Oxygen Compounds

Type of physical science: Chemistry

Field of study: Chemical compounds

Compounds in which the element oxygen plays an important chemical or structural part include, in addition to the ubiquitous oxides, a wide variety of inorganic and organic peroxides, many of which find application in chemical synthesis and in industrial processes.

Overview

Oxygen is, by any standard of comparison, the most abundant element in the earth's crust, constituting 47 weight percent, or 63 atom percent, of the crust exclusive of oceans and atmosphere. The compounds of oxygen may be divided into three or four types: first, the oxyanions of dozens of metals and nonmetals, which are eliminated from this discussion because they are better treated under the chemistry of the other element of the anion. Second, the oxides themselves, which are of chemical and sometimes industrial importance. Third, hydrogen peroxide and many inorganic and organic peroxides, which have an interesting chemistry and, in some cases, valuable industrial applications. Fourth, a miscellaneous category that includes ozone and a variety of peroxy radicals that are important in atmospheric chemistry.

Oxides of all the elements are known, except for the lower atomic number rare gases.

This fact led to oxygen being proposed as an atomic weight standard in the first half of the nineteenth century, with an arbitrarily assigned value of 100. (Hydrogen comes out at about 6.3 on this scale.) That is, determination of the weight percent of oxygen and the other element of the oxide would give a direct ratio of atomic weights. This was not possible for any other element, because no other element combined with all the others. The decision to give hydrogen a weight value of one, however, was essentially complete by the mid-nineteenth century.

Oxides of the metals (which constitute the larger part of the elements) are ionic, the O^2- ion combining with a metal ion of a charge of 1+ to 4+ to make a crystal whose structure is dictated by the relative sizes and numbers of positive and negative ions. Because of the relatively high charges of the ions, many of these crystals have very favorable energies of formation, resulting in high melting points and hardnesses, and low solubility in water and other solvents.

Alumina (Al₂O₃), for example, in the form of emery, is used as an abrasive, as is silica sand (SiO₂). Many metal oxides are sufficiently stable to be found in nature and are used as ores in the production of the free metals. Examples include hematite and magnetite, which are used to make iron and steel; bauxite for aluminum; and cassiterite for tin.

Peroxides are oxygen compounds that contain the peroxy linkage, -O-O-. All of them, organic or inorganic, may be thought of as derivatives of hydrogen peroxide, H₂O₂ or H-O-O-H, with one or both hydrogen atoms replaced by an atom, ion, or larger inorganic or organic structure. The metal peroxides in general incorporate the peroxide anion, (O-O)^2-, into an ionic crystal. Metals forming such compounds are the group I and group II elements (alkali metals and alkaline earths), plus a few outliers such as zinc, cadmium, and mercury. The elements of groups III A to VI A also form peroxides, but not as salts of metal ions. Rather, the peroxy linkage is incorporated into the oxyanion(s) of the element. The most important of these peroxyanions are the peroxyborates, with a cyclic diperoxide structure: and the peroxymono- and disulfates, SO₅^2- and S₂O₈^2-, with structures (O₃S-O-O-)^2- and O₃S-O-O-SO₃. (These compounds are often referred to as "perborates" and "persulfates," an unfortunate usage that confuses them with ions such as permanganate and perchlorate, which do not contain the peroxy linkage.) In addition to these relatively well-known peroxy compounds, there are peroxymono- and dicarbonates, CO₄^2- and C₂O₆^2-, with structures (O₂C-O-O-)^2- and (O₂C-O-O-CO₂)^2-; and a peroxydiphosphate with a structure analogous to that of the peroxydisulfate. Although some of these compounds are made by electrolytic oxidation on a platinum anode, most are made by adding concentrated solutions of hydrogen peroxide to the nonperoxy acid (boric, sulfuric, phosphoric, carbonic), or its ammonium, potassium, or sodium salt. The resulting peroxy salts are solids that hydrolyze in water to give back the hydrogen peroxide. As such, they can be thought of as a stable dry powder form in which to deliver hydrogen peroxide, which is otherwise a liquid that gradually decomposes on storage; they are used in this way in many commercial and industrial formulations.

Superoxides, with the ion O^2-, are known. Cations associated with the superoxide ion are those of the alkali metals and alkaline earths. The only superoxide of commercial importance is potassium superoxide, which is used in self-contained breathing apparatus for mines, and undersea and space applications. Moisture in exhaled breath breaks down the compound into oxygen gas as well as potassium hydroxide, which absorbs exhaled carbon dioxide.

A wide variety of organic compounds form peroxides (Org-O-O-Org) or hydroperoxides (Org-O-O-H), some by reaction with hydrogen peroxide, many by direct air oxidation. The more important types of compounds are hydroperoxide, with a structure of ROOH; peroxides, with a structure of ROOR; peroxyacids, with structures of RC(=O)OOH and RSO2OOH; diacyl peroxides, with structures of RC(=O)OOC(=O)R, ROC(=O)OOC(=O)OR', and RSO₂OOSO₂R′, and peroxy esters, with a structure of RC(=O)OOR'.

Hundreds of organic peroxy compounds have been synthesized and characterized, including many that lie outside the above list of types of compounds. Most are less stable than the inorganic peroxides, many decomposing explosively. This may be why only a few organic peroxides have major commercial or industrial applications.

Applications

By far, the most widely used of the peroxides is hydrogen peroxide itself. It is manufactured by electrolytic oxidation of sulfuric acid or ammonium bisulfate to peroxydisulfate, followed by steam hydrolysis and stripping of the H₂O₂ with regeneration of starting materials. It is also made by catalytic oxidation of certain organic substrates, followed by cleavage to recover the peroxide. It is available for industrial uses in concentrations (in water) from about 30 percent up to nearly 100 percent.

Hydrogen peroxide is a strong oxidizing agent in alkaline solution (interestingly, a weak reducing agent in acid solution), and is exploited in its bulk uses. Bleaching, for example, is an oxidative process in which large color-producing molecules are snipped apart at one or more oxidation-sensitive bonds. The fragments cannot absorb visible light as the large molecule could; consequently, they have no color and bleaching has taken place. Tens of thousands of tons of hydrogen peroxide are used annually in mill bleaching of natural-fiber fabrics, some synthetics, and paper pulp. Substantial amounts of hydrogen peroxide are also used in cleaning up municipal and industrial effluents, to oxidize cyanides, sulfides, and other toxic and malodorous materials to harmless forms. About one-third of the annual production in the United States is used to produce other chemicals, mostly the peroxyborates and peroxysulfates, but also a variety of organics and some pharmaceuticals. Pure hydrogen peroxide is used as an oxidant in liquid-fueled rockets.

Sodium peroxide is competitive with hydrogen peroxide in many bleaching and oxidizing applications. It has the advantage of direct manufacture by dry-air oxidation of sodium metal, and somewhat simpler shipping and handling, as a solid, than the liquid hydrogen peroxide. (Both compounds must be kept away from oxidizable materials in storage, to avoid exothermic and sometimes explosive reactions.) The disadvantage of sodium peroxide is that it is corrosive to metal equipment and hydrolyzes in water to equimolar hydrogen peroxide and sodium hydroxide solution; not all applications can withstand this strong base. In addition to its major uses in bleaching and oxidation, the compound is used in small quantities in denture cleansers.

Of the other metal peroxides, magnesium and calcium peroxide are used in antacid formulations, and the calcium compound is used as a dough conditioner in baking, as a seed and grain disinfectant, and in dentifrices. Zinc peroxide is used in pharmacology to treat infections and skin lesions, and in the rubber industry as a cross-linker, filler, and pigment.

The peroxyanions like peroxyborates and peroxydicarbonates find wide application (as the sodium salts) in household laundry bleaches (the "nonchlorine" bleaches). In Europe, detergents contain these bleaching agents as well. They are used in industrial bleaching, and the peroxyborates are found in denture cleansers.

Peroxymonosulfuric acid (Caro's acid) is both a strong acid and a powerful oxidizing agent that is too reactive for widespread industrial application. It is used in Europe for wastewater treatment. Peroxydisulfuric acid, usually as the diammonium or dipotassium salt, is widely used as a free-radical initiator in emulsion polymerization, in metal cleaning and etching, in hair bleaching, in foodstuff bleaching, and as a laboratory reagent.

Although many organic peroxides are known, only a few are stable enough to be produced and consumed in industrial quantities. Peracetic acid is prepared as a 40 percent solution in pure acetic acid by reaction of the acetic acid with hydrogen peroxide or by catalytic air-oxidation of acetaldehyde. It is used in both laboratory and manufacturing processes to make epoxides and glycols, including glycerol. It has also been used as a bactericide and fungicide.

The only other organic peracid produced for commercial sale is m-chloroperbenzoic acid.

Other peracids are produced as needed and used immediately in further reactions.

Of the various diacyl peroxides, (di)benzoyl peroxide is that most commonly found in industrial applications. It is used as a free-radical initiator in nonaqueous polymerization reactions, to make polymers from vinyl compounds, dienes, and the like. (Ammonium peroxydisulfate, mentioned above as an initiator for emulsion polymerizations, is used in emulsions because it is water soluble; benzoyl peroxide is not water soluble, but it is solvent soluble.) Benzoyl peroxide is also used in bleaching flour and has been used to bleach other natural products such as waxes, fats, and oils. It is also found in some acne lotions.

Cumene (isopropylbenzene) hydroperoxide was at one time used as a free-radical initiator; it breaks into phenoxy and isopropoxy free radicals. Left to itself, this breakup (in acid solution) produces phenol and acetone, and this is such a useful industrial synthesis for phenol that cumene hydroperoxide's role as an initiator is all but abandoned.

Di-t-butyl peroxide is added to diesel fuels as an ignition accelerator and has been used as a free radical source in many laboratory reactions.

A troublesome application of hydroperoxides is that they can be formed by autoxidation (with atmospheric oxygen) of many types of organic compounds and the products can be explosive. Chemists have long known that containers of ether, opened and then stored for a long time, should be handled very gingerly for fear of detonation, treated with a reducing agent such as an iron[II] compound to reduce the peroxides and preferably disposed of thereafter.

Horror stories abound of hands lost in explosions of old cans of isopropyl ether, and it is indeed possible to smell peroxy compounds in cans of ethyl ether that have been open only a few months. What is less widely known is that not only ethers, but also alcohols, aldehydes, ketones, some alkenes and alkylbenzenes can form hydroperoxides, usually by inserting the -OOH group on a carbon atom α to the functional group. The lower molecular weight compounds of this type (which by their low weight contain higher percentages of active oxygen) can detonate from heating and in extreme cases merely from jarring. Even the higher molecular weight hydroperoxides can explode if the parent compound is distilled and the peroxy derivative accumulates in the heated still pot. As long-stored reagents are most likely to be found in high school and college chemical stockrooms, this problem has been addressed repeatedly in articles in the JOURNAL OF CHEMICAL EDUCATION, where both preventive and corrective measures are discussed.

Context

In nearly all of their applications, hydrogen peroxide and the inorganic and organic peroxides react by fission of the weak O-O bond. At 50 kilocalories per mole in hydrogen peroxide, the bond has about half the strength of a normal covalent single bond between oxygen and any other element. This leads to two major forms of reaction: the oxidative, in which the bond splits with absorption of two electrons from an oxidizable substrate: H₂O₂+2e^-→2OH^-, and the free-radical, in which the bond cleaves homolytically, usually under the influence of heat, to give two radicals: ROOR→2RO.

The oxidative reaction mode leads to the use of peroxy compounds as oxidizers and bleaches. The only serious chemical rival in these applications is hypochlorite ion, ClO-, available as liquid laundry bleach, bleaching powder, or chlorine water. All of these hypochlorite products have the drawback of side reactions that produce off-color, off-flavor, or even toxic by-products, which can be avoided by using peroxides.

In the free-radical reaction mode, the peroxides are initiators in polymerization reactions, and sometimes reagents in direct syntheses. In the first of these uses, few other types of chemicals can compete economically. This fact, together with the myriad industrial and household uses for hydrogen peroxide and other peroxides, makes peroxides a nearly irreplaceable part of chemistry today.

The other and absolutely irreplaceable oxygen compounds are the oxides. Here, the chemistry is not that of laboratory or industrial synthesis, but that of "how things are." The bulk of the earth's crust and the living things it supports contain oxide-type oxygen, just as a matter of how elements react to form compounds. Oxides produce strong, stable inorganic crystals, and oxide bonding covalently can form bridges between atoms, leading to extended structures.

Hence, oxide oxygen is a constituent of nearly everything people see, even in the middle of cities, where much of what is seen is metal, but almost exclusively in the countryside.

The major minerals of the earth's crust can be thought of as oxides of silicon and aluminum, with a relatively small number of other metal atoms as additions or replacements. The oceans are essentially an oxide layer at the surface of the earth. Humans, as water-based life-forms, are about two-thirds oxygen by weight. The same value applies, more or less, to all the other animal and plant life. If all the compounds containing oxide oxygen were removed, there would be little left but the iron-nickel core of the earth.

Principal terms

AUTOXIDATION: the reaction of a chemical compound with atmospheric oxygen, without deliberate intent or effort to bring about a chemical reaction

BLEACHING: the elimination of color from a material by breaking up color-producing molecules into colorless fragments, with or without removal of the fragments

ELECTROLYSIS: oxidation (at the anode) or reduction (at the cathode) brought about by passing a direct current through an electrically conducting solution of a compound

FREE RADICAL: an atom, molecule, or other chemical structure that contains one or more unpaired electrons

HYDROLYSIS: the reaction of a compound with water in such a way that fragments of the water molecule are added to fragments of the compound to make new compounds

OXIDATION: the removal of electron(s) from an atom, molecule, or other chemical structure; also, removal of hydrogen atoms or addition of oxygen atoms

OXIDIZING AGENT: a substance that brings about oxidation in another substance; the oxidizing agent is reduced in the process

OXYANION: a compound anion consisting of a central atom surrounded by oxygen atoms, with the whole structure possessing a negative charge

POLYMERIZATION: any reaction in which a small molecule adds to another of its kind, a third molecule adds to the pair, a fourth adds to the trio, and so on, until a very large molecule (the polymer) is built up

Bibliography

Greenwood, N. N., and A. Earnshaw. CHEMISTRY OF THE ELEMENTS. Elmsford, N.Y.: Pergamon Press, 1984. Chapter 14 deals with oxygen and its compounds and has an excellent discussion of hydrogen peroxide. Other peroxyanions are treated under the central elements such as sulfur and phosphorus. Advanced; technical.

Hall, Richard E. "Peroxides and Peroxy Compounds, Inorganic." In KIRK-OTHMER ENCYCLOPEDIA OF CHEMICAL TECHNOLOGY. 3d ed. Vol. 17. New York: Wiley-Interscience, 1982. Contains a thorough discussion of inorganic peroxides, including industrial and economic aspects. Little discussion about hydrogen peroxide. Technical.

Moeller, Therald. INORGANIC CHEMISTRY: AN ADVANCED TEXTBOOK. New York: Wiley, 1952. Includes discussions on hydrogen peroxide and peroxides. Thorough, yet advanced.

Philips, C. S. G., and R. J. Williams. INORGANIC CHEMISTRY. Vol. 1 in PRINCIPLES AND NONMETALS. New York: Oxford University Press, 1965. Chapter 13, "Oxygen and Oxides," deals very thoroughly with oxide compounds, but has little discussion about peroxides. Technical and advanced, but the chapter introduction contains much material that is readily understandable.

Swern, Daniel, ed. ORGANIC PEROXIDES. 3 vols. Reprint. New York: Wiley-Interscience, 1980. Although dated, this source is treated as authority on the subject. Not yet supplanted.

Peroxyborates, with a cyclic diperoxide structure

The Chemistry of Water Pollution