects Some common side effects were recorded while performing clinical studies such as delayed febrile reaction and flu-like symptoms. The magnitude of the side effects is directly related to the size of the emulsion droplets, if the particles are smaller than 0.2μm it seems to be undetectable for the reticuloendothelial system. These side effects occur when the body is excreting/eliminating the perfluorocarbon. The excretion depends on vapour pressure and lipid solubility of the perfluorocarbons. It usually takes on average

three to four days for the compound perfluoroocytl bromide and eight days for perfluorodichlorooactane. This process is relatively slow since perfluorocarbons are inert to biochemical degradation. The perfluorocarbon will then diffuse back into the blood where they dissolve into plasma lipids. The plasma lipids will then transport the perfluorocarbon molecules to the lungs where they are excreted through exhalation along with other gases.

Treatment of decompression sickness Perfluorocarbons accelerate nitrogen washout after venous gas emboli. Success in the treatment of decompression sickness has been shown in rat, swine, and hamster models. This treatment shows great potential as a future adjunctive therapy for decompression sickness in humans.

Non-medical applications Electrical and electronic applications Perfluorocarbons have high dielectric strengths and high insulating properties, and so can be used in direct contact with high voltage components, either as dielectric fluids, dielectric gases, or as coolants.

Perfluorcarbon tracers Perfluorocarbons can be detected at extremely low levels using electron capture detectors or negative ion mass spectroscopy. They can be released at a certain point and the concentration measured in the surrounding area. Perfluorocarbon tracers (PFTs) have been used to map oil fields, study building ventilation, track pollution, detect cable oil leaks and even recover ransom money.

Cosmetics Inspired by the medical applications, several companies incorporate perfluorocarbons in their cosmetic formulations, claiming the oxygen dissolved in the perfluorocarbon has an anti-aging effect on the skin.

Other applications PFCs are being used in refrigerating units as replacements for CFCs (haloalkanes), often in conjunction with other gases, and as "clean" fire extinguishers. They are used in plasma cleaning of silicon wafers. Perfluorocarbons are also used in high end racing ski waxes due to their hydrophobic nature, which is responsible for reduced friction in wet snow conditions. In fluorous biphase catalysis a perfluorocarbon is used to dissolve a catalyst with a perfluoroalkyl group, while the substrate is dissolved in an organic solvent. At elevated

temperature, the perfluorocarbon and organic solvent become miscible, and so the mixture becomes homogeneous, facilitating the reaction. Upon cooling, the two phases separate, allowing the catalyst to be recovered from the perfluorocarbon, and the product from the organic solvent. PFCs are being used in many unexpected places, including pizza boxes, popcorn bags, lipstick, computer mice, in water repellants, stain repellents and as the lining to non-stick cook ware.

Environmental effects PFCs are extremely potent greenhouse gases, and they are a long-term problem with a lifetime up to 50,000 years. In a 2003 study, the most abundant atmospheric PFC was tetrafluoromethane. The greenhouse warming potential (GWP) of tetrafluoromethane is 6,500 times that of carbon dioxide, and the GWP of hexafluoroethane is 9,200 times that of carbon dioxide. Several governments concerned about the properties of PFCs have already tried to implement international agreements to limit their usage before it becomes a global warming issue. PFCs are one of the classes of compounds regulated in the Kyoto Protocol. The primary source of tetrafluoromethane in the environment is from the production of aluminium by electrolysis of alumina. Aluminium producers are taking effective steps in reducing emissions by better controlling the electrolysis process Two PFC derivatives, perfluorooctanesulfonic acid and perfluorooctanoic acid, have been found to be persistent in the environment and are detected in blood samples all over the world.

Chapter-8

Tetrafluoroborates

Fluoroboric acid Fluoroboric acid

IUPAC name Tetrafluoroboric acid Other names Fluoroboric acid; Hydrogen Tetrafluoroborate; Hydrofluoroboric acid; Borofluoric acid Identifiers CAS number 16872-11-0 ChemSpider 10677839 EC number 240-898-3 UN number 1775 RTECS ED2685000 number Properties Molecular HBF4 formula Molar mass 87.81 g/mol Appearance Clear liquid

Density Melting point Boiling point Solubility in water Acidity (pKa)

variable -90 °C 130 °C Miscible -0.4 Structure

Crystal structure MSDS EU Index EU classification R-phrases S-phrases Flash point Other anions Related compounds

N/A Hazards External MSDS 009-010-00-X Corrosive (C) R34 (S1/2), S26, S27, S45 non-flammable Related compounds Hexafluorophosphoric acid, Triflic acid Potassium fluoroborate, nitrosonium fluoroborate, Hexafluorophosphate, Hydrogen fluoride

Fluoroboric acid (also spelt fluoboric acid) is the chemical compound with the formula HBF4. It is the conjugate acid of tetrafluoroborate. It is available commercially as a solution in water and other solvents such as diethyl ether. With a strength comparable to nitric acid, fluoroboric acid is a strong acid with a weakly coordinating, non-oxidizing conjugate base.

Production Pure fluoroboric acid has never been produced but aqueous solutions of HBF4 can be produced by dissolving boric acid in aqueous hydrofluoric acid solution at 20-25 °C. Three equivalents of HF react to give the intermediate boron trifluoride and the fourth gives fluoroboric acid. B(OH)3 + 4 HF → H3O+ + BF4− + 2 H2O Aqueous solutions of fluoroboric acid can also be prepared by treating impure hexafluorosilicic acid with solid boric acid followed by removal of precipitated silicon dioxide. Anhydrous solutions can be prepared by treatment with acetic anhydride.

Salts Fluoroboric acid is the principal precursor to fluoroborate salts, which are typically prepared by acid-base reactions. The inorganic salts are intermediates in the manufacture

of flame-retardant materials, glazing frits, and in electrolytic generation of boron. HBF4 is also used in aluminum etching and acid pickling.

Applications Organic chemistry HBF4 is used as a catalyst in for alkylations and polymerizations. In carbohydrate protection reactions, ethereal fluoroboric acid is an efficient and cost-effective catalyst for transacetalation and isopropylidenation reactions. Acetonitrile solutions cleave acetals and some ethers, while neat fluoroboric acid removes tert-butoxycarbonyl groups.

Galvanic cells Aqueous HBF4 is used as an electrolyte in galvanic cell oxygen sensor systems which consist of an anode, cathode, and oxygen-permeable membrane. The solution of HBF4 is able to dissolve lead(II) oxide from the anode in the form of lead tetrafluoroborate while leaving the rest of the system unchanged.

Metal plating A mixture of CrO3, HBF4, and sulfonic acids in conjunction with a cathode treatment give tin-plated steel. Tin(I) fluoroborate/fluoroboric acid mixtures and organic reagents are used as the electrolyte in the cathode treatment of the tin plating process. Similar processes of electrodeposition and electrolytic stripping are used to obtain specific metal alloys.

Other fluoroboric acids A series of fluoroboric acids is known in aqueous solutions. The series can be presented as follows:     

H[B(OH)4] H[BF(OH)3] H[BF2(OH)2] H[BF3(OH)] H[BF4]

Lithium tetrafluoroborate Lithium tetrafluoroborate

IUPAC name Lithium tetrafluoroborate Other names Borate(1-), tetrafluoro-, lithium Identifiers CAS number 14283-07-9 ChemSpider 3504162 Properties Molecular formula LiBF4 Molar mass 93.746 g/mol Appearance White/grey crystalline solid Odor odorless Density 0.852 g/cm3 solid Melting point 296.5 °C Boiling point decomp Solubility in water Very soluble Hazards MSDS External MSDS Harmful, causes burns, Main hazards hygroscopic. Related compounds Other anions Tetrafluoroborate, Related compounds Nitrosyl tetrafluoroborate

Lithium tetrafluoroborate is a chemical compound with the formula LiBF4. It can be dissolved in propylene carbonate, dimethoxyethane, and/or gamma-butyrolactone for use as an electrolyte in lithium batteries.

Nitronium tetrafluoroborate Nitronium tetrafluoroborate

Other names nitronium fluoroborate, NO2BF4 Identifiers CAS number 13826-86-3 Properties Molecular formula BNO2F4 Molar mass 132.81 Hazards MSDS R-phrases R34 R42 R43 S-phrases S26 S36 S37 S39 S45 Related compounds Other cations Nitrosonium tetrafluoroborate

Nitronium tetrafluoroborate is an inorganic compound with formula NO2BF4. It is a salt of nitronium cation and tetrafluoroborate anion. It is a hygroscopic, colorless crystalline solid. It is corrosive.

Preparation Nitronium tetrafluoroborate can be prepared by adding a mixture of anhydrous hydrogen fluoride and boron trifluoride to a nitromethane solution of nitric acid or nitrogen pentoxide.

Applications Nitronium tetrafluoroborate is used as a nitration agent.

Nitrosonium tetrafluoroborate Nitrosonium tetrafluoroborate

IUPAC name nitrosonium tetrafluoroborate Other names nitrosyl tetrafluoroborate Identifiers CAS number 14635-75-7 PubChem 151929 Properties Molecular formula BF4NO Molar mass 116.81 g mol−1 Appearance colourless crystalline solid Density 2.185 g cm−3 Melting point 250 °C (sublimes) Solubility in water decomposes

Nitrosonium tetrafluoroborate is a chemical compound with the chemical formula NOBF4. This colourless solid finds use in organic synthesis as a nitrosating agent. NOBF4 is the nitrosonium salt of fluoroboric acid, and is composed of a nitrosonium cation, [NO]+, and a tetrafluoroborate anion, [BF4]−.

Reactions Nitrosonium tetrafluoroborate may be used to prepare metal salts of the type [MII(CH3CN)x][BF4]2 (M = Cr, Mn, Fe, Co, Ni, Cu). The nitrosonium cation acts as the oxidizer, itself being reduced to nitric oxide gas: With ferrocene the ferrocenium tetrafluoroborate is formed.

M + NOBF4 + xCH3CN → [M(CH3CN)x](BF4)2 + NO

Silver tetrafluoroborate Silver tetrafluoroborate

IUPAC name Silver tetrafluoridoborate(1–) Other names Borate(1-), tetrafluoro-, silver(1+) Identifiers CAS number 14104-20-2 PubChem 159722 ChemSpider 140438 Properties Molecular formula AgBF4 Molar mass 194.67 g/mol Appearance Off-white powder Hazards MSDS External MSDS EU classification Corrosive (C)

Silver tetrafluoroborate (AgBF4) sometimes referred to "silver BF-4" is an inorganic compound commonly encountered in inorganic and organometallic chemistry. Similar to silver hexafluorophosphate, it is commonly used to replace halide anions or ligands with the weakly-coordinating tetrafluoroborate anions. The abstraction of the halide is driven by the precipitation of the appropriate silver halide.

Triethyloxonium tetrafluoroborate Triethyloxonium tetrafluoroborate

IUPAC name Triethyloxonium tetrafluoroborate Identifiers CAS number 368-39-8 Properties Molecular formula C6H15BF4O Molar mass 189,99 g/mol Melting point 91–92 °C Boiling point dec. Hazards R-phrases 34-40 S-phrases 23-24/25-26-36/37/39-45

Triethyloxonium tetrafluoroborate is the organic oxonium compound with the formula [(CH3CH2)3O]BF4. It is often called Meerwein's reagent after its discoverer Hans Meerwein. Also well known and commercially available is the related trimethyloxonium tetrafluoroborate. The compounds are exceptionally strong alkylating agents. Aside from the BF4− salt, many related derivatives are available with varying solubilities and stabilities.

Synthesis Triethyloxonium tetrafluoroborate is prepared from boron trifluoride, diethyl ether, and epichlorohydrin: 4 Et2O·BF3 + 2 Et2O + 3 C2H3(O)CH2Cl → 3 Et3O+BF4− + B[(OCH(CH2Cl)CH2OEt]3

The trimethyloxonium salt is available from dimethyl ether via an analogous route. These salts do not have long shelf-lives at room temperature. These salts degrade by hydrolysis: [(CH3CH2)3O]+BF4− + H2O → (CH3CH2)2O + CH3CH2OH + HBF4 The propensity of trialkyloxoniums to undergo alkyl-exchange may be utilized to the chemists' advantage. For example, trimethyloxonium tetrafluoroborate, which reacts sluggishly due to low solubility in most compatible solvents may be converted in-situ to higher alkyl/more soluble oxoniums, thereby speeding up alkylation reactions.

Structure The compound features pyramidal oxonium cation and a tetrahedral fluoroborate anion. Reflecting its ionic character, the salt dissolves in polar but inert solvents such as dichloromethane, sulfur dioxide, and nitromethane.

Safety Triethyloxonium tetrafluoroborate is a strong alkylating agent, although the hazards are diminished because it is non-volatile. It releases strong acid upon contact with water. The properties of the methyl derivative are similar.

Use Alkylating agent for nucleophilic functional groups in organic synthesis.

Tropylium tetrafluoroborate Tropylium tetrafluoroborate

CAS number ChemSpider Molecular formula Molar mass

Identifiers 27081-10-3 134784 Properties C7H7BF4 177.94 g mol−1

Tropylium tetrafluoroborate is an organic compound containing the tropylium cation and the non-coordinating tetrafluoroborate counteranion. It is a stable salt which is commercially available.

This compound may be prepared by the reaction of cycloheptatriene with phosphorus trichloride, followed by tetrafluoroboric acid.

Chapter-9

Fluoride Compounds

Ammonium bifluoride Ammonium bifluoride

IUPAC name Ammonium hydrogen fluoride Other names Ammonium acid fluoride Ammonium hydrofluoride Ammonium difluoride Ammonium hydrogendifluoride Ammonium hydrogen difluoride Identifiers CAS number 1341-49-7 ChemSpider 21241205 Properties Molecular formula H5F2N Molar mass 57.04 g mol−1 Exact mass 57.039005575 g mol-1 Appearance White crystals Density 1.50 g cm-3 Melting point 126 °C, 399 K, 259 °F (decomposes) Boiling point 240 °C, 513 K, 464 °F Solubility in water very soluble

Solubility in alcohol slightly soluble Refractive index (nD) 1.390 Structure Crystal structure Cubic, related to the CsCl structure Coordination [NH4]+ cation: tetrahedral [HF2]− anion: linear geometry Related compounds Other cations potassium bifluoride Related compounds ammonium fluoride

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.

Structure Ammonium bifluoride, as its name indicates, contains a bifluoride, or hydrogen(difluoride) anion: HF2−. This centrosymmetric triatomic anion features the strongest known hydrogen bond, with a F−H length of 114 pm. and a bond energy greater than 155 kJ mol−1. In solid [NH4][HF2], each ammonium cation is surrounded by four fluoride centers in a tetrahedron, with hydrogen - fluorine hydrogen bonds present between the hydrogen atoms of the ammonium ion and the fluorine atoms. Solutions contain tetrahedral [NH4]+ cations and linear [HF2]− anions.

Production and applications Ammonium bifluoride is a component of some etchants. It attacks silica component of glass: SiO2 + 4 [NH4][HF2] → SiF4 + 4 [NH4]F + 2 H2O Potassium bifluoride is a related more commonly used etchant. Ammonium bifluoride has been considered as an intermediate in the production of hydrofluoric acid from hexafluorosilicic acid. Thus, hexafluorosilicic acid is hydrolyzed to give ammonium fluoride, which thermally decomposes to give the bifluoride: H2SiF6 + 6 NH3 + 2 H2O → SiO2 + 6 NH4F 2 NH4F → NH3 + [NH4]HF2 The resulting ammonium bifluoride is converted to the sodium bifluoride, which thermally decomposes to release HF.

Ammonium fluoride Ammonium fluoride

IUPAC name Ammonium fluoride Other names Neutral ammonium fluoride Identifiers CAS number 12125-01-8 ChemSpider 23806 EC number 235-185-9 UN number 2505 RTECS number BQ6300000 Properties Molecular NH4F formula Molar mass 37.037 g/mol White crystalline solid Appearance hygroscopic Density 1.009 g/cm3 Melting point 100 °C (decomp) Solubility in water 45.3 g/100 ml (25 °C) slightly soluble in alcohol, insoluble in Solubility liquid ammonia Structure Crystal structure Wurtzite structure (hexagonal) Hazards MSDS ICSC 1223 EU Index 009-006-00-8 EU classification Toxic (T) R-phrases R23/24/25 S-phrases (S1/2), S26, S45 Flash point Non-flammable

Other anions Other cations Related compounds

Related compounds Ammonium chloride Ammonium bromide Ammonium iodide Sodium fluoride Potassium fluoride Ammonium bifluoride

Ammonium fluoride is the inorganic compound with the formula NH4F. It crystallizes as small colourless prisms, having a sharp saline taste, and is exceedingly soluble in water.

Reactions Ammonium fluoride adopts the wurtzite crystal structure, in which both the ammonium cations and the fluoride anions are stacked in ABABAB... layers, each being tetrahedrally surrounded by four of the other. There are NH...F hydrogen bonds between the anions and cations. On passing hydrogen fluoride gas (in excess) through the salt, ammonium fluoride absorbs the gas to form the addition compound ammonium hydrogen fluoride. The reaction occurring is: NH4F + HF → NH4HF2 It sublimes when heated - a property common among ammonium salts. In the sublimation, the salt decomposes to ammonia and hydrogen fluoride , and the two gases recombine to give ammonium fluoride, i.e. the reaction is reversible: [NH4]F ↔ NH3 + HF

Uses This substance is commonly called "commercial ammonium fluoride". The word "neutral" is sometimes added to "ammonium fluoride" to represent the neutral salt [NH4]F vs the "acid salt" (NH4HF2), which contains a higher percentage. The acid salt is usually used in preference to the neutral salt in the etching of glass and related silicates. This property is shared among all soluble fluorides. For this reason it cannot be handled in glass test tubes or apparatus during laboratory work. It is also used for preserving wood, as a mothproofing agent, in printing and dying textiles, and as an antiseptic in breweries.

Dioxygen difluoride Dioxygen difluoride

Preferred IUPAC name Dioxygen difluoride Systematic name Fluorooxy hypofluorite Other names Difluorine dioxide Fluorine dioxide Perfluoroperoxide Identifiers Abbreviations FOOF CAS number 7783-44-0 PubChem 123257 ChemSpider 109870 ChEBI CHEBI:47866 Gmelin Reference 1570 Properties Molecular formula O2F2 Molar mass 69.996 g·mol−1 Melting point −154 °C, 119 K, -245 °F −57 °C, 216 K, -71 °F Boiling point (extrapolated) Solubility in other decomp. solvents Related compounds H2O2 OF2 Related compounds FClO2 S2Cl2

Dioxygen difluoride is a compound with the formula O2F2. It exists as an orange solid that melts into a red liquid at −163 °C It is a strong oxidant and decomposes into OF2 and oxygen even at −160 °C (4% per day).

Preparation Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17 mmHg is optimal) to an electric discharge of 25–30 mA at 2.1–2.4 kV. This is basically the reaction used for the first synthesis by Otto Ruff in

1933. Another synthesis involves mixing O2 and F2 in a stainless steel vessel cooled to −196 °C, followed by exposing the elements to 3 MeV bremsstrahlung for several hours.

Structure and electronic description In O2F2, oxygen is assigned the unusual oxidation state of +1. In most of its other compounds, oxygen has an oxidation state of −2. The structure of dioxygen difluoride resembles that of hydrogen peroxide, H2O2, in its large dihedral angle, which approaches 90°. This geometry conforms with the predictions of VSEPR theory. The O−O bond length is within 2 pm of the 120.7 pm distance for the O=O double bond in dioxygen, O2.

The bonding within dioxygen difluoride has been the subject of considerable speculation over the years, particularly because of the very short O–O distance and the long O–F distances. Bridgeman has proposed a scheme which essentially has an O–O triple bond and an O–F single bond that is destabilised and lengthened by repulsion between the lone pairs on the fluorine atoms and the π-orbitals of the O–O bond. Repulsion involving the fluorine lone pairs is also responsible for the long and weak covalent bonding in the fluorine molecule.

Reactivity The overarching property of this unstable compound is its oxidizing power, despite the fact that all reactions must be conducted near −100 °C. With BF3 and PF5, it gives the corresponding dioxygenyl salts: 2 O2F2 + 2 PF5 → 2 [O2]+[PF6]− + F2 It converts uranium and plutonium oxides into the corresponding hexafluorides.

Fluorosulfuric acid Fluorosulfuric acid

Preferred IUPAC name Sulfurofluoridic acid Systematic name Fluoranesulfonic acid Other names Fluorosulfonic acid Identifiers CAS number 7789-21-1 , 29171-24-2 (2H) PubChem 24603 , 168865 (2H) ChemSpider 23005 , 147714 (2H) EC number 232-149-4 UN number 1777 MeSH Fluorosulfonic+acid RTECS number LP0715000 Properties Molecular formula FSO3H Molar mass 100.06 g/mol Appearance Colorless liquid Density 1.84 g/cm3, liquid Melting point −87.3 °C Boiling point 165.5 °C Solubility in water Soluble −10 Acidity (pKa) Structure Molecular shape tetrahedral Dipole moment N/A Hazards MSDS ICSC 0996 EU Index 016-018-00-7 Harmful (Xn) EU classification Corrosive (C) Oxidant(O) R-phrases R20, R35,R8,R5 S-phrases (S1/2), S26, S45 Related compounds Related compounds Antimony pentafluoride

Trifluoromethanesulfonic acid Hydrofluoric acid

Fluorosulfuric acid (IUPAC name: sulfurofluoridic acid) is the inorganic compound with the formula HSO3F. It is one of the strongest acids commercially available and is a superacid. The formula HFSO3 emphasizes its relationship to sulfuric acid, H2SO4; HSO3F is a tetrahedral molecule.

Chemical properties Fluorosulfuric acid is a free-flowing colorless liquid. It is soluble in polar organic solvents (e.g. nitrobenzene, acetic acid, and ethyl acetate), but poorly soluble in nonpolar solvents such as alkanes. Reflecting its strong acidity, it dissolves almost all organic compounds that are even weak proton acceptors. FSO3H hydrolyzes slowly to HF and sulfuric acid. The related triflic acid (CF3SO3H) retains the high acidity of FSO3H but is more hydrolytically stable.

Production Fluorosulfuric acid is prepared by the reaction of HF and sulfur trioxide: SO3 + HF → FSO3H Alternatively, KHF2 or CaF2 can be treated with oleum at 250 °C. Once freed from HF by sweeping with an inert gas, FSO3H can be distilled in a glass apparatus.

Super-acids FSO3H is one of the strongest known simple Brønsted acids, although recent work on carborane-based acids have led to still stronger acids. It has an H0 value of −15.1 compared to −12 for sulfuric acid. The combination of FSO3H and the Lewis acid antimony pentafluoride produces "Magic acid," which is a far stronger protonating agent. These acids all fall into the category of "superacids", acids stronger than 100% sulfuric acid.

Applications FSO3H is useful for regenerating mixtures of HF and H2SO4 for etching lead glass. FSO3H isomerizes alkanes and the alkylation of hydrocarbons with alkenes, although it is unclear if such applications are of commercial importance. It can also be used as a laboratory fluorinating agent.

Safety Fluorosulfuric acid is considered to be highly toxic and corrosive. It hydrolyzes to release HF. Addition of water to FSO3H can be violent, similar to the addition of water to sulfuric acid. Addition of fluorosulfuric acid to water is much more violent process than addition of sulfuric acid.

Formyl fluoride Formyl fluoride

IUPAC name Formyl fluoride Other names Formic acid fluoride Identifiers CAS number 1493-02-3 PubChem 15153 ChemSpider 14424 Properties Molecular formula CHFO Molar mass 48.02 g/mol Appearance Colourless gas Melting point -142 °C Boiling point –29 °C Solubility in water decomposes chlorocarbons, Solubility in other solvents Freons Structure Dipole moment 2.02 D Hazards Main hazards toxic Related compounds Formic acid Related compounds Hydrogen fluoride Carbonyl fluoride

Formyl fluoride is the organic compound with the formula FC(O)H.

Preparation FC(O)H was first reported in the 1934. Among the many preparations, a typical one involves the reaction of sodium formate with benzoyl fluoride (generated in situ from KHF2 and benzoyl chloride): NaOC(O)H + C6H5C(O)F → FC(O)H + C6H5CO2Na

Structure The molecule is planar; C-O and C-F distances are 1.18 and 1.34 A, respectively.

Reactions HC(O)F decomposes autocatalytically near room temperature to carbon monoxide and hydrogen fluoride: HC(O)F → HF + CO Because of the compound’s sensitivity, reactions are conducted at low temperatures and samples are often stored over anhydrous alkali metal fluorides, e.g. potassium fluoride which absorbs HF. Benzene (and other arenes) react with formyl fluoride in the presence of boron trifluoride to give benzaldehyde. In a related reaction, formyl chloride is implicated in GattermannKoch formylation reaction. The reaction of formyl fluoride/BF3 with perdeuteriobenzene (C6D6) exhibits a kinetic isotope effect of 2.68, similar to the isotope effect observed in Friedel-Crafts acetylation of benzene. Formylation of benzene with a mixture of CO and hexafluoroantinomic acid however, exhibits no isotope effect (C6H6 and C6D6 react at the same rate), indicating that this reaction involves a more reactive formylating agent, possibly CHO+. Formyl fluoride undergoes the reactions expected of an acyl halide: alcohols and carboxylic acids are converted to formate esters and mixed acid anhydrides, respectively.

Oxygen difluoride Oxygen difluoride

Other names difluorine monoxide fluorine monoxide oxygen fluoride hypofluorous anhydride Identifiers CAS number 7783-41-7 PubChem 24547 ChemSpider 22953 RTECS number RS2100000 Properties Molecular formula OF2 Molar mass 53.9962 g/mol colorless gas, pale yellow liquid when Appearance condensed Density 1.9 g/cm3 Melting point −223.8 °C Boiling point −144.8 °C Solubility in other 68 mL gaseous OF2 in 1 L (0 °C) solvents Thermochemistry Std enthalpy of 24.5 kJ mol−1 formation ΔfHo298 Related compounds HFO O2F2 Related compounds NHF2 NF3 SCl2

Oxygen difluoride is the chemical compound with the formula OF2. As predicted by VSEPR theory, the molecule adopts a "V" shaped structure like H2O, but it has very different properties, being a strong oxidizer.

Preparation Oxygen difluoride was first reported in 1929; it was obtained by the electrolysis of molten potassium fluoride and hydrofluoric acid containing small quantities of water. The modern preparation entails the reaction of fluorine with a dilute aqueous solution of sodium hydroxide, with sodium fluoride as a side-product: 2 F2 + 2 NaOH → OF2 + 2 NaF + H2O

Reactions Its powerful oxidizing properties are suggested by the oxidation number of +2 for the oxygen atom, which is unusual. Above 200 °C, OF2 decomposes to oxygen and fluorine via a radical mechanism. OF2 reacts with many metals to yield oxides and fluorides. Nonmetals also react: phosphorus reacts with OF2 to form PF5 and POF3; sulfur gives SO2 and SF4; and unusually for a noble gas, xenon reacts, yielding XeF4 and xenon oxyfluorides. Oxygen difluoride reacts very slowly with water to form hydrofluoric acid: OF2 (aq) + H2O (aq) → 2 HF (aq) + O2 (g) Oxygen difluoride oxidizes sulfur dioxide to sulfur trioxide: OF2 + SO2 → SO3 + F2 However, in the presence of UV radiation the products are sulfuryl fluoride, SO2F2, and pyrosulfuryl fluoride, S2O5F2: OF2 + 2 SO2 → S2O5F2

Safety OF2 is a dangerous chemical, as is the case for any strongly oxidizing gas.

Hexafluoride

Hexafluoride-forming elements A hexafluoride is a chemical compound with the general formula XF6. Sixteen elements are known to form stable hexafluorides. Nine of these elements are transition metals, three are actinides, and four are nonmetals or metalloids.

Properties Physical properties

Octahedral structure of SF6 Most hexafluorides are molecular compounds with low melting and boiling points. Four hexafluorides (S, Se, Te, and W) are gases at room temperature (25 °C) and a pressure of 1 atm, two are liquid (Re, Mo), and the others are volatile solids. The p-block and group 6 hexafluorides are colorless, but the other hexafluorides have colors ranging from yellow to orange, red, brown, and black. The molecular geometry is generally octahedral, though it is sometimes distorted. XeF6 is a fluxional molecule with a distorted octahedral structure, which is, according to VSEPR theory, caused by the non-bonding lone pair. In the solid state, XeF6 has a complex structure involving tetramers and hexamers. According to quantum chemical calculations,

ReF6 and RuF6 should have tetragonally distorted structures (where the two bonds along one axis are longer or shorter than the other four), but this has not been verified experimentally. Compound Sulfur hexafluoride Selenium hexafluoride Tellurium hexafluoride Xenon hexafluoride Molybdenum hexafluoride Technetium hexafluoride Ruthenium hexafluoride Rhodium hexafluoride Tungsten hexafluoride Rhenium hexafluoride Osmium hexafluoride Iridium hexafluoride Platinum hexafluoride Uranium hexafluoride Neptunium hexafluoride Plutonium hexafluoride

m.p (°C)

b.p. (°C)

subl.p. (°C) −50.8 −46.6

solid ρ (g Bond Color cm−3) (pm) 1.88 (−50 146.06 156.4 colorless °C) 167– 192.95 colorless 170 MW

−38.9 −37.6

241.59

49.5

75.6

245.28 3.56

17.5

34.0

209.94

37.4

55.3

≈ 70 17.1

18.5

33.7

33.4

47.5

44

53.6

61.3

69.1 56.5

colorless colorless

3.50 181.7 (−140 °C) 3.58 (212) 181.2 (−140 °C) 3.68 215.07 181.8 (−140 °C) 3.71 216.91 182.4 (−140 °C) 4.86 297.85 182.6 (−140 °C) 4.94 300.20 182.3 (−140 °C) 5.09 304.22 182.9 (−140 °C) 5.11 306.21 183.4 (−140 °C) 5.21 309.07 184.8 (−140 °C)

54

2.3

184

colorless yellow dark brown black colorless yellow yellow yellow deep red

351.99 5.09

199.6

white

54.4

55.18

(358)

198.1

orange

52

62

(356) 5.08

197.1

brown

Chemical properties The hexafluorides have a wide range of chemical reactivity. Sulfur hexafluoride is nearly inert and non-toxic. It has several applications due to its stability, dielectric properties,

and high density. Selenium hexafluoride is nearly as unreactive as SF6, but tellurium hexafluoride is toxic, not very stable and can be hydrolyzed by water within 1 day. In contrast, metal hexafluorides are corrosive, readily hydrolyzed and may react violently with water. Some of them can be used as fluorinating agents. The metal hexafluorides have a high electron affinity, which makes them strong oxidizing agents. Platinum hexafluoride in particular is notable for its ability to oxidize the dioxygen molecule, O2, to form dioxygenyl hexafluoroplatinate, and for being the first compound that was observed to react with xenon.

Applications Some metal hexafluorides find applications due to their volatility. Uranium hexafluoride is used in the uranium enrichment process to produce fuel for nuclear reactors. Fluoride volatility can also be exploited for nuclear fuel reprocessing. Tungsten hexafluoride is used in the production of semiconductors through the process of chemical vapor deposition.