Peroxide

General formula of a peroxide.

The peroxides are substances that present an oxygen-oxygen bond and that contain oxygen in the oxidation state −2. The general formula of peroxides is Metal + O₂^2-. They generally behave as oxidizing substances.

In contact with combustible material they can cause fires or even explosions. However, when faced with strong oxidants such as permanganate, they can act as reductants, oxidizing to elemental oxygen. It is important to note that peroxide is charged.

In a few words, they are oxides that have a greater amount of oxygen than a normal oxide and in their structure they show a simple non-polar covalent bond between oxygen and oxygen.

History

As probably the first synthetically produced peroxide compound, Alexander v. Humboldt in 1799 in an attempt to decompose air with barium peroxide. It was not until 19 years later that Thénard realized that this salt could be used to produce a previously unknown compound he called oxidized water, now known as hydrogen peroxide.In 1811 Thénard and Gay-Lussac presented the first peroxide. In sustained research efforts over the next several decades, hydrogen peroxide and its salts were investigated. In the search for a use, the bleaching effect of the compound on natural dyes was recognized early on. However, industrial use initially failed because only weakly concentrated and contaminated barium peroxide solutions could be produced. The first industrial plant for the synthesis of hydrogen peroxide was built in Berlin in 1873. Only after the discovery of the synthesis of hydrogen peroxide by electrolysis of sulfuric acid could improved processes be developed on an electrochemical basis. The first factory to use this method was opened in 1908 in Weißenstein in Carinthia. The anthraquinone process that is still in use today was developed by IG Farben in Ludwigshafen in the 1930s. With more modern synthesis processes and the expansion of the field of application, the annual production of hydrogen peroxide increased sharply from 35,000 t in 1950 to 100,000 t in 1960 to 300,000 t in 1970. In 1998 there was a world production capacity of 2,700,000 t per year.

Occurrence

In the environment

The mineral studtite containing oxide on uranofane.

Peroxides are usually very reactive, so there are few natural occurrences. In addition to hydrogen peroxide, this includes some natural plant substances such as a prostaglandin peroxide derivative and ascaridol. Hydrogen peroxide occurs naturally in surface water, groundwater, and in the Earth's atmosphere. The formation takes place here through the action of light or natural catalytically active substances from the water. Seawater contains 0.5 to 14 μg/l, fresh water 1 to 30 μg/l, and air 0.1 to 1 ppb.

Two minerals are known to contain peroxide, studtite and metastudite. These are uranyl peroxides with different amounts of water of crystallization in the structure. Unstable peroxide is created during the radiolysis of water by alpha radiation from uranium. Apart from natural uranium deposits, these compounds also form on the surface of radioactive waste and therefore their stability could be important for the ultimate storage of uranium waste.

Summary

The best known peroxide and main starting compound in the synthesis of other peroxides is hydrogen peroxide (H₂O₂). Today it is usually obtained by autoxidation of naphthohydroquinone. Formerly, the formation of barium peroxide or the hydrolysis of persulfates were used, which in turn were generated by electrolysis of sulfates in aqueous solution with high current densities per electrode surface.

Many organic substances can be converted to hydroperoxides in autoxidation reactions in the presence of light and atmospheric oxygen. Especially dangerous is the formation from ethers since these are transformed very easily and the peroxides are usually enriched in the residue of a subsequent distillation. There they can produce very strong explosions. Many of the most tragic laboratory accidents are due to this type of reaction. Therefore, before distilling larger amounts of these solvents, the presence of peroxides must be tested with paper impregnated with potassium iodide and starch. The formation of a bluish or dark color indicates the presence of peroxide. (The peroxide oxidizes the iodide to elemental iodine, which, in turn, forms an inclusion complex of the characteristic dark color with the starch.)

Properties

The O-O bond is unstable due to the degree of oxidation of oxygen equal to -1 (exception). Therefore, the functional group is very reactive and can react as an oxidant (most common case) or as a reducing agent to achieve more stable degrees of oxidation. Another property of this group is its ability to form radicals by homolytic cleavage of the O-O bond. This cleavage can be initiated thermally, catalytically, or by UV. A peroxide is characterized in particular by its half-life temperature T ½ (given by a half-life t ½ of 10 h and 1 h), its mass rate of active oxygen, its self-accelerating decomposition temperature (SADT), its maximum temperature storage and use temperature range; some parameters allow classification according to stability.

Some molecules are considered very dangerous because they present significant fire and explosion hazards. Some grades incorporate a phlegmator. It is generally recommended to store peroxides separately to avoid any reaction with other molecules.

Uses

Analysis parameter in the food industry

In the food industry it is used as a parameter to measure the quality of oils and fats, which are susceptible to becoming rancid or decomposing, using the technique of measuring the peroxide index (or peroxidation index); peroxidation being one of the causes that causes rejectable characteristics in the quality of food (such as oils, among others).

Chemical applications

Likewise, the applications of peroxides are very versatile (from the hairdresser's where they are used in dyes to lighten hair to rocket fuel).

In the chemical industry they are used to obtain epoxides, in various oxidation reactions, as initiators of radical reactions, for example, to harden polyesters or in the manufacture of glycerol from hydroxypropene alcohol. Peroxy-trifluoroacetic acid (F₃C–C(=O)–O–O–H) is a very powerful disinfectant and is used as such in the pharmaceutical industry. In dentistry it is used for teeth whitening, either applied in gel or in strips impregnated with peroxide in concentrations of 9%, 16% and 25%.

Oxidation state

The oxidation state of oxygen in peroxide groups is -1

Analytics

Peroxides give an orange coloration with solutions of titanium oxide in concentrated sulfuric acid.

With potassium dichromate they form blue chromium (VI) peroxide that can be extracted with ethyl ether.