Hydrogen fluoride
Hydrogen fluoride (HF) is a diatomic inorganic compound composed of hydrogen and fluorine, existing as a colorless gas or fuming liquid under standard conditions with a strong, irritating odor. With a molecular weight of 20.01 g/mol, it exhibits an anomalously high boiling point of 19.5 °C (67.1 °F) and melting point of −83.6 °C (−118.4 °F) due to extensive hydrogen bonding, despite being the lightest hydrogen halide.[1][2] Highly soluble in water, it forms hydrofluoric acid solutions that are among the strongest known weak acids, with a pKa of approximately 3.17.[3] Hydrogen fluoride is produced industrially by reacting calcium fluoride (fluorite) with sulfuric acid, yielding HF gas that is then purified and liquefied for storage and transport in steel cylinders.[1] Its primary applications include the manufacture of fluorocarbons such as refrigerants and propellants, production of aluminum fluoride used in the electrolytic smelting of aluminum, and glass etching for decorative or industrial purposes like silicon wafer processing in semiconductors.[1] Additionally, it serves as a catalyst in petroleum alkylation to produce high-octane gasoline and is used in the synthesis of pharmaceuticals and agrochemicals.[4] Despite its industrial importance, hydrogen fluoride poses severe health and safety risks, acting as a highly corrosive agent that causes deep tissue burns by penetrating skin and reacting with calcium and magnesium ions, potentially leading to bone necrosis and systemic toxicity.[5] Inhalation of its vapors irritates the respiratory tract and eyes even at low concentrations (above 3 ppm), and concentrated exposures can result in pulmonary edema, cardiac arrhythmias, or death without prompt treatment using calcium gluconate.[6] It also corrodes glass, silica, and most metals, necessitating specialized handling equipment like polyethylene or Teflon containers.[7]Properties
Physical properties
Hydrogen fluoride has the molecular formula HF and a molar mass of 20.006 g/mol. It appears as a colorless gas at standard conditions or as a colorless, fuming liquid when condensed below its boiling point.[8][3] The compound exhibits a melting point of -83.6 °C and a boiling point of 19.5 °C. The density of the liquid phase is 0.99 g/cm³ at the boiling point. Hydrogen fluoride displays a high dielectric constant of 83.6 at 0 °C, reflecting its polar nature; its viscosity is approximately 0.26 mPa·s at 19.5 °C, lower than that of water (0.89 mPa·s at 20 °C) but with a notably lower surface tension of about 10 mN/m compared to water's 72 mN/m.[8][9] Due to strong hydrogen bonding, hydrogen fluoride undergoes polymerization in both liquid and gaseous phases. In the gas phase, it forms cyclic dimers and higher oligomers, while in the liquid phase, it adopts zigzag chain structures. This association influences its physical behavior, such as elevating the boiling point relative to other hydrogen halides.[10][11] Thermodynamic properties include a heat of vaporization of 30.7 kJ/mol at the boiling point. The specific heat capacity of the gas phase is approximately 29.1 J/mol·K at 298 K, while for the liquid phase, it is around 52 J/mol·K near 0 °C.[12][13] Spectroscopic characteristics feature a prominent infrared absorption band for the H-F stretching vibration at approximately 4000 cm⁻¹ in the monomeric form, shifting to lower wavenumbers in associated species. In nuclear magnetic resonance, the ¹⁹F chemical shift for gaseous monomeric HF is 46.85 ppm relative to SiF₄, and the ¹H shift is about 8.4 ppm relative to external standards.[14][15]| Property | Value | Conditions |
|---|---|---|
| Molecular formula | HF | - |
| Molar mass | 20.006 g/mol | - |
| Appearance | Colorless gas/liquid | Room temp./below b.p. |
| Melting point | -83.6 °C | 1 atm |
| Boiling point | 19.5 °C | 1 atm |
| Density (liquid) | 0.99 g/cm³ | At boiling point |
| Dielectric constant | 83.6 | 0 °C |
| Heat of vaporization | 30.7 kJ/mol | At boiling point |
| Specific heat (gas) | 29.1 J/mol·K | 298 K |
| IR H-F stretch | ~4000 cm⁻¹ | Gas monomer |
| ¹⁹F NMR shift | 46.85 ppm | Gas monomer vs. SiF₄ |
Chemical properties
Hydrogen fluoride (HF) is characterized by its exceptional acidity, which varies significantly between its anhydrous and aqueous forms. In aqueous solution, HF behaves as a weak acid with a pKa of 3.17 at 25 °C, dissociating partially to H⁺ and F⁻, unlike the other hydrogen halides (HCl, HBr, HI), which are strong acids in water. This relative weakness stems from the strong H–F bond, with a bond dissociation energy of 569 kJ/mol at 298 K—the highest among the HX series due to fluorine's high electronegativity (4.0 on the Pauling scale)—and the strong hydrogen bonding of the fluoride ion with water molecules, which stabilizes F⁻ and hinders complete dissociation.[8] In contrast, anhydrous HF is an extremely strong acid, functioning as a protogenic solvent with a Hammett acidity function H₀ of approximately −11, comparable to concentrated sulfuric acid; this level of acidity arises from its ability to readily donate protons in the absence of water's leveling effect.[16] Anhydrous HF exhibits low solubility in non-polar solvents such as hydrocarbons, reflecting its high polarity and tendency to form hydrogen bonds, but it is fully miscible with water and polar organic solvents like ethanol and diethyl ether. With water, it forms a binary azeotrope consisting of ~38% HF by weight, boiling at 112 °C. Thermally, HF is stable under ambient conditions and does not readily decompose to its elements, 2 HF → H₂ + F₂, without very high temperatures or specific conditions such as catalysis; this underscores its stability in typical storage and handling scenarios due to the robust H–F bond.[17] A notable chemical reactivity of HF is its corrosive action on glass and silica-containing materials, where it etches silicon dioxide via the reaction SiO₂ + 6 HF → H₂SiF₆ + 2 H₂O, forming fluosilicic acid; this property necessitates storage in fluoropolymer or metal containers like steel or aluminum. In its pure liquid form, anhydrous HF undergoes limited autoionization: 3 HF ⇌ H₂F⁺ + HF₂⁻, producing ionic species (fluoronium and bifluoride ions) that confer a low electrical conductivity of about 10⁻⁵ S/cm at 0 °C, with an autoionization constant K ≈ 10⁻¹². Among the hydrogen halides, HF occupies a unique position, as its intermolecular hydrogen bonding—stronger than van der Waals forces in HCl, HBr, and HI—dominates its liquid and solid-state properties, influencing viscosity, boiling point, and reactivity patterns.[2]History
Discovery and early studies
The initial preparation of hydrogen fluoride is attributed to German glass cutter Heinrich Schwanhard in 1670, who reacted fluorspar (CaF₂) with sulfuric acid to produce fumes capable of etching glass, though he did not isolate the compound. Schwanhard noted the extreme danger of inhaling these corrosive vapors, marking the first documented use of the substance in a practical application.[18] In 1771, Swedish chemist Carl Wilhelm Scheele advanced the understanding of the compound by systematically heating fluorspar with concentrated sulfuric acid, yielding an impure form of hydrofluoric acid that he termed "fluoric acid." Scheele's experiments highlighted its powerful etching effect on glass vessels, confirming its corrosive nature and distinguishing it from other mineral acids. Although his product contained water, this work represented the first scientific isolation and description of the acid in larger quantities.[18] By 1810, British chemist Humphry Davy conducted electrolysis on samples of hydrofluoric acid, demonstrating that it consisted of hydrogen combined with a novel element analogous to chlorine. Davy named the acid "fluohydric acid" based on this binary composition, shifting the prevailing view from oxide-based theories and laying groundwork for recognizing fluorine as an independent element. His studies, however, came at personal cost, as exposure caused severe burns and health issues.[18] In the mid-19th century, French chemist Edmond Frémy prepared anhydrous hydrogen fluoride gas by distilling potassium bifluoride with sulfuric acid, enabling detailed examination of its properties. Frémy emphasized its intense corrosiveness, capable of attacking metals, glass, and organic tissues, and observed anomalies like its unexpectedly high boiling point (19°C), which hinted at intermolecular associations later understood as hydrogen bonding.[2][19] During the 1830s, investigations by chemists including the Knox brothers revealed hydrogen fluoride's gaseous state at room temperature and its tendency to form dense, irritating fumes in air due to hydrolysis. These studies, building on earlier work, underscored the compound's volatility and toxicity, with experimenters suffering near-fatal poisoning from inhalation.[18]Industrial development
Commercial production of hydrogen fluoride emerged in the early 20th century, relying on the reaction of concentrated sulfuric acid with fluorspar (calcium fluoride, CaF₂) to generate HF gas, which is then purified by distillation. This method allowed for scalable output to meet growing industrial needs. The process's efficiency and the abundance of fluorspar deposits facilitated integration into chemical manufacturing, marking the transition from laboratory-scale synthesis to economic viability.[20] The 1940s represented a pivotal expansion driven by wartime demands, particularly the Manhattan Project's requirement for large-scale HF to produce uranium hexafluoride (UF₆) for gaseous diffusion-based uranium enrichment. Facilities scaled up dramatically, with Harshaw Chemical alone delivering over 1,615 tons of UF₆ derived from HF to support the effort. Researchers at institutions like Purdue University optimized processes leveraging commercially available HF, underscoring its role as an accessible precursor in high-stakes applications. This period catalyzed infrastructure investments that propelled HF into a cornerstone of the fluorochemical sector.[21][22] Following World War II, HF production surged in the 1950s amid booming demand for fluorocarbons, including refrigerants like chlorofluorocarbons (CFCs) and materials such as polytetrafluoroethylene (PTFE, known as Teflon). DuPont, for instance, ramped up Teflon commercialization in 1950, with output exceeding 2 million pounds annually by the mid-decade, heavily reliant on HF as a fluorinating agent. This era's innovations in polymer and refrigerant synthesis solidified HF's industrial footprint, fueling applications in consumer goods and electronics.[23][24] By the 2020s, global HF production capacity approached 2.4 million metric tons annually, reflecting sustained growth in sectors like electronics and pharmaceuticals. However, this expansion followed heightened safety measures implemented after 1980s incidents, notably the 1987 Texas refinery release of approximately 24,000 kg of HF, which exposed nearby communities and prompted the U.S. Occupational Safety and Health Administration (OSHA) to enact Process Safety Management standards in 1992 specifically addressing anhydrous HF hazards. These regulations mandated risk assessments, emergency planning, and engineering controls to mitigate releases in high-risk industries like petroleum refining.[25][6][26]Synthesis and production
Laboratory synthesis
The classic laboratory method for preparing hydrogen fluoride involves heating powdered calcium fluoride (CaF₂) with concentrated sulfuric acid (H₂SO₄) in a corrosion-resistant apparatus, such as platinum or Teflon-lined vessels, to generate HF gas. The reaction proceeds as follows: \ce{CaF2 + H2SO4 -> 2HF + CaSO4} This process typically yields HF of 90-95% purity and is suitable for small-scale production ranging from grams to kilograms.[27] Alternative methods include thermal decomposition of potassium bifluoride (KHF₂) by heating to produce anhydrous HF gas: \ce{KHF2 -> KF + HF}, or the reaction of sodium fluoride (NaF) with concentrated hydrochloric acid (HCl) to generate HF gas in situ for immediate use in research applications.[28] Purification of the crude HF is achieved through distillation under anhydrous conditions to eliminate water and volatile impurities like SO₂ or SiF₄; apparatus constructed from corrosion-resistant materials such as copper or Monel metal is employed to withstand HF's aggressive nature.[29] Due to HF's extreme corrosivity, toxicity, and ability to cause severe burns, all synthesis must occur in a properly functioning fume hood with adequate ventilation, and calcium oxide (CaO) or similar fluoride scavengers should be on hand for neutralizing spills or residues. Personal protective equipment, including HF-specific gloves and eyewear, is essential.[30]Industrial production
The primary industrial production of hydrogen fluoride (HF) involves the reaction of acid-grade fluorspar (CaF₂, typically ≥97% purity) with concentrated sulfuric acid (H₂SO₄, 96-98%) in a rotary kiln reactor.[31][32] This endothermic process occurs at temperatures of 225-265°C, generating HF gas (95-99% purity) and calcium sulfate (CaSO₄) as a solid byproduct according to the equation: \mathrm{CaF_2 + H_2SO_4 \rightarrow CaSO_4 + 2HF} [33][31] The HF gas is then separated from the reaction mixture, while the gypsum byproduct is filtered, cooled, and often neutralized for use in construction materials or land application.[32][34] Commercial plants typically operate at capacities ranging from 10,000 to 50,000 tons of HF per year, with global production approximately 1.2 million metric tons as of 2024, primarily driven by demand in fluorochemical manufacturing.[35] Alternative methods, such as recovery from fluorosilicic acid (a byproduct of phosphate fertilizer production from apatite ore) or recycling of fluoroorganic wastes, account for less than 10% of total output due to lower scalability and higher costs compared to the fluorspar route.[31][36] For purification, the crude HF gas is absorbed in water to yield aqueous hydrofluoric acid (49-70% concentration) for direct industrial use, or further processed via fractional distillation to produce anhydrous HF with purity levels of 99.8% or higher.[31][37] In the distillation step, impurities like water, sulfuric acid, and trace metals (e.g., arsenic) are removed under controlled conditions to meet specifications for sensitive applications such as semiconductor etching.[37][38]Reactions and chemical behavior
Bonding and structure
Hydrogen fluoride (HF) is a linear diatomic molecule characterized by a single covalent bond between the hydrogen and fluorine atoms. The experimental bond length of the H–F bond is 0.917 Å, reflecting the strong attraction due to fluorine's high electronegativity. This polarity imparts a significant dipole moment of 1.86 D to the molecule, making HF the most polar among the hydrogen halides.[39][40] The electronic structure of HF is well-described by molecular orbital theory, where the primary sigma bonding orbital arises from the end-to-end overlap of the hydrogen 1s atomic orbital and the fluorine 2p_z atomic orbital along the molecular axis. The fluorine atom contributes three lone pairs in its 2s and 2p_x, 2p_y orbitals, which remain largely non-bonding but play a crucial role in intermolecular interactions. These lone pairs enable HF to form strong hydrogen bonds, the strongest among the HX series (X = F, Cl, Br, I), due to fluorine's compact electron cloud and high electronegativity. In the gas phase, HF predominantly exists as monomers at low pressures but forms dimers at higher concentrations, with the hydrogen-bonded (HF)_2 structure featuring a nearly linear F–H···F arrangement and an F···F distance of approximately 2.5 Å.[41] In the liquid phase, extensive hydrogen bonding leads to the formation of polymeric chains denoted as (HF)_n, with average chain lengths of 6 to 7 molecules, though clusters up to n ≈ 10 are observed. These associations result in a zig-zag configuration of bent hydrogen bonds, explaining HF's anomalously high boiling point of 19.5 °C compared to the trend in hydrogen halides. Quantum mechanical calculations, such as those using Hartree-Fock methods, reproduce the H–F bond strength with dissociation energies around 0.36 hartree (approximately 980 kJ/mol at the equilibrium geometry), though post-Hartree-Fock corrections are needed for experimental accuracy of 565 kJ/mol. The vibrational frequency of the H–F stretch, measured at 4138 cm⁻¹, further underscores the bond's rigidity and strength.[42][40] In the solid state, HF crystallizes in an orthorhombic lattice (space group Cmcm) composed of infinite, unbranched zigzag chains of hydrogen-bonded molecules, with F–H···F angles near 180° within chains and interchain distances governed by van der Waals interactions between fluorine atoms.[43] This polymeric network persists across phases, highlighting the pervasive influence of hydrogen bonding on HF's structure.Reactions with other halides
Hydrogen fluoride exhibits distinct reactivity compared to the other hydrogen halides (HCl, HBr, HI) primarily due to the strong H–F bond and fluorine's high electronegativity. The bond dissociation energies decrease down the group: HF (569 kJ/mol) > HCl (431 kJ/mol) > HBr (366 kJ/mol) > HI (299 kJ/mol), reflecting the increasing atomic size of the halogen and weaker orbital overlap.[44] This trend influences acidity; in the gas phase, acidity increases from HF to HI due to decreasing bond strength, but in aqueous solution, HF is the weakest acid (pKa = 3.17) while HCl (pKa ≈ -7), HBr (pKa ≈ -9), and HI (pKa ≈ -10) are strong acids, as the small, highly basic F⁻ ion forms strong hydrogen bonds that limit dissociation.[44][45] Thermal stability follows the bond strength trend, with HF being the most stable and least prone to decomposition into elements upon heating, unlike HI which decomposes readily. HF is also the least volatile, with a boiling point of 19.5°C due to extensive hydrogen bonding, contrasting with the lower boiling points of HCl (-85.1°C), HBr (-66.8°C), and HI (-35.4°C) where London dispersion forces dominate.[44][44] In redox behavior, HF is notably inert to oxidation, resisting reactions that would liberate fluorine, whereas HI acts as a strong reducing agent, readily oxidizing to I₂ (e.g., with permanganate or air) due to the weak H–I bond facilitating electron donation. This reducing power increases from HF to HI, as weaker bonds allow easier cleavage and oxidation of the halide. Halogen exchange reactions highlight HF's ability to displace less electronegative halogens from their salts, driven by fluorine's preference for stronger bonds. For instance, HF reacts with silver chloride to form silver fluoride and HCl: AgCl + HF → AgF + HCl, a method used historically to prepare anhydrous AgF.[46] A key difference arises in anhydrous versus aqueous environments: anhydrous HF forms the bifluoride ion [HF₂]⁻ through strong hydrogen bonding (F–H–F), a symmetric species absent in the other halides which do not exhibit such dimerization.[47]| Property | HF | HCl | HBr | HI |
|---|---|---|---|---|
| Boiling point (°C) | 19.5 | -85.1 | -66.8 | -35.4 |
| pKa (in water) | 3.17 | ≈ -7 | ≈ -9 | ≈ -10 |
| Reactivity with metals (e.g., Zn) | Slow (weak acid) | Moderate | Fast | Very fast (strongest acid) |