Names | |||
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IUPAC name Boron trifluoride | |||
Systematic IUPAC name Trifluoroborane | |||
Other names Boron fluoride, Trifluoroborane | |||
Identifiers | |||
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3D model (JSmol) | |||
ChEBI | |||
ChemSpider | |||
ECHA InfoCard | 100.028.699 | ||
EC Number |
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PubChem CID | |||
RTECS number |
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UNII | |||
UN number | compressed: 1008. boron trifluoride dihydrate: 2851. | ||
CompTox Dashboard (EPA) | |||
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Properties | |||
BF3 | |||
Molar mass | 67.82 g/mol (anhydrous) 103.837 g/mol (dihydrate) | ||
Appearance | colorless gas (anhydrous) colorless liquid (dihydrate) | ||
Odor | Pungent | ||
Density | 0.00276 g/cm3 (anhydrous gas) 1.64 g/cm3 (dihydrate) | ||
Melting point | −126.8 °C (−196.2 °F; 146.3 K) | ||
Boiling point | −100.3 °C (−148.5 °F; 172.8 K) | ||
exothermic decomposition [1] (anhydrous) very soluble (dihydrate) | |||
Solubility | soluble in benzene, toluene, hexane, chloroform and methylene chloride | ||
Vapor pressure | >50 atm (20 °C) [2] | ||
0 D | |||
Thermochemistry | |||
Heat capacity (C) | 50.46 J/(mol·K) | ||
Std molar entropy (S⦵298) | 254.3 J/(mol·K) | ||
Std enthalpy of formation (ΔfH⦵298) | −1137 kJ/mol | ||
Gibbs free energy (ΔfG⦵) | −1120 kJ/mol | ||
Hazards [3] [4] | |||
GHS labelling: | |||
Danger | |||
H314, H330, H335, H373 | |||
P260, P280, P303+P361+P353, P304+P340, P305+P351+P338, P310, P403+P233 | |||
NFPA 704 (fire diamond) | |||
Flash point | Nonflammable | ||
Lethal dose or concentration (LD, LC): | |||
LC50 (median concentration) | 1227 ppm (mouse, 2 hr) 39 ppm (guinea pig, 4 hr) 418 ppm (rat, 4 hr) [5] | ||
NIOSH (US health exposure limits): | |||
PEL (Permissible) | C 1 ppm (3 mg/m3) [2] | ||
REL (Recommended) | C 1 ppm (3 mg/m3) [2] | ||
IDLH (Immediate danger) | 25 ppm [2] | ||
Safety data sheet (SDS) | ICSC | ||
Related compounds | |||
Other anions | |||
Other cations | |||
Related compounds | Boron monofluoride | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). |
Boron trifluoride is the inorganic compound with the formula BF3. This pungent, colourless, and toxic gas forms white fumes in moist air. It is a useful Lewis acid and a versatile building block for other boron compounds.
The geometry of a molecule of BF3 is trigonal planar. Its D3h symmetry conforms with the prediction of VSEPR theory. The molecule has no dipole moment by virtue of its high symmetry. The molecule is isoelectronic with the carbonate anion, CO2−3.
BF3 is commonly referred to as "electron deficient," a description that is reinforced by its exothermic reactivity toward Lewis bases.
In the boron trihalides, BX3, the length of the B–X bonds (1.30 Å) is shorter than would be expected for single bonds, [7] and this shortness may indicate stronger B–X π-bonding in the fluoride. A facile explanation invokes the symmetry-allowed overlap of a p orbital on the boron atom with the in-phase combination of the three similarly oriented p orbitals on fluorine atoms. [7] Others point to the ionic nature of the bonds in BF3. [8]
BF3 is manufactured by the reaction of boron oxides with hydrogen fluoride:
Typically the HF is produced in situ from sulfuric acid and fluorite (CaF2). [9] Approximately 2300-4500 tonnes of boron trifluoride are produced every year. [10]
For laboratory scale reactions, BF3 is usually produced in situ using boron trifluoride etherate, which is a commercially available liquid.[ how? ]
Laboratory routes to the solvent-free materials are numerous. A well documented route involves the thermal decomposition of diazonium salts of [BF4]−: [11]
It forms by treatment of a mixture boron trioxide and sodium tetrafluoroborate with sulfuric acid: [12]
Alternatively, boron tribromide converts various organofluorine compounds to organobromines, evolving the trifluoride gas: [13]
Anhydrous boron trifluoride has a boiling point of −100.3 °C and a critical temperature of −12.3 °C, so that it can be stored as a refrigerated liquid only between those temperatures. Storage or transport vessels should be designed to withstand internal pressure, since a refrigeration system failure could cause pressures to rise to the critical pressure of 49.85 bar (4.985 MPa). [14]
Boron trifluoride is corrosive. Suitable metals for equipment handling boron trifluoride include stainless steel, monel, and hastelloy. In presence of moisture it corrodes steel, including stainless steel. It reacts with polyamides. Polytetrafluoroethylene, polychlorotrifluoroethylene, polyvinylidene fluoride, and polypropylene show satisfactory resistance. The grease used in the equipment should be fluorocarbon based, as boron trifluoride reacts with the hydrocarbon-based ones. [15]
Unlike the aluminium and gallium trihalides, the boron trihalides are all monomeric. They undergo rapid halide exchange reactions:
Because of the facility of this exchange process, the mixed halides cannot be obtained in pure form.
Boron trifluoride is a versatile Lewis acid that forms adducts with such Lewis bases as fluoride and ethers:
Tetrafluoroborate salts are commonly employed as non-coordinating anions. The adduct with diethyl ether, boron trifluoride diethyl etherate, or just boron trifluoride etherate, (BF3·O(CH2CH3)2) is a conveniently handled liquid and consequently is widely encountered as a laboratory source of BF3. [16] Another common adduct is the adduct with dimethyl sulfide (BF3·S(CH3)2), which can be handled as a neat liquid. [17]
All three lighter boron trihalides, BX3 (X = F, Cl, Br) form stable adducts with common Lewis bases. Their relative Lewis acidities can be evaluated in terms of the relative exothermicities of the adduct-forming reaction. Such measurements have revealed the following sequence for the Lewis acidity:
This trend is commonly attributed to the degree of π-bonding in the planar boron trihalide that would be lost upon pyramidalization of the BX3 molecule. [18] which follows this trend:
The criteria for evaluating the relative strength of π-bonding are not clear, however. [7] One suggestion is that the F atom is small compared to the larger Cl and Br atoms. As a consequence, the bond length between boron and the halogen increases while going from fluorine to iodine hence spatial overlap between the orbitals becomes more difficult. The lone pair electron in pz of F is readily and easily donated and overlapped to empty pz orbital of boron. As a result, the pi donation of F is greater than that of Cl or Br.
In an alternative explanation, the low Lewis acidity for BF3 is attributed to the relative weakness of the bond in the adducts F3B−L. [19] [20]
Yet another explanation might be found in the fact that the pz orbitals in each higher period have an extra nodal plane and opposite signs of the wave function on each side of that plane. This results in bonding and antibonding regions within the same bond, diminishing the effective overlap and so lowering the π-donating blockage of the acidity. [21]
Boron trifluoride reacts with water to give boric acid and fluoroboric acid. The reaction commences with the formation of the aquo adduct, H2O−BF3, which then loses HF that gives fluoroboric acid with boron trifluoride. [22]
The heavier trihalides do not undergo analogous reactions, possibly due to the lower stability of the tetrahedral ions [BCl4]− and [BBr4]−. Because of the high acidity of fluoroboric acid, the fluoroborate ion can be used to isolate particularly electrophilic cations, such as diazonium ions, that are otherwise difficult to isolate as solids.
Boron trifluoride is most importantly used as a reagent in organic synthesis, typically as a Lewis acid. [10] [23] Examples include:
Other, less common uses for boron trifluoride include:
Boron trifluoride was discovered in 1808 by Joseph Louis Gay-Lussac and Louis Jacques Thénard, who were trying to isolate "fluoric acid" (i.e., hydrofluoric acid) by combining calcium fluoride with vitrified boric acid. The resulting vapours failed to etch glass, so they named it fluoboric gas. [27] [28]
A Lewis acid (named for the American physical chemist Gilbert N. Lewis) is a chemical species that contains an empty orbital which is capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base, because it can donate its lone pair of electrons. Trimethylborane [(CH3)3B] is a Lewis acid as it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, a lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3. The terminology refers to the contributions of Gilbert N. Lewis.
Diborane(6), commonly known as diborane, is the chemical compound with the formula B2H6. It is a highly toxic, colorless, and pyrophoric gas with a repulsively sweet odor. Given its simple formula, borane is a fundamental boron compound. It has attracted wide attention for its electronic structure. Several of its derivatives are useful reagents.
The Brønsted–Lowry theory (also called proton theory of acids and bases) is an acid–base reaction theory which was first developed by Johannes Nicolaus Brønsted and Thomas Martin Lowry independently in 1923. The basic concept of this theory is that when an acid and a base react with each other, the acid forms its conjugate base, and the base forms its conjugate acid by exchange of a proton (the hydrogen cation, or H+). This theory generalises the Arrhenius theory.
Antimony pentafluoride is the inorganic compound with the formula SbF5. This colourless, viscous liquid is a strong Lewis acid and a component of the superacid fluoroantimonic acid, formed upon mixing liquid HF with liquid SbF5 in 1:1 ratio. It is notable for its strong Lewis acidity and the ability to react with almost all known compounds.
Boron trichloride is the inorganic compound with the formula BCl3. This colorless gas is a reagent in organic synthesis. It is highly reactive towards water.
Boron tribromide, BBr3, is a colorless, fuming liquid compound containing boron and bromine. Commercial samples usually are amber to red/brown, due to weak bromine contamination. It is decomposed by water and alcohols.
Boron trifluoride etherate, strictly boron trifluoride diethyl etherate, or boron trifluoride–ether complex, is the chemical compound with the formula BF3O(C2H5)2, often abbreviated BF3OEt2. It is a colorless liquid, although older samples can appear brown. The compound is used as a source of boron trifluoride in many chemical reactions that require a Lewis acid. The compound features tetrahedral boron coordinated to a diethylether ligand. Many analogues are known, including the methanol complex.
Tetrafluoroborate is the anion BF−
4. This tetrahedral species is isoelectronic with tetrafluoroberyllate (BeF2−
4), tetrafluoromethane (CF4), and tetrafluoroammonium (NF+
4) and is valence isoelectronic with many stable and important species including the perchlorate anion, ClO−
4, which is used in similar ways in the laboratory. It arises by the reaction of fluoride salts with the Lewis acid BF3, treatment of tetrafluoroboric acid with base, or by treatment of boric acid with hydrofluoric acid.
Selenium tetrafluoride (SeF4) is an inorganic compound. It is a colourless liquid that reacts readily with water. It can be used as a fluorinating reagent in organic syntheses (fluorination of alcohols, carboxylic acids or carbonyl compounds) and has advantages over sulfur tetrafluoride in that milder conditions can be employed and it is a liquid rather than a gas.
Tris(pentafluorophenyl)borane, sometimes referred to as "BCF", is the chemical compound (C6F5)3B. It is a white, volatile solid. The molecule consists of three pentafluorophenyl groups attached in a "paddle-wheel" manner to a central boron atom; the BC3 core is planar. It has been described as the “ideal Lewis acid” because of its high thermal stability and the relative inertness of the B-C bonds. Related fluoro-substituted boron compounds, such as those containing B−CF3 groups, decompose with formation of B-F bonds. Tris(pentafluorophenyl)borane is thermally stable at temperatures well over 200 °C, resistant to oxygen, and water-tolerant.
Boron compounds are compounds containing the element boron. In the most familiar compounds, boron has the formal oxidation state +3. These include oxides, sulfides, nitrides, and halides.
Bromine compounds are compounds containing the element bromine (Br). These compounds usually form the -1, +1, +3 and +5 oxidation states. Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X2/X− couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.
Fluoroboric acid or tetrafluoroboric acid is an inorganic compound with the simplified chemical formula H+[BF4]−. Solvent-free tetrafluoroboric acid has not been reported. The term "fluoroboric acid" usually refers to a range of compounds including hydronium tetrafluoroborate, which are available as solutions. The ethyl ether solvate is also commercially available, where the fluoroboric acid can be represented by the formula [H( 2O)n]+[BF4]−, where n is 2.
Thiazyl fluoride, NSF, is a colourless, pungent gas at room temperature and condenses to a pale yellow liquid at 0.4 °C. Along with thiazyl trifluoride, NSF3, it is an important precursor to sulfur-nitrogen-fluorine compounds. It is notable for its extreme hygroscopicity.
Sodium tetrafluoroborate is an inorganic compound with formula NaBF4. It is a salt that forms colorless or white rhombic crystals and is soluble in water (108 g/100 mL) but less soluble in organic solvents.
Cyclopentadienylindium(I), C5H5In, is an organoindium compound containing indium in the +1 oxidation state. Commonly abbreviated to CpIn, it is a cyclopentadienyl complex with a half-sandwich structure. It was the first (1957) low-valent organoindium compound reported.
Boron monofluoride or fluoroborylene is a chemical compound with the formula BF, one atom of boron and one of fluorine. It is an unstable gas, but it is a stable ligand on transition metals, in the same way as carbon monoxide. It is a subhalide, containing fewer than the normal number of fluorine atoms, compared with boron trifluoride. It can also be called a borylene, as it contains boron with two unshared electrons. BF is isoelectronic with carbon monoxide and dinitrogen; each molecule has 14 electrons.
Borane is an inorganic compound with the chemical formula BH
3. Because it tends to dimerize or form adducts, borane is very rarely observed. It normally dimerizes to diborane in the absence of other chemicals. It can be observed directly as a continuously produced, transitory, product in a flow system or from the reaction of laser ablated atomic boron with hydrogen.
Platinum tetrafluoride is the inorganic compound with the chemical formula PtF
4. In the solid state, the compound features platinum(IV) in octahedral coordination geometry.
Tetrahalodiboranes are a class of diboron compounds with the formula B2X4. These compounds were first discovered in the 1920s, but, after some interest in the middle of the 20th century, were largely ignored in research. Compared to other diboron compounds, tetrahalodiboranes are fairly unstable and historically have been difficult to prepare; thus, their use in synthetic chemistry is largely unexplored, and research on tetrahalodiboranes has stemmed from fundamental interest in their reactivity. Recently, there has been a resurgence in interest in tetrahalodiboranes, particularly in diboron tetrafluoride as a reagent to promote doping of silicon with B+ for use in semiconductor devices.
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