# 18.5 Occurrence, preparation, and compounds of hydrogen  (Page 5/9)

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${\text{2H}}_{2}\text{S}\left(g\right)+{\text{O}}_{2}\left(g\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{2S}\left(s\right)+{\text{2H}}_{2}\text{O}\left(l\right)$

This oxidation process leads to the removal of the hydrogen sulfide found in many sources of natural gas. The deposits of sulfur in volcanic regions may be the result of the oxidation of H 2 S present in volcanic gases.

Hydrogen sulfide is a weak diprotic acid that dissolves in water to form hydrosulfuric acid. The acid ionizes in two stages, yielding hydrogen sulfide ions, HS , in the first stage and sulfide ions, S 2− , in the second. Since hydrogen sulfide is a weak acid, aqueous solutions of soluble sulfides and hydrogen sulfides are basic:

${\text{S}}^{2-}\left(aq\right)+{\text{H}}_{2}\text{O}\left(l\right)\phantom{\rule{0.2em}{0ex}}⇌\phantom{\rule{0.2em}{0ex}}{\text{HS}}^{\text{−}}\left(aq\right)+{\text{OH}}^{\text{−}}\left(aq\right)$
${\text{HS}}^{\text{−}}\left(aq\right)+{\text{H}}_{2}\text{O}\left(l\right)\phantom{\rule{0.2em}{0ex}}⇌\phantom{\rule{0.2em}{0ex}}{\text{H}}_{2}\text{S}\left(g\right)+{\text{OH}}^{\text{−}}\left(aq\right)$

## Halogen hydrogen compounds

Binary compounds containing only hydrogen and a halogen are hydrogen halides . At room temperature, the pure hydrogen halides HF, HCl, HBr, and HI are gases.

In general, it is possible to prepare the halides by the general techniques used to prepare other acids. Fluorine, chlorine, and bromine react directly with hydrogen to form the respective hydrogen halide. This is a commercially important reaction for preparing hydrogen chloride and hydrogen bromide.

The acid-base reaction between a nonvolatile strong acid and a metal halide will yield a hydrogen halide. The escape of the gaseous hydrogen halide drives the reaction to completion. For example, the usual method of preparing hydrogen fluoride is by heating a mixture of calcium fluoride, CaF 2 , and concentrated sulfuric acid:

${\text{CaF}}_{2}\left(s\right)+{\text{H}}_{2}{\text{SO}}_{4}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{CaSO}}_{4}\left(s\right)+\text{2HF}\left(g\right)$

Gaseous hydrogen fluoride is also a by-product in the preparation of phosphate fertilizers by the reaction of fluoroapatite, Ca 5 (PO 4 ) 3 F, with sulfuric acid. The reaction of concentrated sulfuric acid with a chloride salt produces hydrogen chloride both commercially and in the laboratory.

In most cases, sodium chloride is the chloride of choice because it is the least expensive chloride. Hydrogen bromide and hydrogen iodide cannot be prepared using sulfuric acid because this acid is an oxidizing agent capable of oxidizing both bromide and iodide. However, it is possible to prepare both hydrogen bromide and hydrogen iodide using an acid such as phosphoric acid because it is a weaker oxidizing agent. For example:

${\text{H}}_{3}{\text{PO}}_{4}\left(l\right)+{\text{Br}}^{\text{−}}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{HBr}\left(g\right)+{\text{H}}_{2}{\text{PO}}_{4}{}^{\text{−}}\left(aq\right)$

All of the hydrogen halides are very soluble in water, forming hydrohalic acids. With the exception of hydrogen fluoride, which has a strong hydrogen-fluoride bond, they are strong acids. Reactions of hydrohalic acids with metals, metal hydroxides, oxides, or carbonates produce salts of the halides. Most chloride salts are soluble in water. AgCl, PbCl 2 , and Hg 2 Cl 2 are the commonly encountered exceptions.

The halide ions give the substances the properties associated with X ( aq ). The heavier halide ions (Cl , Br , and I ) can act as reducing agents, and the lighter halogens or other oxidizing agents will oxidize them:

${\text{Cl}}_{2}\left(aq\right)+{\text{2e}}^{-}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{2Cl}}^{\text{−}}\left(aq\right)\phantom{\rule{5em}{0ex}}E\text{°}=1.36\phantom{\rule{0.2em}{0ex}}\text{V}$
${\text{Br}}_{2}\left(aq\right)+{\text{2e}}^{-}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{2Br}}^{\text{−}}\left(aq\right)\phantom{\rule{5em}{0ex}}E\text{°}=1.09\phantom{\rule{0.2em}{0ex}}\text{V}$
${\text{I}}_{2}\left(aq\right)+{\text{2e}}^{-}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{2I}}^{\text{−}}\left(aq\right)\phantom{\rule{5em}{0ex}}E\text{°}=0.54\phantom{\rule{0.2em}{0ex}}\text{V}$

For example, bromine oxidizes iodine:

${\text{Br}}_{2}\left(aq\right)+\text{2HI}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{2HBr}\left(aq\right)+{\text{I}}_{2}\left(aq\right)\phantom{\rule{5em}{0ex}}E\text{°}=0.55\phantom{\rule{0.2em}{0ex}}\text{V}$

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