Hydrogen iodide

Strongest hydrohalic acid and versatile reducing agent.

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Hydrogen iodide, with the formula HI, is a diatomic molecule belonging to the hydrogen halide family. When dissolved in water, it forms hydroiodic acid (also called hydriodic acid), a strong acid. It is important to distinguish between the two: under standard conditions, hydrogen iodide is a gas, while hydroiodic acid is the gas dissolved in water. The two forms can be converted back and forth. HI serves as a key source of iodine and a reducing agent in both organic and inorganic synthesis.

As a colorless gas, hydrogen iodide reacts with oxygen to produce water and iodine. In moist air, it forms a mist of hydroiodic acid. It is extremely soluble in water—one liter of water can absorb 425 liters of HI gas, and the most concentrated solution contains only four water molecules for every molecule of HI.

Hydroiodic acid, the aqueous solution, is commercially available at 48–57% HI by mass. This solution forms an azeotrope that boils at 127 °C, consisting of 57% HI and 43% water. Its high acidity comes from the large size of the iodide ion, which spreads the negative charge over a wide volume. This weakens the attraction between the proton and the anion, making it easier for the proton to dissociate. As a result, HI is the strongest acid among the hydrohalides, with an acid dissociation constant (Ka) around 10¹⁰, compared to roughly 10⁹ for HBr and 10⁶ for HCl.

Industrially, HI is made by reacting iodine with hydrazine, which also produces nitrogen gas. When done in water, the HI can be purified by distillation. Anhydrous HI can be prepared by reacting iodine with tetrahydronaphthalene. Another method distills HI from a solution of sodium iodide or another alkali iodide treated with phosphorus pentoxide, which acts as a dehydrating agent and yields phosphoric acid. Concentrated sulfuric acid cannot be used to acidify iodides because it oxidizes iodide to elemental iodine. Historically, HI was produced by oxidizing hydrogen sulfide with aqueous iodine. HI can also be made by simply combining hydrogen and iodine gases, a reaction often catalyzed by platinum to yield high-purity samples. For many years, this was thought to be a simple bimolecular reaction, but irradiating the gas mixture with light at about 578 nm—the wavelength that breaks iodine molecules apart—greatly speeds up the process. This supports a mechanism where iodine atoms first form, then attach to hydrogen molecules and break the H–H bond. In the lab, another route involves hydrolyzing phosphorus triiodide (PI₃), made by reacting iodine with phosphorus; the PI₃ then reacts with water to give HI and phosphorous acid.

Solutions of hydrogen iodide are easily oxidized by air, turning brown due to the formation of iodine. Like HCl and HBr, HI adds to alkenes following Markovnikov and anti-Markovnikov rules. In organic chemistry, HI converts primary alcohols into alkyl iodides via an SN2 substitution, where the iodide ion replaces the hydroxyl group (which leaves as water). HI is sometimes preferred over other hydrogen halides for this reaction. HI (or HBr) can also cleave ethers, commonly breaking aryl-alkyl ethers into phenols and alkyl iodides. For example, diethyl ether splits into two equivalents of ethyl iodide. The reaction is regioselective, with iodide attacking the less hindered carbon of the ether.

Early in organic chemistry, HI was widely used as a reducing agent. In the 19th century, chemists tried to reduce benzene to cyclohexane with HI at high temperatures but instead obtained methylcyclopentane, a rearranged product. Kiliani reported that hydroiodic acid reduces sugars and other polyols by cleaving multiple hydroxyl groups, though yields and reproducibility were often poor. For benzyl alcohols and alcohols with alpha-carbonyl groups, HI reduction can give useful yields of the corresponding hydrocarbons. This process can be made catalytic in HI by adding red phosphorus, which reduces any iodine that forms.

Commercially, iodine is obtained from iodide-rich brines. The process starts by converting iodide into hydroiodic acid, which is then oxidized to iodine. The iodine is recovered by evaporation or adsorption.

Quick Facts

State at standard conditions: colorless gas
Solubility in water: 425 liters of HI gas per liter of water
Boiling point of azeotrope: 127 °C (57% HI, 43% water)
Acid dissociation constant (ka): approximately 10^10
Commercial concentration: 48–57% HI by mass

Facts from the source article.

Lore & Background

HI is a colorless gas that reacts with oxygen to give water and iodine. With moist air, HI gives a mist (or fumes) of hydroiodic acid. It is exceptionally soluble in water, giving hydroiodic acid. One liter of water will dissolve 425 liters of HI gas. The solution forms an azeotrope boiling at 127 °C with 57% HI, 43% water. The high acidity is caused by the dispersal of the ionic charge over the anion. The iodide ion radius is much larger than the other common halides, which results in the negative charge being dispersed over a large volume. This weaker H+···I− interaction in HI facilitates dissociation of the proton from the anion and is the reason HI is the strongest acid of the hydrohalides. Ka ≈ 10^10.

Chemical Identity and Acidic Dominance

Hydrogen iodide is a colorless diatomic gas at standard conditions, but its encounter with water produces one of the most formidable acids in the hydrohalide family. The gaseous molecule and its aqueous solution—hydroiodic acid—are distinct physical states of the same chemical entity, yet they behave very differently. HI dissolves in water with remarkable ease: one liter of water can absorb roughly 425 liters of the gas, yielding a solution so concentrated that only about four water molecules accompany each HI unit. The reason hydroiodic acid outperforms its lighter siblings as an acid is geometric. The iodide ion is substantially larger than chloride or bromide, so its negative charge is smeared across a much greater volume. This dilution of charge weakens the electrostatic attraction between the proton and the anion, making proton release far easier. The acid dissociation constant reflects this, sitting near 10 to the tenth power—two orders above hydrobromic acid and four above hydrochloric. In commerce, the acid appears at 48 to 57 percent by mass and forms a well-defined azeotrope boiling at 127 degrees Celsius.

Synthetic Routes and the Photodissociation Puzzle

Producing hydrogen iodide in the laboratory or on an industrial scale has never been a single-path affair. The most common industrial route couples iodine with hydrazine, a reaction that simultaneously liberates harmless nitrogen gas and, when run in water, allows the HI to be purified by simple distillation. For anhydrous material, chemists react iodine with tetrahydronaphthalene, stripping four hydrogens and releasing the naphthalene ring as a byproduct. An older, historically important method oxidizes hydrogen sulfide with aqueous iodine, depositing elemental sulfur while generating HI in solution. Perhaps the most elegant route is the direct combination of hydrogen and iodine gases, often catalyzed by platinum to yield high-purity product. For decades this was assumed to be a straightforward bimolecular collision. Then researchers discovered that irradiating the mixture at 578 nanometers—the dissociation energy of the iodine bond—dramatically accelerates the reaction. The light first splits I2 into two free iodine atoms, which then sandwich the hydrogen molecule and cleave its H–H bond, revealing a stepwise mechanism far richer than a simple one-step collision.

Reactivity in Organic Synthesis

In the organic chemist's toolkit, hydrogen iodide occupies a versatile niche. It adds across carbon–carbon double bonds following the same Markovnikov and anti-Markovnikov rules that govern its lighter halide cousins. More practically, HI converts primary alcohols into alkyl iodides through an SN2 displacement in which the iodide ion replaces the hydroxyl group as water departs. The reagent also cleaves ethers, a reaction that is regioselective because the iodide preferentially attacks the less sterically crowded carbon; aryl–alkyl ethers, for instance, split into phenols and alkyl iodides. HI's role as a reducing agent dates to the nineteenth century, when chemists attempted to hydrogenate benzene into cyclohexane at high temperature. The effort backfired, yielding the rearranged product methylcyclopentane instead. Kiliani later showed that hydroiodic acid can strip multiple hydroxyl groups from sugars and polyols, though yields and reproducibility were often poor. More reliably, HI reduces benzyl alcohols and alpha-carbonyl alcohols to their parent hydrocarbons, and the process can be made catalytic by adding red phosphorus to re-reduce the iodine that forms as a byproduct.

Industrial Iodine Production and the Chemistry of Aging

Beyond the laboratory, hydrogen iodide plays a pivotal intermediary role in the commercial extraction of iodine. The process begins with iodide-rich brines, where the iodide is first converted into hydroiodic acid. That acid is then oxidized back to elemental iodine, which is recovered by evaporation or adsorption. The HI thus serves as a chemical shuttle, carrying iodine from a dissolved ionic form into a state where it can be cleanly separated. In the bottle, however, hydroiodic acid is not static. Exposure to atmospheric oxygen steadily oxidizes the HI, producing water and elemental iodine. The iodine in turn complexes with excess HI to form triiodide, HI3, a brown species that tints aged solutions dark. This slow self-oxidation is why freshly prepared solutions appear clear while stored ones gradually darken. The phenomenon also underscores a practical limitation: hydroiodic acid is a reagent that must be handled with awareness of its own reactivity toward the very air in which it sits.