Hydrogen iodide

Hydrogen iodide (HI) is a diatomic molecule. Aqueous solutions of HI are known as hydroiodic acid or hydriodic acid, a strong acid. Hydrogen iodide and hydroiodic acid are, however, different in that the former is a gas under standard conditions; whereas, the other is an aqueous solution of said gas. They are interconvertible. HI is used in organic and inorganic synthesis as one of the primary sources of iodine and as a reducing agent.

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Properties of hydrogen iodide

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, the final solution having only four water molecules per molecule of HI.^[1]

Hydroiodic acid

Once again, although chemically related, hydroiodic acid is not pure HI but a mixture containing it. Commercial "concentrated" hydroiodic acid usually contains 48% - 57% HI by mass. The solution forms an azeotrope boiling at 127 °C with 57% HI, 43% water. Hydroiodidic acid is one of the strongest of all the common halide acids, despite the fact that the electronegativity of iodine is weaker than the rest of the other common halides. The high acidity is caused by the dispersal of the ionic charge over the anion. The iodide ion is much larger than the other common halides which results in the negative charge being dispersed over a large space. By contrast, a chloride ion is much smaller, meaning its negative charge is more concentrated, leading to a stronger interaction between the proton and the chloride ion. This weaker H⁺---I⁻ interaction in HI facilitates dissociation of the proton from the anion.

HI(g) + H2O(l)	⇌ H3O+(aq) + I–(aq)	(Ka ≈ 1010)
HBr(g) + H2O(l)	⇌ H3O+(aq) + Br–(aq)	(Ka ≈ 109)
HCl(g) + H2O(l)	⇌ H3O+(aq) + Cl–(aq)	(Ka ≈ 108)

Preparation

The industrial preparation of HI involves the reaction of I₂ with hydrazine, which also yields nitrogen gas.^[2]

2 I₂ + N₂H₄ → 4 HI + N₂

When performed in water, the HI must be distilled.

HI can also be distilled from a solution of NaI or other alkali iodide in concentrated phosphoric acid (note that sulfuric acid will not work for acidifying iodides as it will oxidize the iodide to elemental iodine).

Additionally HI can be prepared by simply combining H₂ and I₂. This method is usually employed to generate high purity samples.

H₂ + I₂ → 2 HI

For many years, this reaction was considered to involve a simple bimolecular reaction between molecules of H₂ and I₂. However, when a mixture of the gases is irradiated with the wavelength of light equal to the dissociation energy of I₂, about 578 nm, the rate increases significantly. This supports a mechanism whereby I₂ first dissociates into 2 iodine atoms, which each attach themselves to a side of an H₂ molecule and break the H -- H bond:^[3]

H₂ + I₂ + 578 nm radiation → H₂ + 2 I → I - - - H - - - H - - - I → 2 HI

In the laboratory, another method involves hydrolysis of PI₃, the iodine equivalent of PBr₃. In this method, I₂ reacts with phosphorus to create phosphorus triiodide, which then reacts with water to form HI and phosphorous acid.

3 I₂ + 2 P + 6 H₂O → 2 PI₃ + 6 H₂O → 6 HI + 2 H₃PO₃

Key reactions and applications

• HI will undergo oxidation if left open to air according to the following pathway:'

4 HI + O₂ → 2H₂O + 2 I₂

HI + I₂ → HI₃

HI₃ is dark brown in color, which makes aged solutions of HI often appear dark brown.

• Like HBr and HCl, HI add to alkenes⁴:

HI + H₂C=CH₂ → H₃CCH₂I

HI is subject to the same Markovnikov and anti-Markovnikov guidelines as HCl and HBr.

• HI reduces certain α-substituted ketones and alcohols replacing the α substituent with a hydrogen atom.^[4]

References

1. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
2. ^ Greenwood, N.N. and A. Earnshaw. The Chemistry of the Elements. 2nd ed. Oxford: Butterworth-Heineman. p 809-815. 1997.
3. ^ Holleman, A.F. Wiberg, E. Inorganic Chemistry. San Diego: Academic Press. p 371, 432-433. 2001.
4. ^ Breton, G. W., P. J. Kropp, P. J.; Harvey, R. G. “Hydrogen Iodide” in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. DOI: 10.1002/047084289.

See also: Nishikata, E., T.; Ishii, and T. Ohta. “Viscosities of Aqueous Hydrochloric Acid Solutions, and Densities and Viscosities of Aqueous Hydroiodic Acid Solutions”. J. Chem. Eng. Data. 26. 254-256. 1981.