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Hydrogen halide
Hydrogen halide
Hydrogen halide
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In chemistry, hydrogen halides (hydrohalic acids when in the aqueous phase) are diatomic, inorganic compounds that function as Arrhenius acids. The formula is HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, astatine, or tennessine.[1] All known hydrogen halides are gases at standard temperature and pressure.[2]

Compound Chemical formula Bond length
d(H−X) / pm
(gas phase)		 model Dipole
μ / D		 Aqueous phase (acid) Aqueous Phase pKa values
hydrogen fluoride
(fluorane)		 HF 1.86 hydrofluoric acid 3.1
hydrogen chloride
(chlorane)		 HCl 1.11 hydrochloric acid −3.9
hydrogen bromide
(bromane)		 HBr 0.788 hydrobromic acid −5.8
hydrogen iodide
(iodane)		 HI 0.382 hydroiodic acid −10.4 [3]
hydrogen astatide
astatine hydride
(astatane)		 HAt −0.06 hydroastatic acid ?
hydrogen tennesside
tennessine hydride
(tennessane)		 HTs −0.24 ? hydrotennessic acid ?[4]

Comparison to hydrohalic acids

[edit]

The hydrogen halides are diatomic molecules with no tendency to ionize in the gas phase (although liquified hydrogen fluoride is a polar solvent somewhat similar to water). Thus, chemists distinguish hydrogen chloride from hydrochloric acid. The former is a gas at room temperature that reacts with water to give the acid. Once the acid has formed, the diatomic molecule can be regenerated only with difficulty, but not by normal distillation. Commonly the names of the acid and the molecules are not clearly distinguished such that in lab jargon, "HCl" often means hydrochloric acid, not the gaseous hydrogen chloride.

Occurrence

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Hydrogen chloride, in the form of hydrochloric acid, is a major component of gastric acid.

Hydrogen fluoride, chloride and bromide are also volcanic gases.

Synthesis

[edit]

The direct reaction of hydrogen with fluorine and chlorine gives hydrogen fluoride and hydrogen chloride, respectively. Industrially these gases are, however, produced by treatment of halide salts with sulfuric acid. Hydrogen bromide arises when hydrogen and bromine are combined at high temperatures in the presence of a platinum catalyst. The least stable hydrogen halide, HI, is produced less directly, by the reaction of iodine with hydrogen sulfide or with hydrazine.[1]: 809–815 

Physical properties

[edit]
Comparison of the boiling points of hydrogen halides and hydrogen chalcogenides; here it can be seen that hydrogen fluoride breaks trends alongside water.

The hydrogen halides are colourless gases at standard conditions for temperature and pressure (STP) except for hydrogen fluoride, which boils at 19 °C. Alone of the hydrogen halides, hydrogen fluoride exhibits hydrogen bonding between molecules, and therefore has the highest melting and boiling points of the HX series. From HCl to HI the boiling point rises. This trend is attributed to the increasing strength of intermolecular van der Waals forces, which correlates with numbers of electrons in the molecules. Concentrated hydrohalic acid solutions produce visible white fumes. This mist arises from the formation of tiny droplets of their concentrated aqueous solutions of the hydrohalic acid.

Reactions

[edit]

Upon dissolution in water, which is highly exothermic, the hydrogen halides give the corresponding acids. These acids are very strong, reflecting their tendency to ionize in aqueous solution yielding hydronium ions (H3O+). With the exception of hydrofluoric acid, the hydrogen halides are strong acids, with acid strength increasing down the group. Hydrofluoric acid is complicated because its strength depends on the concentration owing to the effects of homoconjugation. As solutions in non-aqueous solvents, such as acetonitrile, the hydrogen halides are only modestly acidic however.

Similarly, the hydrogen halides react with ammonia (and other bases), forming ammonium halides:

HX + NH3 → NH4X

In organic chemistry, the hydrohalogenation reaction is used to prepare halocarbons. For example, chloroethane is produced by hydrochlorination of ethylene:[5]

C2H4 + HCl → CH3CH2Cl

See also

[edit]

References

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Hydrogen halide

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Hydrogen halides are binary inorganic compounds consisting of hydrogen and a halogen atom (fluorine, chlorine, bromine, or iodine) in the −1 oxidation state, with the general formula HX, where X represents the halogen; these include (HF), (HCl), (HBr), and (HI). They exist as colorless, diatomic gases at and standard pressure, exhibiting polar covalent bonding due to the electronegativity difference between hydrogen and the halogen, with bond polarity increasing from HI to HF. Highly soluble in water, they form hydrohalic acids that fully dissociate except for HF, which is a weak acid, and all display sharp, irritating odors and steamy fumes in moist air. The physical properties of hydrogen halides vary systematically down the group. Melting and boiling points generally increase from HCl to HI due to increasing , but HF anomalously has the highest values among them ( 190 K, 293 K) owing to extensive intermolecular hydrogen bonding, forming zig-zag polymeric chains in the solid and liquid states, unlike the other monomeric . Bond dissociation energies decrease from HF (562 kJ/mol) to HI (299 kJ/mol), reflecting weaker H–X bonds as halogen size increases, which correlates with thermal stability trends where HI decomposes most readily. In , acidity strengthens dramatically from HF (K_a ≈ 5.6 × 10^{-4}) to HI (K_a ≈ 5.0 × 10^{10}), driven by progressively weaker H–X bonds and reduced effects on the larger ions, making HCl, HBr, and HI strong acids that fully ionize. Hydrogen halides are prepared industrially by direct combination of gas with the elemental , often requiring controlled conditions to manage exothermic reactions (e.g., H_2 + Cl_2 → 2HCl), or via acid treatment of salts like NaCl with H_2SO_4 for HCl. They exhibit versatile reactivity, including to alkenes following , where the hydrogen adds to the carbon with more hydrogens, and serve as sources of ions in synthesis. Notable chemical behaviors include HF's unique ability to etch due to its reaction with silica (SiO_2 + 4HF → SiF_4 + 2H_2O), while the others are less reactive toward silicates but corrosive to metals. In applications, hydrogen halides are essential in chemical manufacturing: HCl is used in the production of for PVC plastics and in metal cleaning, HBr in the synthesis of organobromine compounds and pharmaceuticals, HI as a , and HF in fluorochemicals like refrigerants and aluminum production. Their aqueous forms, known as hydrohalic acids, are vital reagents, though handling requires caution due to and corrosiveness, with HF posing particular risks from deep tissue penetration via binding to calcium.

Overview

Definition and nomenclature

Hydrogen halides are diatomic molecules composed of one hydrogen atom bonded to one halogen atom, represented by the general formula HX, where X denotes a halogen from group 17 of the periodic table. The halogens include fluorine (F), chlorine (Cl), bromine (Br), and iodine (I); astatine (At) is excluded due to its extreme rarity and radioactivity. These group 17 elements are highly electronegative nonmetals that readily form covalent bonds with hydrogen, resulting in polar diatomic species. As binary compounds consisting solely of hydrogen and a single halogen, hydrogen halides are classified both as binary acids—due to their capacity to release a proton—and as covalent hydrides of the halogens. The nomenclature follows systematic IUPAC conventions, naming them as hydrogen followed by the halogen name: hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), and hydrogen iodide (HI). These abbreviations are widely used in chemical literature for brevity. The discovery of hydrogen halides dates back to the 17th and 18th centuries, marking early advances in and isolation. Hydrogen chloride was first prepared in 1648 by German chemist , who heated with , producing the gas as a byproduct distinct from metal chloride salts. In 1771, Swedish chemist isolated by distilling fluorspar (calcium fluoride) with , noting its unique etching properties on glass and recognizing it as a novel acidic substance separate from salts. These early isolations highlighted the gaseous nature of hydrogen halides at standard conditions and laid the foundation for understanding their chemistry as pure compounds.

Comparison to hydrohalic acids

Hydrohalic acids are the aqueous solutions of hydrogen halides (HX), where X represents a halogen atom (F, Cl, Br, or I). For example, hydrochloric acid is the aqueous solution of hydrogen chloride, denoted as HCl(aq). The anhydrous hydrogen halides exist primarily as molecular, diatomic gases at standard conditions, with hydrogen fluoride exhibiting a notably higher boiling point (19.5 °C) due to extensive hydrogen bonding, while the others (HCl, HBr, HI) are gases with lower boiling points. In contrast, hydrohalic acids are ionized in water, forming hydronium ions and halide ions: HX(g)+H2O(l)→H3O+(aq)+X−(aq)\text{HX(g)} + \text{H}_2\text{O(l)} \rightarrow \text{H}_3\text{O}^+(\text{aq}) + \text{X}^-(\text{aq}) The completeness of this dissociation varies by halide; HF is a weak acid with partial ionization, whereas HCl, HBr, and HI are strong acids that fully dissociate. Behaviorally, hydrogen halides are employed in non-aqueous reactions, such as HCl in pharmaceutical synthesis where must be avoided, and they are typically stored as compressed gases or liquefied under pressure. Hydrohalic acids, however, are used in aqueous chemistry applications like adjustment or metal and are stored as liquid solutions, often at concentrations up to 38% for HCl to form azeotropes. A unique aspect of HF arises from the ion's small size and high , which promotes its reactivity in to form complex s such as the bifluoride ion [HF₂]⁻, enhancing its corrosive properties beyond simple proton donation.

Sources

Natural occurrence

Hydrogen halides occur naturally in the atmosphere primarily through volcanic emissions, where HCl and HF are released as components of volcanic gases. Volcanoes emit HCl at concentrations up to 1-5% by volume in some fumarolic gases, contributing an estimated global flux of 0.4-11 Tg per year, while HF emissions are typically lower, around 0.06-6 Tg annually. HBr and HI are less prominent in volcanic outputs but arise from sea spray aerosols and hydrothermal vents; debromination releases HBr into the , serving as a key source of reactive . In oceanic environments, hydrogen halides exist mainly as dissociated in salts such as NaCl, KBr, and NaI. ions dominate at approximately 19 g/kg in open , comprising over half of the total of 35 g/kg, while and are present at much lower levels (around 65 mg/kg and 60 µg/kg, respectively). Biological sources include the accumulation of in marine algae, particularly seaweeds, which can concentrate iodine up to thousands of times levels, potentially releasing HI or related volatiles during decay or stress. Trace HF arises in through uptake from or atmospheric deposition, with like and grapevines accumulating up to several hundred mg/kg in leaves under natural fluoride exposure. Astatine halides do not occur naturally in significant amounts due to the element's extreme and short half-lives, with trace produced only as a in ores. These natural emissions contribute to environmental processes, including formation from volcanic HCl, which dissolves in rainwater to lower near eruption sites, and limited primarily from and species rather than HF, which remains largely inert in the .

Synthesis

Hydrogen halides can be synthesized through both direct and indirect laboratory methods, as well as large-scale industrial processes. Historically, the preparation of (HCl) dates back to the 17th century, when German chemist heated a mixture of (common salt) and (known as oil of vitriol) to produce HCl gas, marking one of the earliest practical methods for isolating the compound. This approach laid the foundation for subsequent developments in halide synthesis. In laboratory settings, direct synthesis involves the exothermic reaction of hydrogen gas with the elemental halogen:
\ceH2(g)+X2(g)−>2HX(g)\ce{H2(g) + X2(g) -> 2HX(g)}
where X represents F, Cl, Br, or I. For HCl, the reaction is initiated by an or exposure to light, as the mixture is explosive under certain conditions. (HBr) and (HI) require platinum catalysts and elevated temperatures for efficient combination, while (HF) demands specialized equipment due to its extreme reactivity with glass and most metals. Indirect methods typically start from metal halides and a non-volatile acid to liberate the hydrogen halide gas. A common laboratory procedure for HCl involves heating with concentrated at around 500°C:
\ceNaCl(s)+H2SO4(l)−>NaHSO4(s)+HCl(g)\ce{NaCl(s) + H2SO4(l) -> NaHSO4(s) + HCl(g)}
Similar reactions apply to HBr and HI using or iodide, though HI synthesis prefers (H3PO4) over to prevent oxidation of iodide to iodine. These methods allow controlled generation of pure HX gases for experimental use. On an industrial scale, HCl is predominantly produced as a of organic chlorination processes, such as the production of or , generating approximately 10 million tons annually worldwide. The gas is captured and absorbed in to form solutions. In contrast, HF is manufactured by reacting fluorspar (CaF2) with concentrated in a or packed-bed reactor at elevated temperatures:
\ceCaF2(s)+H2SO4(l)−>CaSO4(s)+2HF(g)\ce{CaF2(s) + H2SO4(l) -> CaSO4(s) + 2HF(g)}
The HF vapor is then condensed and purified. HBr and HI see limited industrial production, often via direct synthesis or of phosphorus halides for niche applications. Purification of the resulting hydrogen halides is essential for high-purity applications. HCl and HBr are typically purified by , where impurities like moisture or residual acids are removed under controlled conditions to yield gases. For HF, addresses challenges from its with water (at 38% HF), often involving multiple stages or addition of to break the and produce HF. Safety considerations are paramount during synthesis due to inherent hazards. Direct combination of H2 and X2 (especially Cl2) poses explosion risks from ignition by light, heat, or sparks, necessitating inert atmospheres and remote initiation. HF synthesis and handling require corrosion-resistant or vessels, as it aggressively attacks , , and many alloys, leading to potential leaks or equipment failure.

Properties

Physical properties

The hydrogen halides (HF, HCl, HBr, and HI) are all colorless, nonflammable gases at room temperature and standard pressure, though HF has a boiling point of 19.5°C, rendering it a liquid under conditions slightly below typical room temperature but still a gas above this threshold near standard temperature and pressure (STP). Their melting and boiling points exhibit a general increasing trend from HCl to HI due to rising molecular weights, which enhance London dispersion forces; however, HF displays an anomaly with a significantly higher than expected, attributed to strong intermolecular hydrogen bonding between the highly electronegative fluorine and hydrogen atoms. points follow a similar , with HI highest and HCl lowest, influenced by intermolecular forces and molecular size. The following table summarizes these values:
CompoundMelting Point (°C)Boiling Point (°C)
HF-83.619.5
HCl-114.2-85.1
HBr-88.5-66.8
HI-50.8-35.6
The H–X bond lengths increase down the group from HF to HI due to the enlarging of the halogen atom, leading to weaker orbital overlap and progressively lower bond dissociation energies: HF (570 kJ/mol), HCl (432 kJ/mol), HBr (366 kJ/mol), and HI (298 kJ/mol). This trend arises from the increasing size mismatch between the small hydrogen 1s orbital and the larger p orbitals of the heavier , reducing bond strength. Gas densities at STP increase with molecular weight: HF (0.89 g/L), HCl (1.63 g/L), HBr (3.61 g/L), and HI (5.71 g/L). All hydrogen halides are highly soluble in , forming the corresponding hydrohalic acids; HF is miscible with , achieving concentrations up to approximately 68% by weight at 20°C due to extensive hydrogen bonding. As polar diatomic molecules, the hydrogen halides exhibit moments, making their H–X stretching vibrations active. Vibrational frequencies decrease down the group owing to increasing and : HF (≈4140 cm⁻¹), HCl (≈2990 cm⁻¹), HBr (≈2650 cm⁻¹), and HI (≈2310 cm⁻¹). Standard enthalpies of formation (ΔH_f°) become less negative down the group, reflecting weakening H–X bonds: HF (–271 kJ/mol), HCl (–92 kJ/mol), HBr (–36 kJ/mol), and HI (+26 kJ/mol).

Reactions

Hydrogen halides exhibit varying acidity depending on the phase. In the gas phase, acidity increases down the group from HF to HI due to decreasing H–X bond strength, which facilitates , with HF being the weakest acid and HI the strongest. In , the trend is similar, but all except HF are strong acids that fully dissociate; HF remains weak with a pKa of 3.17, while HCl (pKa ≈ -6.3), HBr (pKa ≈ -8.7), and HI (pKa ≈ -9.3) are leveled by water's solvent properties. Hydrogen halides react with active metals to form metal halides and liberate hydrogen gas, demonstrating their acidic character. For example, the reaction with sodium follows 2HCl + 2Na → 2NaCl + H₂, where the metal reduces the halide to H₂ while oxidizing to the halide salt. Similar reactions occur with other electropositive metals like zinc or magnesium, such as Zn + 2HBr → ZnBr₂ + H₂, proceeding via proton reduction. In water, HCl, HBr, and HI undergo complete to form ions and ions, behaving as strong acids: HX + H₂O → H₃O⁺ + X⁻ (X = Cl, Br, I). HF, however, hydrolyzes only partially due to strong hydrogen bonding between F⁻ and H₂O, resulting in the equilibrium HF + H₂O ⇌ H₃O⁺ + F⁻ with limited dissociation. Oxidation sensitivity increases down the group, with HI and HBr readily oxidized by oxygen to form the free , as in 4HI + O₂ → 2I₂ + 2H₂O, reflecting their reducing power. HCl and HF are more stable toward oxidation due to stronger bonds and lower reducing tendency. This trend correlates with decreasing standard redox potentials for the X₂/X⁻ couples: E°(F₂/F⁻) = +2.87 V, E°(Cl₂/Cl⁻) = +1.36 V, E°(Br₂/Br⁻) = +1.07 V, E°(I₂/I⁻) = +0.54 V, making heavier easier to oxidize. Hydrogen halides do not typically form interhalogens directly, but their interactions with free vary; for instance, HCl shows no reaction with Cl₂ under standard conditions. In , hydrogen halides add to alkenes via , following ; for example, HCl + CH₂=CH₂ → CH₃CH₂Cl, where H adds to the less substituted carbon. HCl also participates in Friedel–Crafts alkylation when generated with AlCl₃, facilitating . HF exhibits unique reactivity, such as etching by reacting with SiO₂: SiO₂ + 4HF → SiF₄ + 2H₂O, dissolving the silica network. Additionally, HF forms stable complexes with amines, like pyridinium poly(), which serve as safer fluorinating agents through hydrogen-bonded networks.

References

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