Fact-checked by Grok 2 weeks ago

Hydrofluoric acid

Hydrofluoric acid is a colorless, fuming of (), a with the and a molecular weight of 20.01 g/. It exists as a gas at but is commonly handled as a up to 70% concentration, exhibiting a strong irritating odor with an odor threshold of 0.03 mg/m³. Despite being a weak acid with a of 3.17, it forms strong hydrogen bonds due to fluorine's high , resulting in a higher of 19.5°C compared to other hydrogen halides. Chemically, hydrofluoric acid is highly reactive and corrosive, particularly to silica-containing materials like , which it dissolves to form and . It reacts exothermically with metals to produce gas and with bases to form salts, but its low in belies its penetrating power into tissues. HF is one of the strongest known acids, capable of protonating compounds, while aqueous solutions are miscible with and denser at about 1.15 g/mL for 49% concentration at 20°C. Hydrofluoric acid is industrially vital, serving as a precursor for fluorochemicals such as refrigerants, , and pharmaceuticals, with about 60% of production used in synthesis. Key applications include and etching, aluminum production, refining, pickling, and petroleum as a catalyst. It is also employed in computer screens, fluorescent bulbs, and high-octane . However, hydrofluoric acid poses severe health risks due to its rapid penetration of and tissues, causing , severe burns, and systemic like and cardiac arrhythmias, even from small exposures. can lead to respiratory damage and , while ingestion affects major organs and may be fatal; exposure limits include an 8-hour TWA of 3 (OSHA PEL). Prevention requires chemical-resistant protective equipment, and treatment involves immediate water rinsing followed by application to bind ions.

Properties

Physical properties

Hydrofluoric acid appears as a colorless, fuming liquid with a strong irritating , particularly at concentrations above 48%. The of anhydrous hydrofluoric acid is 19.5 °C at standard pressure, whereas aqueous solutions exhibit higher boiling points that vary with concentration; for instance, the binary consisting of 38% HF boils at 112 °C. The of a % is 1.15 g/cm³ at 20 °C, with values decreasing as temperature increases. Hydrofluoric acid is infinitely miscible with , , and , and it forms azeotropic mixtures, such as the one with at 38% HF composition. Its high volatility results in significant , for example, 25 mmHg at 20 °C for a % , contributing to the characteristic fuming. Concentrated forms of hydrofluoric acid exhibit thermal stability up to approximately 100 °C, above which they decompose, releasing hydrogen fluoride gas.

Chemical properties

Hydrofluoric acid is the aqueous solution of hydrogen fluoride, denoted as HF(aq), while the anhydrous form has the molecular formula HF. In its pure liquid state, anhydrous HF exhibits extensive hydrogen bonding due to the high electronegativity of fluorine, resulting in a structure composed of zig-zag polymeric chains of HF molecules. The H-F bond is notably strong, with a of 565 kJ/mol, attributed to fluorine's high . Despite this, HF readily forms the , [HF₂]⁻, which contains the strongest known and stabilizes the species in solutions with excess . A key reactive property of HF is its ability to etch silica-containing materials, as demonstrated by the reaction with :
\ce{SiO2 + 4HF -> SiF4 + 2H2O}
This process releases gas and dissolves glass and , making HF uniquely aggressive toward siliceous substrates compared to other acids. HF corrodes most metals by forming corresponding metal fluorides, though it does not affect noble metals such as and . With silicates, it forms stable complexes, including (H₂SiF₆), via reactions like:
\ce{SiO2 + 6HF -> H2SiF6 + 2H2O}
Under controlled conditions, such as in the presence of catalysts, HF also reacts with compounds to yield fluorocarbons, facilitating the introduction of into carbon-based structures. Anhydrous HF serves as a non-aqueous for many ionic compounds owing to its autoprotolysis equilibrium:
\ce{3HF <=> H2F+ + HF2-}
This self-ionization produces an extremely acidic medium (H₀ ≈ -15). Unlike the other hydrogen halides (HCl, HBr, ), which fully dissociate in , HF remains partially undissociated due to hydrogen bonding that stabilizes the undissociated form, rendering it weaker in aqueous media but more potent as an acid in non-protic solvents.

Acidity

Hydrofluoric acid () is classified as a weak in dilute aqueous solutions, characterized by a of 3.17 at 25 °C and an K_a = 6.8 \times 10^{-4}. This indicates partial , where only about 8% of HF molecules ionize in a 0.1 M . The is represented as: \text{HF} \rightleftharpoons \text{H}^+ + \text{F}^- The limited arises from the strong H–F bond (bond energy of 569 kJ/), driven by fluorine's high , combined with the small size of the F⁻ , which leads to strong but entropically unfavorable hydrogen bonding in its hydration shell. The standard change for , \Delta G^\circ \approx 18.1 kJ/ at 25 °C, reflects this endergonic process, where enthalpic gains from bond breaking are offset by a negative change due to the structured of ions, particularly F⁻. Compared to other hydrogen halides (HX: HCl, HBr, ), HF is the weakest acid because the H–F bond strength dominates over the stabilization of the conjugate base, unlike larger halides where weaker bonds and better ion enhance acidity ( values: HCl ≈ -7, HBr ≈ -9, ≈ -10). In concentrated aqueous solutions (> ~5 M), HF's effective acidity increases through ion pairing, forming ions [HF₂]⁻, which reduces the activity of free F⁻ and shifts the equilibrium toward greater proton release, making it behave more strongly. In non-aqueous solvents or its form, HF functions as a strong acid and , fully ionizing due to the absence of water's and its ability to solvate protons efficiently. The /F⁻ buffer system maintains stable around 3.2 and ensures availability of free ions for complexation or reactions, despite the weak initial dissociation; this buffering capacity is utilized in for controlling .

Production

Industrial production

Hydrofluoric acid is produced on an industrial scale primarily through the reaction of acid-grade fluorspar (, CaF₂, containing at least 97% CaF₂) with concentrated . The process involves heating the mixture to 200–300 °C in rotary kilns or, less commonly, reactors, where the endothermic reaction generates gas and as a solid byproduct:
\ce{CaF2 + H2SO4 -> 2HF + CaSO4}
This reaction requires external heating to maintain temperature and proceeds over 30–60 minutes, with the solid () separated as waste material that must be managed due to its volume and environmental impact. The process is energy-intensive owing to the endothermic nature of the reaction, which absorbs approximately 1,400–2,500 kJ per kg of HF produced. The gaseous HF is then captured, condensed, and purified via to remove impurities such as (SO₂), residual (H₂SO₄), water, and trace fluorides like SiF₄ or BF₃. columns separate low-boiling impurities (e.g., SO₂) in overhead streams and high-boiling ones (e.g., H₂SO₄) in bottoms, yielding high-purity HF that can be hydrated to form aqueous solutions or further processed to form through under controlled conditions to achieve 99.9% purity or higher. As of 2024, global production capacity for hydrofluoric acid exceeds 3 million metric tons per year, with actual output around 2–3 million tons annually to meet demand. Major producing countries include (over 60% of exports), , and the , where capacity is approximately 220,000 metric tons per year as of 2022. Alternative production routes, though less prevalent than the fluorspar method, include derivation from fluoroapatite in phosphate rock processing, where fluorosilicic acid (H₂SiF₆) byproduct is decomposed with to yield HF, and fluoride wastes from aluminum operations. These methods contribute a smaller share of supply but support by utilizing industrial byproducts. Industrial hydrofluoric acid is typically marketed as 48–70% aqueous solutions for general use, while HF is produced via additional for specialized applications requiring high purity.

Laboratory preparation

In laboratory settings, hydrofluoric acid is prepared on a small scale, typically ranging from grams to liters, to support or educational applications, with strict emphasis on systems and neutralization setups to manage hazardous vapors. A common method involves the of (KHF₂), where the solid is heated to release gas, which is then absorbed into to form the aqueous acid. \text{KHF}_2 \xrightarrow{\text{heat}} \text{KF} + \text{HF (g)} This decomposition occurs quantitatively at temperatures above 200°C, producing pure HF gas that can be collected and dissolved in distilled water under controlled conditions. The process requires apparatus made from corrosion-resistant materials such as platinum or polytetrafluoroethylene (PTFE, commonly known as Teflon) to prevent reaction with the highly corrosive HF, and distillation is often conducted under reduced pressure to lower boiling points and minimize equipment stress. An alternative approach utilizes the reaction of (NaF) with concentrated in glass-free apparatus to generate gas, followed by or . $2\text{NaF} + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + 2\text{HF} This method avoids the formation of insoluble byproducts like encountered in other preparations and is suitable for producing small quantities of , though it demands similar safety measures including plastic or metal equipment to avert . For obtaining anhydrous HF, aqueous solutions are dried by mixing with concentrated , which forms a water- azeotrope that is distilled off, leaving pure HF that can then be redistilled under reduced pressure. Historically, early laboratory preparations involved distilling HF directly from fluorspar (CaF₂) with , a method pioneered by Frémy in the mid-19th century through heating , but it has been largely supplanted due to significant hazards from uncontrolled gas evolution and equipment failure.

Uses

Synthesis of organofluorine compounds

Hydrofluoric acid () plays a pivotal role in the of organofluorine compounds by facilitating the formation of carbon-fluorine s through direct fluorination, , or as a in key transformations. These processes are essential for producing materials used in pharmaceuticals, refrigerants, and polymers, where the strong C-F imparts stability and unique properties. HF, in particular, serves as both a fluorinating agent and a due to its superacidic nature, enabling reactions that are challenging in conventional media. In the production of chlorofluorocarbons (CFCs), is employed in the vapor-phase catalytic displacement of atoms in chlorocarbons, such as the reaction of with (CCl₄) to yield (CFCl₃) and HCl: \mathrm{CCl_4 + HF \rightarrow CFCl_3 + HCl} This process, a variant of the Swarts reaction, was historically significant for manufacturing refrigerants and propellants before regulatory restrictions. The reaction typically requires catalysts like chlorofluorides to enhance selectivity and efficiency. A key application involves the synthesis of (PTFE), known commercially as Teflon, whose monomer (TFE) is derived from and HF. (CHCl₃) undergoes stepwise fluorination with HF to form (CHClF₂), which is then pyrolyzed at high temperatures (around 700–900°C) to produce TFE: \mathrm{CHCl_3 + 2HF \rightarrow CHClF_2 + 2HCl}, \quad \mathrm{2CHClF_2 \rightarrow CF_2=CF_2 + 2HCl} This multi-step process highlights HF's role in introducing fluorine atoms to aliphatic precursors, yielding a renowned for its chemical inertness and non-stick properties. In pharmaceutical synthesis, HF is indirectly involved in the Balz-Schiemann reaction, a classical method for preparing aryl fluorides from aromatic amines. The process begins with diazotization of the amine to form an aryldiazonium salt, followed by treatment with tetrafluoroboric acid (HBF₄, prepared from HF and ) to isolate the tetrafluoroborate salt (ArN₂⁺ BF₄⁻). then affords the aryl fluoride: \mathrm{ArN_2^+ BF_4^- \rightarrow ArF + N_2 + BF_3} This reaction is widely used to introduce fluorine into drug scaffolds, enhancing metabolic stability, as seen in compounds like fluoroquinolone antibiotics. Yields typically range from 40–70%, with modern variants improving efficiency through solvent optimization. For refrigerants, is crucial in producing hydrofluorocarbons (HFCs), such as (HFC-134a), a common replacement. The synthesis starts with the hydrofluorination of (ClCH=CCl₂) using , often in a catalyzed vapor-phase , proceeding through intermediates like 1-chloro-1,1,2-trifluoroethane to the final product: \mathrm{ClCH=CCl_2 + 4HF \rightarrow CF_3CH_2F + 3HCl} This multi-stage reaction, developed in the late 1980s, achieves high selectivity (>95%) under controlled conditions and supports large-scale production for . Anhydrous HF also acts as a for reactions, particularly in systems combined with Lewis acids like SbF₅, enabling the fluorination of aromatic and aliphatic substrates via or addition to unsaturated bonds. For instance, it facilitates the direct fluorination of to under electrochemical conditions, where HF provides both the medium and source. These methods are valued for their ability to introduce at specific positions, though they require specialized equipment due to HF's corrosivity. The transition from CFCs to HFCs was driven by the 1987 , which phased out ozone-depleting substances, prompting the to develop HF-based processes for HFCs that avoid while maintaining thermodynamic performance. This shift increased global HF demand for organofluorine production, with HFC-134a becoming a flagship example of environmentally adapted synthesis.

Production of inorganic fluorides

Hydrofluoric acid (HF) serves as a key reagent in the industrial synthesis of various inorganic fluorides, enabling the fluorination of metal and non-metal compounds for applications in aluminum production, , nuclear fuels, and . These processes typically involve neutralization or direct reaction of HF with oxides, hydroxides, or carbonates, often under controlled conditions to manage the acid's corrosivity and ensure high purity. The resulting fluorides are essential in ceramics for fluxing agents, in for precursors, and in catalysts for industrial reactions. Aluminum fluoride (AlF₃) is produced on a large scale by neutralizing hydrofluoric acid with aluminum hydroxide, following the reaction: \ce{3HF + Al(OH)3 -> AlF3 + 3H2O} This wet process yields a product used primarily as an additive in aluminum via the Hall-Héroult method, where it lowers the of and improves conductivity. Production is predominantly from HF-based routes, supporting the aluminum industry's demand for high-purity grades. Sodium fluoride (NaF) is manufactured through the neutralization of hydrofluoric acid with or : \ce{HF + NaOH -> NaF + H2O} The resulting salt is centrifuged, washed, and dried to achieve pharmaceutical or technical grades. It finds widespread use in to prevent dental caries and in formulations as an anti-cavity agent, with production emphasizing low-impurity variants for applications. Cryolite (Na₃AlF₆) is synthesized by first forming hexafluoroaluminate acid from HF and aluminum hydroxide, followed by reaction with : \ce{Al(OH)3 + 3HF -> H3AlF6 + 3H2O}, \quad \ce{H3AlF6 + 3Na2CO3 -> 2Na3AlF6 + 3CO2 + 3H2O} This synthetic substitutes for scarce natural deposits and is vital in the Hall-Héroult process as the primary solvent for alumina dissolution during aluminum . Industrial production focuses on granular forms to optimize electrolytic performance and reduce in smelters. Uranium tetrafluoride (UF₄), a precursor for , is prepared by hydrofluorination of with HF gas in a : \ce{UO2 + 4HF -> UF4 + 2H2O} The process operates under dilution to facilitate gas-solid contact and remove , yielding green salt (UF₄) that is subsequently converted to UF₆ for enrichment. This method ensures high conversion efficiency in nuclear fuel cycles, with stringent controls for radiological safety. Silicon tetrafluoride (SiF₄) is generated as a volatile product from the reaction of hydrofluoric acid with , often during or processes: \ce{SiO2 + 4HF -> SiF4 + 2H2O} Captured as a in , SiF₄ is recycled by to recover HF and silica, minimizing waste in production. Its gaseous nature allows easy purification for use in and synthesis. Nitrogen trifluoride (NF₃) for is produced via of a molten mixture of , , and hydrofluoric acid (NH₄F–KF–HF), achieving purities up to 99.99% essential for in fabrication. This electrolytic route avoids direct fluorination hazards and supports the demand for ultra-high-purity gases in .

Etching and cleaning

Hydrofluoric acid plays a critical role in manufacturing, where buffered hydrofluoric acid (BHF)—a mixture of hydrofluoric acid and (NH₄F)—is used for the isotropic of (SiO₂) layers on microchips. This process selectively removes oxide films to define circuit patterns, with typical etch rates for thermal SiO₂ reaching approximately 100 nm/min in a 5:1 BHF solution (5 parts 40% NH₄F to 1 part 49% ). The buffering stabilizes the pH and concentration, ensuring uniform and preventing undercutting of masks, which is essential for precise fabrication in integrated circuits. In glass processing, hydrofluoric acid etches surfaces to produce frosted or finishes for decorative items, , and architectural elements, typically using concentrations of 5-10% applied via or formulations to control the depth and uniformity of the etch. This dissolves the silica network in (SiO₂ + 4HF → SiF₄ + 2H₂O), creating a light-diffusing without compromising structural integrity. For industrial cleaning, hydrofluoric acid removes and scales from metals, as in the Fe₂O₃ + 6HF → 2FeF₃ + 3H₂O, and dissolves silicate-based deposits in systems, often in inhibited formulations to protect underlying alloys. Specific applications include texturing silicon wafers in production to create anti-reflective surfaces, where acidic mixtures containing hydrofluoric acid roughen multicrystalline for improved light trapping and efficiency. In phosphate fertilizer processing, hydrofluoric acid facilitates silica removal from phosphate rock by converting impurities to soluble fluorosilicates, enhancing product purity during wet acidulation. protocols emphasize diluting hydrofluoric acid to 1-5% for and operations to minimize exposure risks, followed by neutralization with hydrated lime (Ca(OH)₂) to form insoluble (CaF₂) for safe disposal: 2HF + Ca(OH)₂ → CaF₂ + 2H₂O. Demand from the , particularly semiconductors and display manufacturing, represents a significant portion of global hydrofluoric acid production.

Petroleum refining

In petroleum refining, hydrofluoric acid (HF) serves as a key catalyst in the alkylation process, which combines isobutane with light olefins such as propylene or butenes to produce high-octane alkylate, primarily branched C8 hydrocarbons like 2,2,4-trimethylpentane (C8H18). This reaction occurs in the liquid phase at moderate temperatures of 20–40 °C and elevated pressures, where HF promotes the formation of carbocations from the olefins, facilitating their combination with isobutane to yield branched-chain products that enhance gasoline octane ratings to 95 or higher. The process is highly selective for desirable isomers, contributing to cleaner-burning, high-performance fuels. The two predominant commercial HF alkylation technologies are the Phillips process and the UOP process, both featuring a reactor where hydrocarbons and HF are intimately mixed, followed by separation in a and HF regeneration through and purification to remove and contaminants. In these systems, HF is continuously recycled, achieving over 99% recovery rates with minimal byproducts such as and n-butane, which are separated via . Worldwide, approximately 50 HF units operate, primarily in the United States where about 48 facilities produce roughly 1 million barrels per day of alkylate, representing a significant portion of global high-octane blending stock. Compared to sulfuric acid (H2SO4) alkylation, HF offers advantages including lower overall acid consumption, reduced sensitivity to temperature variations, easier regeneration without the need for frequent acid replacement, and less equipment corrosion in optimized systems. However, HF's extreme toxicity and potential for vapor cloud formation have prompted a phasedown, driven by heightened regulations in the 2010s, such as EPA risk assessments and state-level mandates for enhanced mitigation like water neutralization systems. These concerns were amplified following the 1987 Texas City incident, where a major HF release from a Marathon refinery exposed thousands and led to stricter environmental controls and inventory limits at HF units. Emerging alternatives to liquid HF include solid acid catalysts, which eliminate handling risks associated with corrosive acids and are under continued , though broader commercial adoption remains limited as of 2025. In 2025, environmental groups petitioned the U.S. EPA to prohibit HF use in refineries under the Toxic Substances Control Act due to ongoing safety risks, adding to regulatory scrutiny.

Health and safety

Toxicity and hazards

Hydrofluoric acid () is exceptionally hazardous due to its dual corrosive and systemic toxic effects, primarily from the penetration and of fluoride ions into biological s. Upon dermal contact, HF rapidly diffuses through the skin and underlying tissues, often without immediate pain, allowing deep penetration before occurs. The released fluoride ions (F⁻) strongly bind to divalent cations such as calcium (Ca²⁺) and magnesium (Mg²⁺), precipitating them as insoluble salts and depleting essential electrolytes in cells and extracellular fluids. This leads to cellular disruption, , and progressive tissue destruction that can extend to if untreated. Systemic effects from significant absorption, particularly dermal exposures exceeding 2-3% of total with concentrated solutions, include severe and hypomagnesemia, which can trigger life-threatening cardiac arrhythmias, renal failure, and multi-organ dysfunction. has an acute oral LD50 of approximately 25–50 mg/kg in rats (GHS Category 2), underscoring its high toxicity via and potential for rapid systemic uptake. Inhalation of HF vapor causes immediate irritation to the , potentially progressing to laryngeal spasm, , and delayed-onset due to corrosive damage and fluid accumulation in the lungs. The fuming nature of HF in humid air increases the risk of inadvertent during handling. Ocular exposure results in severe corneal ulceration and opacification, often leading to permanent vision impairment or blindness from rapid penetration and ion binding in eye tissues. Chronic low-level exposure through repeated dermal contact or inhalation can accumulate fluoride, resulting in fluorosis—a condition marked by with bone densification, brittleness, joint stiffness, and potential dental mottling. The carcinogenicity of HF remains unclassifiable; inorganic fluorides employed in , encompassing HF-derived compounds, are designated Group 3 by the International Agency for Research on Cancer (IARC), indicating inadequate evidence for human carcinogenicity. In contrast to other mineral acids like sulfuric or hydrochloric acid, which typically produce coagulative necrosis that confines damage to the surface, HF uniquely causes delayed pain onset—sometimes up to 24 hours post-exposure—and facilitates deeper tissue liquefaction through ongoing fluoride diffusion and electrolyte sequestration. Regulatory exposure limits reflect HF's potency: the Occupational Safety and Health Administration (OSHA) establishes a permissible exposure limit (PEL) of 3 ppm as a ceiling value, while the National Institute for Occupational Safety and Health (NIOSH) defines the immediately dangerous to life or health (IDLH) concentration at 30 ppm.

Treatment and first aid

Immediate for hydrofluoric acid () exposure focuses on rapid to minimize penetration and subsequent administration of calcium-based antidotes to counteract systemic effects by free ions. Workplaces handling must maintain gel as a mandatory component, per occupational safety standards requiring its availability for prompt topical application. For skin exposure, the initial step is to rinse the affected area with copious lukewarm for 15-30 minutes to dilute and remove the acid, followed immediately by application of 2.5-10% gel or jelly, massaged into the skin every 10-15 minutes until pain subsides, as this binds ions locally. In severe cases involving large surface areas (>20% body surface) or signs of systemic toxicity such as , intravenous administration of 10-20 mL of 10% solution is indicated, with continuous electrocardiogram (ECG) monitoring for arrhythmias like prolongation. Eye exposure requires immediate irrigation with normal saline or a solution for at least 20-30 minutes, using 1-2 liters of fluid while holding eyelids open, followed by topical 1% drops if corneal damage is suspected; urgent consultation is essential for ongoing evaluation and potential subconjunctival calcium injection. For exposure, provide supplemental to support respiration and prevent , while nebulizing 2.5% may be attempted; benzodiazepines such as are used to control any resultant seizures. In cases of ingestion, administer 4-12 ounces of water or milk if the patient is conscious to dilute the acid, but avoid due to the high risk of esophageal or gastric perforation from the corrosive nature of HF. In hospital settings, protocols include serial monitoring of serum calcium and magnesium levels every 6 hours, along with ECG and assessment for at least 24 hours; is employed for severe renal failure or refractory , as outlined in updated CDC/ATSDR guidelines from the . All exposed individuals should receive even if , given the potential for delayed .

Handling and regulations

Hydrofluoric acid () requires specialized storage to prevent corrosion and accidental release due to its reactivity with many materials. It should be stored in containers lined with or (PTFE, commonly known as Teflon) to ensure compatibility and avoid degradation. Storage areas must be kept away from sources of or moisture to minimize the risk of leaks or reactions that could generate hazardous fumes, and facilities should maintain sufficient to keep airborne concentrations below 3 parts per million (ppm), aligning with occupational exposure limits. Secondary containment, such as spill trays or bunds, is essential in all storage locations to capture potential leaks. Personal protective equipment (PPE) is critical for safe handling of to protect against its severe corrosive and toxic effects. Gloves made from or Viton provide effective resistance to permeation for extended periods, while face shields and chemical-resistant aprons or suits are required to shield the face, eyes, and body. clothing must be avoided, as it readily absorbs and can prolong skin contact, exacerbating injury. All PPE should be inspected regularly and stored separately from work areas to prevent contamination. Transportation of HF is governed by international and national regulations to mitigate risks during shipping. Aqueous solutions of HF (up to 60% concentration) are classified under United Nations number UN 1790 as a Class 8 corrosive substance, requiring appropriate placards, labeling, and packaging in corrosion-resistant containers such as steel drums with protective linings. Anhydrous HF, handled as hydrogen fluoride gas, falls under UN 1056 and is subject to additional restrictions due to its toxicity (Class 2.3) and corrosivity. Shippers must comply with Department of Transportation (DOT) rules, including segregation from incompatible materials and emergency response information. In the event of a spill, immediate and neutralization are necessary to limit exposure and environmental impact. Small spills should be neutralized using (calcium oxide) or soda ash () to form less hazardous salts, followed by absorption with inert materials like or sand before disposal. Larger spills require evacuation and professional response teams. HF waste, including spill residues, is designated under EPA code D002 for its corrosivity (pH ≤2), and must be managed through permitted treatment, storage, and disposal facilities. HF is subject to stringent regulatory oversight in multiple jurisdictions to ensure safe management and minimize environmental release. , it is listed on the Toxic Substances Control Act (TSCA) inventory, with ongoing evaluations of its risks in industrial applications. Under the EU's REACH regulation, HF is classified as a due to its severe health hazards, imposing authorization and restriction requirements for uses like in refineries. Emissions are controlled under the US Clean Air Act as a hazardous air pollutant, mandating maximum achievable control technology for sources exceeding thresholds. As of 2025, environmental groups have petitioned the EPA under TSCA Section 21 to prohibit HF use in domestic oil refining, citing unreasonable risks, leading to litigation and potential phaseout rules. Workers handling HF, particularly in high-risk processes like units, must undergo comprehensive training as required by OSHA's standard (29 CFR 1910.119). This includes initial and every three years on HF-specific hazards, safe work practices, emergency procedures, and process equipment integrity. Facilities must develop written safety programs, conduct hazard analyses, and ensure mechanical integrity to prevent releases.

History

Discovery and early development

Hydrofluoric acid was first isolated in 1771 by Swedish chemist Carl Wilhelm Scheele through the distillation of fluorspar (calcium fluoride, CaF₂) with concentrated sulfuric acid, producing a corrosive liquid that etched glass vessels. Scheele noted its unique ability to dissolve silicates, distinguishing it from other mineral acids, though he did not fully characterize its composition. In the early , British chemist , building on André-Marie Ampère's hypothesis, proposed that fluoric acid consisted of hydrogen combined with a novel element he termed "," based on its etymological root from fluorspar. This naming reflected early observations of its -corroding properties, which complicated handling and led to experiments in lead or vessels as alternatives to . Advances in the mid-19th century included French chemist Edmond Frémy's 1856 preparation of hydrofluoric acid by distilling , yielding a purer form free of water that intensified its reactivity. This work paved the way for Henri Moissan's 1886 electrolytic isolation of elemental from a molten of and hydrofluoric acid, confirming the acid's structure as (HF) and earning Moissan the 1906 . Early applications were limited by the acid's extreme hazards, including severe burns and toxicity, but it found niche use in etching gemstones such as (a variety of beryl) to create decorative frosted effects, often on a small scale in artisanal settings. Pre-20th century production remained artisanal, primarily via Scheele's fluorspar-sulfuric acid method or occasional alternatives like cryolite decomposition, yielding modest quantities for laboratory and craft purposes. Key scientific milestones included Leopold Gmelin's 1819 Handbuch der theoretischen Chemie, which systematically documented hydrofluoric acid's properties and reactions, establishing it as a foundational reference in . These developments underscored the acid's dual role as a powerful and perilous substance, shaping cautious early research protocols.

Industrial adoption and notable incidents

Hydrofluoric acid (HF) saw significant industrial adoption in the early 20th century, particularly in the production of chlorofluorocarbons (CFCs) such as Freon, which was developed by Thomas Midgley Jr. in 1928 as a non-toxic refrigerant alternative. HF served as a key fluorinating agent in the synthesis of these compounds, enabling the rapid commercialization of refrigeration and air conditioning technologies. During World War II, HF played a critical role in the Manhattan Project for uranium enrichment, where it was used to convert uranium oxide to uranium tetrafluoride and subsequently to uranium hexafluoride gas for gaseous diffusion processes. Post-1950s, HF use expanded dramatically in petroleum refining through processes to produce high-octane . patented and commercialized the first HF in 1942, marking a shift from sulfuric acid-based methods due to HF's efficiency and recyclability, which spurred widespread adoption in U.S. refineries during the post-war economic boom. By the 1970s, surging demand in the further drove HF consumption, as it became essential for silicon wafers and cleaning semiconductor surfaces in the burgeoning manufacturing sector. Several notable incidents underscored HF's hazards and influenced industrial practices. On October 30, 1987, a construction accident at the refinery in , ruptured an HF storage tank, releasing a toxic vapor cloud that affected nearby residents and workers, resulting in over 700 medical treatments for respiratory irritation and burns, with thousands evacuated. In June 2019, explosions at the Philadelphia Energy Solutions refinery released approximately 5,000 pounds of HF, along with a 38,000-pound vessel fragment that landed off-site, prompting evacuations and highlighting vulnerabilities in aging infrastructure, though no fatalities occurred. More recently, between October 2021 and June 2024, multiple releases occurred at the Geismar facility in due to ruptures, affecting workers and prompting a CSB . The final , released on May 27, 2025, highlighted systemic failures. These events contributed to stricter measures, including the development of OSHA's standard in 1992, which was partly motivated by -related risks and the Clean Air Act Amendments of 1990. The amendments also prompted EPA assessments, such as the 1993 on HF accident risks, fueling ongoing discussions about phasing out HF in refining to mitigate potential catastrophic releases. Globally, HF production has grown substantially since the mid-20th century, from modest scales in the 1950s to over 3 million metric tons annually by the 2020s, largely driven by demand for fluoropolymers and electronics applications. Regulatory frameworks evolved accordingly, with HF registered under the EU's REACH regulation effective from 2007, requiring detailed data and risk management. In the U.S., the EPA has conducted ongoing risk evaluations, including a 2022 recommendation from the and Hazard Investigation Board to prioritize HF under TSCA for risk evaluation.

Cultural references

Hydrofluoric acid () has gained notoriety in popular media for its extreme corrosiveness, most prominently in the television series (2008–2013), where characters use it to dissolve human remains in a , leading to the acid eating through the and floor. This depiction, while dramatized, heightened public awareness of HF's ability to penetrate tissues and cause severe, delayed burns, though scientific analysis later clarified that concentrated HF cannot fully dissolve as rapidly or completely as shown. The scene's impact extended to educational discussions on , influencing viewer perceptions of acids in narratives. In literature, HF appears in forensic and thriller contexts, such as Jeff Lindsay's Dexter book series (2004–2015), where it is referenced in methods for evidence disposal, perpetuating myths about its efficacy in "melting" bodies without trace. These portrayals often exaggerate HF's speed and totality of dissolution for narrative tension, drawing from real chemical properties but amplifying them into plot devices that explore moral and ethical dilemmas in . Such references have contributed to broader cultural fascination with acids in , blending fact with to underscore the substance's dangers. Documentaries and films from the have further explored HF's risks through educational lenses, including the MythBusters episode in 2013 testing the Breaking Bad acid bath scenario, which demonstrated that HF fails to dissolve a carcass (as a proxy) or breach a bathtub within realistic timeframes, using enhanced acids like sulfuric for comparison. Educational videos since 2010, such as analyses by chemistry educators, have dissected these scenes to correct misconceptions while highlighting HF's unique toxicity, often tying briefly to its real hazards like bone dissolution and systemic that inspire such fiction. These productions have served as informal announcements, promoting safer chemical understanding amid entertainment. Cultural misconceptions portray HF as an "instant death acid" capable of immediate, total destruction, fueled by urban legends and media tropes that ignore its weak acidity and slower action compared to stronger corrosives like . These myths, often spread via online forums and anecdotal tales, overestimate HF's body-dissolving prowess, leading to unnecessary fear; CDC fact sheets counter this by emphasizing that while HF causes insidious, penetrating burns without initial pain, it requires specific medical intervention like rather than evoking instant lethality. Such clarifications help demystify HF, distinguishing its fluoride ion toxicity from sensationalized invincibility. In art and scientific illustration, HF has historical ties to 19th-century glass etching techniques, where it was used to create intricate designs by selectively dissolving layers of , as seen in works like the 1867 Greek Revival vase at the , showcasing its role in producing frosted, decorative surfaces. This application transformed HF from an industrial tool into an artistic medium, enabling precise patterns in luxury glassware during the . Modern biohazard art incorporates HF motifs symbolically, appearing in safety posters and digital illustrations that depict corrosive warnings alongside hazard symbols, reflecting contemporary concerns over chemical spills in creative expressions of environmental risk. Art safety guidelines highlight HF's presence in studio materials, urging precautions to prevent its dangers from infiltrating artistic practice. As of 2025, on have addressed in the wake of recent industrial incidents, such as the January episode of Chemical Processing's process safety series discussing recurring accidents involving corrosive releases and the need for enhanced emergency protocols. These discussions, including analyses of thermal breakdowns producing byproducts in refrigerants on the Industrial Strength , emphasize organizational learning from spills to mitigate threats, bridging cultural awareness with preventive strategies.