Peroxides are compounds in which there is an oxygen-oxygen bond, -O-O-. Oxygen atoms stick together in pairs to form oxygen gas, O2, and in threes to make ozone, O3, but do not form longer chains, as sulphur does so readily. These gases are not considered to be peroxides. With hydrogen, oxygen forms the remarkable compound H2O, water. In 1818, Thénard discovered that it made another hydride, H2O2, hydrogen peroxide, which was similar to water but had peculiar properties. This article will look at these oxygen compounds, with an emphasis on hydrogen peroxide.

Oxygen is a simple molecule, with nuclear charge +6 and electron configuration 1s22s22p4. Two of the p-electrons usually form a lone pair, and the remaining two are available for making bonds. The oxygen atom is a very different animal from the diatomic oxygen molecule. The atom, a free radical, is very active, while the molecule is sleepy and inert. The oxygen molecule is peculiar, however. Although it is generally represented with a double bond, O=O, this cannot be the whole truth, since oxygen gas is paramagnetic, having a magnetic moment, and so must have unpaired electrons. The bond is actually a single bond with two electrons in a bonding orbital, plus two three-electron bonds, which accounts for the eight valence electrons that are available, as well as the unpaired electrons. The bond length is 0.121 nm, shorter than an O-O single bond, and requiring 5.08 eV to break, or 117 kcal/mole. Because of this, O2 is very stable, remaining whole until high in the atmosphere. It is still active enough to support combustion, however, when it is picked apart by free radicals from the fuel.

Short-wavelength (less than about 250 nm) ultraviolet radiation from the sun can dissociate diatomic oxygen, creating free oxygen atoms that can stick to an oxygen molecule to form ozone, O3. The name comes from the Greek ozw, "I smell," alluding to the sharp and refreshing odor of the gas. Ozone can often be smelled around electrical machinery, since it is made in electrical sparks, and in the vicinity of short-wavelength UV. Because it was a bactericide, it was considered beneficial, and people sought out resorts that advertised ozone in their air. However, it is irritating, and there is no benefit to breathing it at all. In fact, it combines with nitrogen oxides and unburned hydrocarbons to form an irritating component of smog, peroxyacetyl nitrate, called PAN. To its credit, it absorbs ultraviolet at longer, but still dangerous, wavelengths than oxygen or nitrogen, and a thin layer of it around 25-35 km altitude shields the surface of the earth. Chlorine ions are very effective in combining with ozone, enlarging the "ozone hole" that appears in polar regions. Ozone has the form of an isosceles triangle, with an apicial angle of 116.8° and bond length 0.128 nm. The bond is a resonance between single and double bonds, which helps to stabilize the molecule. The heat of formation of ozone is 34.4 kcal/mole, so it is certainly a stable molecule. However, it is tempting for it to give up an oxygen free radical and form the even more stable O2, so it is an excellent oxidizing agent in the strict sense, making a good bleach. Liquid ozone is a dark blue-black, they say.

We must mention a few facts about water here for comparison with the other compounds. Water is an extraordinary compound, liquid between 0°C and 100°C, when most other compounds of such small molecular weight (18) are gases. The length of the O-H bonds is 0.0965 nm, and the angle between them is 104° 31'. This angle is the p-orbital angle of 90° increased by the effect of the repulsion between the protons. The molecule has an electric dipole of 1.84 debye (a debye is 10-10 esu of charge times 10-8 cm, or 10-18 esu, the expected order of magnitude of molecular dipole moments). The lone pair of electrons are as close as they can get to the protons, but still the dipole moment is large. The attraction of a proton on one water molecule for the lone pair of electrons on another leads to strong hydrogen bonding in water. Water has a complex structure involving dodecahedral arrangements that is responsible for its unusual physical properties. The dipoles of water make possible the tearing down of ionic crystals as they cradle and shield the ions, so that salts are soluble in water. Water ionizes slightly to make hydrogen and hydroxyl ions available, HOH → H+ + OH-. Most traditional chemistry is chemistry in aqueous solutions. The heat of formation of water is 68.32 kcal/mole. Note that the heats of formation quoted are from the diatomic elements, so they include the expensive breaking of those bonds to form the new ones. Water is only a clear liquid to our eyes; it absorbs light strongly in the infrared and ultraviolet, and is clear only in the narrow visual spectrum. Our eyes evolved under water.

Finally we come to our main subject, hydrogen peroxide. The photograph in the title shows a model of its structure. It would be better expressed as (OH)2 than as H2O2. The angles between the O-O bond and the O-H bonds is 103°, again not far from the ideal 90° for p-orbitals. The O-H bonds are not in the same plane, but in orthogonal planes. This shows that there is no free rotation around the O-O bond, whose length is 0.149 nm, considerably longer than the stronger bonds in O2 and O3. Its heat of formation is 44.8 kcal/mole, so it is stable if on its own. It is very much like water, melting at -0.89°C and boiling at 151.4°C. If anything, hydrogen bonding is even stronger in peroxide than in water. The dipole moment is 2.1 debye, a bit larger than water's. Its density is 1.438 g/cc, half again as heavy as water, and more viscous, since it is described as "syrupy." Remarkably, its index of refraction is exactly equal to that of water, 1.333. Hydrogen peroxide is one of the few interesting chemicals that can easily be bought over the counter, in the form of a 3% aqueous solution intended for use as an antiseptic. This solution stores well, and is useful in the laboratory. A 6% solution is sold for bleaching hair, while the standard industrial concentration is 27.5%. An 87% solution can also be manufactured for special uses. Hydrogen peroxide is shipped in tank-car lots as an industrial chemical. Hydrogen peroxide also occurs in small amounts in honey, making it a sticky antiseptic for wounds. It dissolves readily in water, alcohol and ether.

Hydrogen peroxide can decompose into water and oxygen, H2O2 → H2O + (1/2)O2, releasing 26.04 kcal/mole, or 766 cal/g, or 1380 Btu/lb. It does not do this spontaneously, since the reaction involves creating oxygen atoms, which is energetically very unfavorable. That is, there is a high barrier to the progress of the reaction, although the final result releases energy. The reaction is catalyzed by metals, most acids (phosphoric seems to preserve the peroxide), and even surface roughness. Kept cool in the dark, in a clean, smooth bottle, peroxide will survive. In ether solution, peroxide is more stable than in aqueous solution. If strongly heated, or exposed to a platinum catalyst, concentrated peroxide will decompose explosively into steam and oxygen.

The bleaching action of hydrogen peroxide results from its attack on the electron-rich alternating single and double bonds in the chromophores of dye molecules. It splits a double bond and replaces it by an epoxy link, rendering the chromophore colorless. An epoxy link is an -O- bridge between two carbon atoms that are already bound to each other.

Ultraviolet light can create hydrogen peroxide photochemically from water and oxygen by the inverse of the reaction above. This can also contribute to smog and the production of PAN, in addition to photochemical ozone. The odor of hydrogen peroxide is irritating, and it has an astringent taste. Nevertheless, the 3% solution is used as a gargle, so it is definitely not a poison. Swallowing it may lead to greater difficulties, but probably no more than internal irritation. One of the great merits of peroxide is that it leaves no hazardous residues.

87% hydrogen peroxide, or HTP (high-test peroxide), can be used as a monopropellant, releasing heat when it encounters a catalyst and producing hot steam mixed with oxygen. This was the source of the power for the fuel pumps of the V-2 rocket missile of World War II, whose main fuel was alcohol and liquid oxygen. The first catalyst was calcium permanganate, but later silver mesh proved superior. The industrial process for making concentrated hydrogen peroxide, now widely used, was developed in connection with this use. An exhaust speed of 4060 ft/sec can be produced by hydrogen peroxide, which is a specific impulse of 126 sec., rather low but useful for special purposes, such as guidance jets. When mixed with JP-4 (jet fuel, similar to kerosene) to take advantage of the oxygen, the specific impulse rises to 200-230 sec., a creditable amount.

Some Lewis structures for peroxides are shown at the right. The diagram for O2 is inadequate, but does show the two unpaired electrons. The O-- ion does not exist in aqueous solution, and neither do the O2- or O2--ions, but they do exist in crystals and make salts, called superoxides and peroxides. We'll see below why they do not survive in water. Potassium superoxide, KO2, often known as potassium tetroxide, is an example of a superoxide. In water, it decomposes to hydrogen peroxide, oxygen, and potassium hydroxide. There are many peroxides, such as sodium peroxide, Na2O2, and barium peroxide BaO2, whose structure is shown in the figure. When heated, peroxides evolve oxygen, but in an acid solution yield hydrogen peroxide. Barium peroxide can be formed by heating barium metal in oxygen under pressure. Superoxides and peroxides are all brightly colored. Hydrogen peroxide ionizes slightly as a weak acid, H2O2 → H+ + O2H-, with a dissociation constant of 2.4 x 10-12, somewhat more readily than water.

Taking the oxidation number of hydrogen as +1, the oxidation number of oxygen in hydrogen peroxide is -1. In superoxides, it is -1/2, but this is peculiar. In O2, the oxidation number is 0, while in water oxygen shows an oxidation number of -2. Therefore, the oxygen in peroxide can be oxidized from -1 to 0, or reduced from -1 to -2. Hydrogen peroxide can be either a reducing agent in the first case, when O2 is evolved, or an oxidizing agent in the second case, when water is formed instead. This remarkable behavior makes hydrogen peroxide a versatile reagent. A typical reaction as an oxidizing agent is Mn(OH)2 + H2O2 → MnO2 + 2H2O, where Mn is oxidized from +2 in the pale pink hydroxide to +4 in the black oxide. Hydrogen peroxide is usually oxidizing in an alkaline solution. On the other hand, it is a reducing agent in 2MnO4- + 5H2O2 + 6H+ → 2Mn++ + 5O2 + 8H2O. Here the permanganate ion, in which Mn is +7, is reduced to manganous ion, where the Mn is +2. The 10 electrons necessary all come from the 5H2O2, where the oxygen goes from -1 to 0. Note the presence of hydrogen ion, which takes care of the oxygen from the permanganate. Reduction generally occurs in acid solution.

When a peroxide of an alkali or alkaline earth metal is dropped into water, a vigorous reaction ensues resulting in the formation of solutions of the hydroxides and the evolution of oxygen gas. The mechanism of this reaction, illustrated at the left, is probably as follows: first, the water dissoves the metal ions, freeing the ion O2--. This ion has a lot of electrons, and is attacked at one end by a proton of a water molecule. In organic chemistry, this is called an electrophilic attack. Electrons are attracted to that end of the ion, and the oxygen at the other end departs, leaving an electron behind, becoming a free radical. One oxygen has been oxidized, the other reduced, in this step. The proton bonds to the oxygen it has approached, leaving the OH- behind. The result is the creation of metal ions, two hydroxyl ions, and a free O. The very reactive free O can either engage in oxidation, or if there is nothing else to do, combine with another to form O2, releasing a good deal of energy. The O2- ion of a superoxide reacts in the same way. Curiously, alkali metals prefer to form the superoxides and peroxides when they react with oxygen. It is more difficult to get the ordinary oxides.

Hydrogen peroxide is used three times in the standard qualitative analysis scheme. First, it is used in HCl solution to make sure that tin is oxidized from Sn++ to Sn++++ to achieve a clean separation of tin with the other members of the Cu-As group. Arsenite is also oxidized to arsenate. When the peroxide has done its job, it can be boiled off so it will not interfere with later procedures. Next, in NaOH solution, it oxidizes Mn++ to Mn++++, Co++ to Co+++ and Cr(OH)4- (+3, pale green) to CrO4-- (+6, yellow), so that the aluminium group can be separated from the nickel group. Finally, it is used in an acidic ether solution to reduce dichromate to the blue peroxide CrO5 in the confirming test for chromium. For quantitative determination of hydrogen peroxide, it is titrated against potassium permanganate solution, which it bleaches.

The sulphur in sulphuric acid is already in its highest oxidation state, +6, so when an oxidizing agent like H2O2 acts on sulphuric acid, the only possibility is the oxidation of the oxygen from -2 to -1, creating peroxy acids, such as are shown in the diagram on the right. These compounds are stable indefinitely in the dry state. They are strong oxidizing agents as well as drying agents. They can be prepared by dissolving SO2 in hydrogen peroxide, or by electrolysis of sulphuric acid at low temperature, since oxidation occurs at the anode. The salts of peroxydisulphuric acid are quite useful. Peroxysulphuric acid is also known as Caro's acid, and is of lesser utility. In water, these compounds yield sulphuric acid and hydrogen peroxide.

Ethyl ether, the common ether that was used as an anaesthetic and is a solvent much used in organic chemistry, presents an additional hazard beyond its volatility and low ignition temperature of the heavy vapor, which is explosive over a wide range of concentration (2%-36%). On long storage, especially in the presence of sunlight, ether, which is C2H5-O-C2H5, forms unstable peroxides. These will explode at the slightest excuse, perhaps breaching the storage vessel and causing a dangerous fire. Barium, strontium, sodium and potassium peroxides should be stored away from combustible material, and water should not be used on fires in their vicinity. They are not dangerous themselves, but only through their contribution of oxygen for combustion, and the strong alkaline solutions they make in water. 27.5% hydrogen peroxide is also an oxygen source, but water can be used on fires in its presence. Solutions of smaller concentration are, as we have said above, no hazard at all.


A BBCi web search for "hydrogen peroxide" turned up 20,200 pages of results, many quite interesting, dealing mainly with either the propellant or the antiseptic applications.

M. J. Sienko and R. A. Plane, Chemical Principles and Properties (New York: McGraw-Hill, 1974). pp. 430-433.

W. N. Jones, Jr., Inorganic Chemistry (Philadelphia: Blakiston, 1947). p. 417 and pp. 456-461.

L. Pauling, General Chemistry (New York: Dover, 1970). pp. 155, 227, 275 and 502.

P. W. Atkins, Molecules (New York: Freeman, 1987). p. 31. A planar molecule model is shown.

C. H. Sorum, Introduction to Semi-Micro Qualitative Analysis, 2nd ed. (New York: Prentice-Hall, 1953).

H. S. Seifert, editor, Space Technology (New York: John Wiley & Sons, 1959). Chapters 14 and 15. Rocket propulsion and propellants.

B. Jaffe, Crucibles: The Story of Chemistry, 4th ed. (New York: Dover, 1978). p. 98.

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Composed by J. B. Calvert
Created 21 January 2003
Last revised 4 September 2012