Hydrogen halide

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

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. Hydrogen chloride is a gas at room temperature that reacts with water to give hydrochloric acid; once the acid has formed, the hydrogen chloride can be regenerated, but only with difficulty and not by normal distillation. Often, the names of the acid and the molecules are not clearly distinguished, and in lab jargon, "HCl" often means hydrochloric acid, not the gaseous hydrogen chloride.

Occurrence and production

Hydrogen fluoride, chloride, and bromide are volcanic gases. ^{[5] [6]}

The hydrogen halides can be produced by many routes industrially and in the laboratory. Focusing on the most abundant compound, hydrogen chloride is mainly produced as a side product in production of chlorocarbons. ^[7] Hydrogen fluoride is a byproduct of the production of phosphoric acid. Fluorine, chlorine, and bromine react with hydrogen gas to give HF, HCl, and HBr. These gases can also be produced by treatment of halide salts with sulfuric acid. 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

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

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. However, as solutions in non-aqueous solvents, such as acetonitrile, the hydrogen halides are only moderately acidic.

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: ^[8]

C₂H₄ + HCl → CH₃CH₂Cl

See also

• Chemistry portal

• Pseudohalogen
• Hypohalous acid
• Group 13 hydrides
• Group 14 hydrides
• Group 15 hydrides
• Group 16 hydrides

References

1. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. doi:10.1016/C2009-0-30414-6. ISBN   978-0-08-037941-8.
2. The Acidity of the Hydrogen Halides. (2020, August 21). Retrieved May 5, 2021, from https://chem.libretexts.org/@go/page/3699
3. Schmid, Roland; Miah, Arzu M. (2001). "The Strength of the Hydrohalic Acids". Journal of Chemical Education. 78 (1). American Chemical Society (ACS): 116. Bibcode:2001JChEd..78..116S. doi: 10.1021/ed078p116 . ISSN   0021-9584.
4. de Farias, Robson Fernandes (January 2017). "Estimation of some physical properties for tennessine and tennessine hydride (TsH)". Chemical Physics Letters. 667: 1–3. Bibcode:2017CPL...667....1D. doi:10.1016/j.cplett.2016.11.023.
5. "Volcanic gases can be harmful to health, vegetation and infrastructure". U.S. Geological Survey . Retrieved 2026-02-02.
6. Edwards, Brock A.; Kushner, D. Skye; Outridge, Peter M.; Wang, Feiyue (2021-02-25). "Fifty years of volcanic mercury emission research: Knowledge gaps and future directions". Science of the Total Environment. 757 143800. doi:10.1016/j.scitotenv.2020.143800. ISSN   0048-9697. PMID   33280881.
7. Austin, Severin; Glowacki, Arndt (2000). "Hydrochloric Acid". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a13_283. ISBN   978-3-527-30385-4.
8. M. Rossberg et al. "Chlorinated Hydrocarbons" in Ullmann’s Encyclopedia of Industrial Chemistry, 2006, Wiley-VCH, Weinheim. doi : 10.1002/14356007.a06_233.pub2