Hydrogen Compounds
Other than the noble gases, each of the nonmetals forms compounds with hydrogen. For brevity, we will discuss only a few hydrogen compounds of the nonmetals here.
Nitrogen Hydrogen Compounds
Ammonia, NH3, forms naturally when any nitrogen-containing organic material decomposes in the absence of air. The laboratory preparation of ammonia is by the reaction of an ammonium salt with a strong base such as sodium hydroxide. The acid-base reaction with the weakly acidic ammonium ion gives ammonia, illustrated in the figure below. Ammonia also forms when ionic nitrides react with water. The nitride ion is a much stronger base than the hydroxide ion:
\({\text{Mg}}_{3}{\text{N}}_{2}(s)\;+\;{\text{6H}}_{2}\text{O}(l)\;⟶\;\text{3Mg}{(\text{OH})}_{2}(s)+{\text{2NH}}_{3}(g)\)
The commercial production of ammonia is by the direct combination of the elements in the Haber process:
\({\text{N}}_{2}(g)+{\text{3H}}_{2}(g)\stackrel{\;\text{catalyst}\;}{\text{⇌}}{\text{2NH}}_{3}(g)\phantom{\rule{5em}{0ex}}\text{Δ}H\text{°}=\text{−92 kJ}\)
The structure of ammonia is shown with a central nitrogen atom and three hydrogen atoms.
Ammonia is a colorless gas with a sharp, pungent odor. Smelling salts utilize this powerful odor. Gaseous ammonia readily liquefies to give a colorless liquid that boils at −33 °C. Due to intermolecular hydrogen bonding, the enthalpy of vaporization of liquid ammonia is higher than that of any other liquid except water, so ammonia is useful as a refrigerant. Ammonia is quite soluble in water (658 L at STP dissolves in 1 L H2O).
The chemical properties of ammonia are as follows:
- Ammonia acts as a Brønsted base, as discussed in the tutorial on acid-base chemistry. The ammonium ion is similar in size to the potassium ion; compounds of the two ions exhibit many similarities in their structures and solubilities.
- Ammonia can display acidic behavior, although it is a much weaker acid than water. Like other acids, ammonia reacts with metals, although it is so weak that high temperatures are necessary. Hydrogen and (depending on the stoichiometry) amides (salts of \({\text{NH}}_{2}{}^{\text{−}}),\) imides (salts of NH2−), or nitrides (salts of N3−) form.
- The nitrogen atom in ammonia has its lowest possible oxidation state (3−) and thus is not susceptible to reduction. However, it can be oxidized. Ammonia burns in air, giving NO and water. Hot ammonia and the ammonium ion are active reducing agents. Of particular interest are the oxidations of ammonium ion by nitrite ion, \({\text{NO}}_{2}{}^{\text{−}},\) to yield pure nitrogen and by nitrate ion to yield nitrous oxide, N2O.
- There are a number of compounds that we can consider derivatives of ammonia through the replacement of one or more hydrogen atoms with some other atom or group of atoms. Inorganic derivations include chloramine, NH2Cl, and hydrazine, N2H4:
Chloramine, NH2Cl, results from the reaction of sodium hypochlorite, NaOCl, with ammonia in basic solution. In the presence of a large excess of ammonia at low temperature, the chloramine reacts further to produce hydrazine, N2H4:
\({\text{NH}}_{3}(aq)+{\text{OCl}}^{\text{−}}(aq)\;⟶\;{\text{NH}}_{2}\text{Cl}(aq)+{\text{OH}}^{\text{−}}(aq)\)
\({\text{NH}}_{2}\text{Cl}(aq)+{\text{NH}}_{3}(aq)+{\text{OH}}^{\text{−}}(aq)\;⟶\;{\text{N}}_{2}{\text{H}}_{4}(aq)+{\text{Cl}}^{\text{−}}(aq)+{\text{H}}_{2}\text{O}(l)\)
Anhydrous hydrazine is relatively stable in spite of its positive free energy of formation:
\({\text{N}}_{2}(g)+{\text{2H}}_{2}(g)\;⟶\;{\text{N}}_{2}{\text{H}}_{4}(l)\phantom{\rule{5em}{0ex}}{\text{Δ}G}_{\text{f}}^{°}=149.2\;\text{kJ}\;{\text{mol}}^{-1}\)
Hydrazine is a fuming, colorless liquid that has some physical properties remarkably similar to those of H2O (it melts at 2 °C, boils at 113.5 °C, and has a density at 25 °C of 1.00 g/mL). It burns rapidly and completely in air with substantial evolution of heat:
\({\text{N}}_{2}{\text{H}}_{4}(l)+{\text{O}}_{2}(g)\;⟶\;{\text{N}}_{2}(g)+{\text{2H}}_{2}\text{O}(l)\phantom{\rule{5em}{0ex}}\text{Δ}H\text{°}=-621.5\;\text{kJ}\;{\text{mol}}^{-1}\)
Like ammonia, hydrazine is both a Brønsted base and a Lewis base, although it is weaker than ammonia. It reacts with strong acids and forms two series of salts that contain the \({\text{N}}_{2}{\text{H}}_{5}{}^{+}\) and \({\text{N}}_{2}{\text{H}}_{6}{}^{2+}\) ions, respectively. Some rockets use hydrazine as a fuel.
Phosphorus Hydrogen Compounds
The most important hydride of phosphorus is phosphine, PH3, a gaseous analog of ammonia in terms of both formula and structure. Unlike ammonia, it is not possible to form phosphine by direct union of the elements. There are two methods for the preparation of phosphine. One method is by the action of an acid on an ionic phosphide. The other method is the disproportionation of white phosphorus with hot concentrated base to produce phosphine and the hydrogen phosphite ion:
\(\text{AlP}(s)+{\text{3H}}_{3}{\text{O}}^{\text{+}}(aq)\;⟶\;{\text{PH}}_{3}(g)+{\text{Al}}^{3+}(aq)+{\text{3H}}_{2}\text{O}(l)\)
\({\text{P}}_{4}(s)+{\text{4OH}}^{\text{−}}(aq)+{\text{2H}}_{2}\text{O}(l)\;⟶\;{\text{2HPO}}_{3}{}^{2-}(aq)+{\text{2PH}}_{3}(g)\)
Phosphine is a colorless, very poisonous gas, which has an odor like that of decaying fish. Heat easily decomposes phosphine \(({\text{4PH}}_{3}\;⟶\;{\text{P}}_{4}+{\text{6H}}_{2}),\) and the compound burns in air. The major uses of phosphine are as a fumigant for grains and in semiconductor processing.
Like ammonia, gaseous phosphine unites with gaseous hydrogen halides, forming phosphonium compounds like PH4Cl and PH4I. Phosphine is a much weaker base than ammonia; therefore, these compounds decompose in water, and the insoluble PH3 escapes from solution.
Sulfur Hydrogen Compounds
Hydrogen sulfide, H2S, is a colorless gas that is responsible for the offensive odor of rotten eggs and of many hot springs. Hydrogen sulfide is as toxic as hydrogen cyanide; therefore, it is necessary to exercise great care in handling it. Hydrogen sulfide is particularly deceptive because it paralyzes the olfactory nerves; after a short exposure, one does not smell it.
The production of hydrogen sulfide by the direct reaction of the elements (H2 + S) is unsatisfactory because the yield is low. A more effective preparation method is the reaction of a metal sulfide with a dilute acid. For example:
\(\text{FeS}(s)+{\text{2H}}_{3}{\text{O}}^{\text{+}}(aq)\;⟶\;{\text{Fe}}^{2+}(aq)+{\text{H}}_{2}\text{S}(g)+{\text{2H}}_{2}\text{O}(l)\)
It is easy to oxidize the sulfur in metal sulfides and in hydrogen sulfide, making metal sulfides and H2S good reducing agents. In acidic solutions, hydrogen sulfide reduces Fe3+ to Fe2+, \({\text{MnO}}_{4}{}^{-}\) to Mn2+, \({\text{Cr}}_{2}{\text{O}}_{7}{}^{2-}\) to Cr3+, and HNO3 to NO2. The sulfur in H2S usually oxidizes to elemental sulfur, unless a large excess of the oxidizing agent is present. In which case, the sulfide may oxidize to \({\text{SO}}_{3}{}^{2-}\) or \({\text{SO}}_{4}{}^{2-}\) (or to SO2 or SO3 in the absence of water):
\({\text{2H}}_{2}\text{S}(g)+{\text{O}}_{2}(g)\;⟶\;\text{2S}(s)+{\text{2H}}_{2}\text{O}(l)\)
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 H2S 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, S2−, in the second. Since hydrogen sulfide is a weak acid, aqueous solutions of soluble sulfides and hydrogen sulfides are basic:
\({\text{S}}^{2-}(aq)+{\text{H}}_{2}\text{O}(l)\;⇌\;{\text{HS}}^{\text{−}}(aq)+{\text{OH}}^{\text{−}}(aq)\)
\({\text{HS}}^{\text{−}}(aq)+{\text{H}}_{2}\text{O}(l)\;⇌\;{\text{H}}_{2}\text{S}(g)+{\text{OH}}^{\text{−}}(aq)\)
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, CaF2, and concentrated sulfuric acid:
\({\text{CaF}}_{2}(s)+{\text{H}}_{2}{\text{SO}}_{4}(aq)\;⟶\;{\text{CaSO}}_{4}(s)+\text{2HF}(g)\)
Gaseous hydrogen fluoride is also a by-product in the preparation of phosphate fertilizers by the reaction of fluoroapatite, Ca5(PO4)3F, 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:
\({\text{H}}_{3}{\text{PO}}_{4}(l)+{\text{Br}}^{\text{−}}(aq)\;⟶\;\text{HBr}(g)+{\text{H}}_{2}{\text{PO}}_{4}{}^{\text{−}}(aq)\)
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, PbCl2, and Hg2Cl2 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:
\({\text{Cl}}_{2}(aq)+{\text{2e}}^{-}\;⟶\;{\text{2Cl}}^{\text{−}}(aq)\phantom{\rule{5em}{0ex}}E\text{°}=1.36\;\text{V}\)
\({\text{Br}}_{2}(aq)+{\text{2e}}^{-}\;⟶\;{\text{2Br}}^{\text{−}}(aq)\phantom{\rule{5em}{0ex}}E\text{°}=1.09\;\text{V}\)
\({\text{I}}_{2}(aq)+{\text{2e}}^{-}\;⟶\;{\text{2I}}^{\text{−}}(aq)\phantom{\rule{5em}{0ex}}E\text{°}=0.54\;\text{V}\)
For example, bromine oxidizes iodine:
\({\text{Br}}_{2}(aq)+\text{2HI}(aq)\;⟶\;\text{2HBr}(aq)+{\text{I}}_{2}(aq)\phantom{\rule{5em}{0ex}}E\text{°}=0.55\;\text{V}\)
Hydrofluoric acid is unique in its reactions with sand (silicon dioxide) and with glass, which is a mixture of silicates:
\({\text{SiO}}_{2}(s)+\text{4HF}(aq)\;⟶\;{\text{SiF}}_{4}(g)+{\text{2H}}_{2}\text{O}(l)\)
\({\text{CaSiO}}_{3}(s)+\text{6HF}(aq)\;⟶\;{\text{CaF}}_{2}(s)+{\text{SiF}}_{4}(g)+{\text{3H}}_{2}\text{O}(l)\)
The volatile silicon tetrafluoride escapes from these reactions. Because hydrogen fluoride attacks glass, it can frost or etch glass and is used to etch markings on thermometers, burets, and other glassware.
The largest use for hydrogen fluoride is in production of hydrochlorofluorocarbons for refrigerants, in plastics, and in propellants. The second largest use is in the manufacture of cryolite, Na3AlF6, which is important in the production of aluminum. The acid is also important in the production of other inorganic fluorides (such as BF3), which serve as catalysts in the industrial synthesis of certain organic compounds.
Hydrochloric acid is relatively inexpensive. It is an important and versatile acid in industry and is important for the manufacture of metal chlorides, dyes, glue, glucose, and various other chemicals. A considerable amount is also important for the activation of oil wells and as pickle liquor—an acid used to remove oxide coating from iron or steel that is to be galvanized, tinned, or enameled. The amounts of hydrobromic acid and hydroiodic acid used commercially are insignificant by comparison.