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H2 o2 equation. H2O2 - what is this substance




The well-known formula of the basis of life - water. Its molecule consists of two hydrogen atoms and one oxygen, which is written as H2O. If there is twice as much oxygen, then a completely different substance will turn out - H2O2. What is it and how will the resulting substance differ from its “relative” of water?

H2O2 - what is this substance?

Let's dwell on it in more detail. H2O2 is the formula for hydrogen peroxide, yes, the one used to treat scratches, white. Hydrogen peroxide H2O2 - scientific.

A 3% peroxide solution is used for disinfection. In pure or concentrated form, it causes chemical burns to the skin. A thirty percent peroxide solution is otherwise called perhydrol; it was previously used in hairdressing salons to bleach hair. The skin burned by him also becomes white.

Chemical properties of H2O2

Hydrogen peroxide is a colorless liquid with a "metallic" taste. It is a good solvent and is easily soluble in water, ether, alcohols.

Three and six percent peroxide solutions are usually prepared by diluting a thirty percent solution. When concentrated H2O2 is stored, the substance decomposes with the release of oxygen, so it should not be stored in tightly sealed containers in order to avoid an explosion. With a decrease in the concentration of peroxide, its stability increases. Also, to slow down the decomposition of H2O2, various substances can be added to it, for example, phosphoric or salicylic acid. To store solutions of strong concentration (more than 90 percent), sodium pyrophosphate is added to the peroxide, which stabilizes the state of the substance, and aluminum vessels are also used.

H2O2 in chemical reactions can be both an oxidizing agent and a reducing agent. More often, however, peroxide exhibits oxidizing properties. Peroxide is considered to be an acid, but a very weak one; salts of hydrogen peroxide are called peroxides.

as a method of obtaining oxygen

The decomposition reaction of H2O2 occurs when a substance is exposed to high temperature (more than 150 degrees Celsius). The result is water and oxygen.

Reaction formula - 2 H2O2 + t -> 2 H2O + O2

The oxidation state of H in H 2 O 2 and H 2 O \u003d +1.
The oxidation state of O: in H 2 O 2 \u003d -1, in H 2 O \u003d -2, in O 2 \u003d 0
2 O -1 - 2e -> O2 0

O -1 + e -> O -2
2 H2O2 = 2 H2O + O2

Decomposition of hydrogen peroxide can also occur at room temperature if a catalyst (a chemical that speeds up the reaction) is used.

In laboratories, one of the methods for obtaining oxygen, along with the decomposition of berthollet salt or potassium permanganate, is the reaction of peroxide decomposition. In this case, manganese (IV) oxide is used as a catalyst. Other substances that accelerate the decomposition of H2O2 are copper, platinum, sodium hydroxide.

The history of the discovery of peroxide

The first steps towards the discovery of peroxide were made in 1790 by the German Alexander Humboldt, when he discovered the transformation of barium oxide into peroxide when heated. That process was accompanied by the absorption of oxygen from the air. Twelve years later, the scientists Tenard and Gay-Lussac conducted an experiment on the combustion of alkali metals with an excess of oxygen, as a result of which sodium peroxide was obtained. But hydrogen peroxide was obtained later, only in 1818, when Louis Tenard studied the effect of acids on metals; for their stable interaction, a low amount of oxygen was needed. Conducting a confirmatory experiment with barium peroxide and sulfuric acid, the scientist added water, hydrogen chloride and ice to them. After a short time, Tenar found small solidified drops on the walls of the container with barium peroxide. It became clear that it was H2O2. Then they gave the resulting H2O2 the name "oxidized water". This was hydrogen peroxide - a colorless, odorless, hardly evaporable liquid that dissolves other substances well. The result of the interaction of H2O2 and H2O2 is a dissociation reaction, the peroxide is soluble in water.

An interesting fact is that the properties of the new substance were quickly discovered, allowing it to be used in restoration work. Tenard himself, using peroxide, restored the painting by Raphael, which had darkened with time.

Hydrogen peroxide in the 20th century

After a thorough study of the resulting substance, it began to be produced on an industrial scale. At the beginning of the twentieth century, an electrochemical technology for the production of peroxide was introduced, based on the electrolysis process. But the shelf life of the substance obtained by this method was small, about a couple of weeks. Pure peroxide is unstable, and for the most part it was produced in a thirty percent concentration for bleaching fabrics and in three or six percent for domestic use.

Scientists in Nazi Germany used peroxide to create a liquid fuel rocket engine that was used for defense purposes in World War II. As a result of the interaction of H2O2 and methanol / hydrazine, a powerful fuel was obtained, on which the aircraft reached speeds of more than 950 km / h.

Where is H2O2 used now?

  • in medicine - for the treatment of wounds;
  • in the pulp and paper industry, the bleaching properties of the substance are used;
  • in the textile industry, natural and synthetic fabrics, furs, wool are bleached with peroxide;
  • as rocket fuel or its oxidizer;
  • in chemistry - to produce oxygen, as a foaming agent for the production of porous materials, as a catalyst or hydrogenating agent;
  • for the production of disinfectants or cleaning agents, bleaches;
  • for bleaching hair (this is an outdated method, since the hair is severely damaged by peroxide);

Hydrogen peroxide can be successfully used to solve various household problems. But only 3% hydrogen peroxide can be used for these purposes. Here are some ways:

  • To clean surfaces, pour peroxide into a container with a spray bottle and spray on contaminated areas.
  • To disinfect objects, they must be wiped with an undiluted solution of H2O2. This will help cleanse them of harmful microorganisms. Sponges for washing can be soaked in water with peroxide (proportion 1:1).
  • To bleach fabrics when washing white things, add a glass of peroxide. You can also rinse white fabrics in water mixed with a glass of H2O2. This method restores whiteness, prevents fabrics from yellowing and helps remove stubborn stains.
  • To combat mold and mildew, mix peroxide and water in a spray bottle in a ratio of 1:2. Spray the resulting mixture onto infected surfaces and clean them with a brush or sponge after 10 minutes.
  • You can update the darkened grout in the tile by spraying peroxide on the desired areas. After 30 minutes, you need to carefully rub them with a stiff brush.
  • To wash dishes, add half a glass of H2O2 to a full basin of water (or a sink with a closed drain). Cups and plates washed in such a solution will shine with cleanliness.
  • To clean your toothbrush, you need to dip it in an undiluted 3% peroxide solution. Then rinse under strong running water. This method disinfects the hygiene item well.
  • To disinfect purchased vegetables and fruits, spray a solution of 1 part peroxide and 1 part water on them, then rinse them thoroughly with water (can be cold).
  • In the suburban area with the help of H2O2, you can fight plant diseases. You need to spray them with a peroxide solution or soak the seeds shortly before planting in 4.5 liters of water mixed with 30 ml of forty percent hydrogen peroxide.
  • To revive aquarium fish, if they are poisoned by ammonia, suffocated when aeration is turned off, or for another reason, you can try placing them in water with hydrogen peroxide. It is necessary to mix 3% peroxide with water at the rate of 30 ml per 100 liters and place it in the resulting mixture of lifeless fish for 15-20 minutes. If they do not come to life during this time, then the remedy did not help.

Even as a result of vigorous shaking of a water bottle, a certain amount of peroxide is formed in it, since the water is saturated with oxygen during this action.

Fresh fruits and vegetables also contain H2O2 until they are cooked. During heating, boiling, roasting and other processes with an accompanying high temperature, a large amount of oxygen is destroyed. That is why cooked foods are considered not so useful, although some amount of vitamins remains in them. Freshly squeezed juices or oxygen cocktails served in sanatoriums are useful for the same reason - due to oxygen saturation, which gives the body new strength and cleanses it.

The dangers of peroxide when ingested

After the above, it may seem that peroxide can be specifically taken orally, and this will benefit the body. But that's not the case at all. In water or juices, the compound is found in minimal amounts and is closely related to other substances. Taking “unnatural” hydrogen peroxide inside (and all peroxide bought in a store or produced as a result of chemical experiments on your own cannot be considered natural in any way, besides, it has too high a concentration compared to natural) can lead to life-threatening and health-threatening consequences. To understand why, you need to turn to chemistry again.

As already mentioned, under certain conditions, hydrogen peroxide breaks down and releases oxygen, which is an active oxidizing agent. can occur when H2O2 collides with peroxidase, an intracellular enzyme. The use of peroxide for disinfection is based on its oxidizing properties. So, when a wound is treated with H2O2, the released oxygen destroys the living pathogenic microorganisms that have entered it. It has the same effect on other living cells. If you treat intact skin with peroxide, and then wipe the area with alcohol, you will feel a burning sensation, which confirms the presence of microscopic damage after peroxide. But with the external use of peroxide at a low concentration, there will be no noticeable harm to the body.

Another thing, if you try to take it inside. That substance, which is capable of damaging even relatively thick skin from the outside, enters the mucous membranes of the digestive tract. That is, chemical mini-burns occur. Of course, the released oxidizing agent - oxygen - can also kill harmful microbes. But the same process will occur with the cells of the alimentary tract. If burns as a result of the action of an oxidizing agent are repeated, then atrophy of the mucous membranes is possible, and this is the first step towards cancer. The death of intestinal cells leads to the inability of the body to absorb nutrients, this explains, for example, weight loss and the disappearance of constipation in some people who practice peroxide "treatment".

Separately, it must be said about such a method of using peroxide as intravenous injections. Even if for some reason they were prescribed by a doctor (this can only be justified in case of blood poisoning, when there are no other suitable drugs available), then under medical supervision and with a strict calculation of dosages, there are still risks. But in such an extreme situation, it will be a chance for recovery. In no case should you prescribe yourself injections of hydrogen peroxide. H2O2 poses a great danger to blood cells - erythrocytes and platelets, as it destroys them when it enters the bloodstream. In addition, a deadly blockage of blood vessels by released oxygen can occur - a gas embolism.

Safety measures in handling H2O2

  • Keep out of the reach of children and incapacitated persons. The lack of smell and pronounced taste makes peroxide especially dangerous for them, as large doses can be taken. If the solution is ingested, the consequences of use can be unpredictable. You must immediately consult a doctor.
  • Peroxide solutions with a concentration of more than three percent cause burns if it comes into contact with the skin. The burn area should be washed with plenty of water.

  • Do not allow the peroxide solution to get into the eyes, as their swelling, redness, irritation, and sometimes pain are formed. First aid before going to the doctor - plentiful rinsing of the eyes with water.
  • Store the substance in such a way that it is clear that it is H2O2, that is, in a container with a sticker to avoid accidental misuse.
  • Storage conditions that extend its life are a dark, dry, cool place.
  • Do not mix hydrogen peroxide with any liquids other than pure water, including chlorinated tap water.
  • All of the above applies not only to H2O2, but to all preparations containing it.

Water (hydrogen oxide) is a binary inorganic compound with the chemical formula H 2 O. The water molecule consists of two hydrogen atoms and one oxygen, which are interconnected by a covalent bond.

Hydrogen peroxide.


Physical and chemical properties

The physical and chemical properties of water are determined by the chemical, electronic and spatial structure of H 2 O molecules.

The H and O atoms in the H 2 0 molecule are in their stable oxidation states, respectively +1 and -2; therefore, water does not exhibit pronounced oxidizing or reducing properties. Please note: in metal hydrides, hydrogen is in the -1 oxidation state.



The H 2 O molecule has an angular structure. H-O bonds are very polar. There is an excess negative charge on the O atom, and excess positive charges on the H atoms. In general, the H 2 O molecule is polar, i.e. dipole. This explains the fact that water is a good solvent for ionic and polar substances.



The presence of excess charges on H and O atoms, as well as unshared electron pairs at O ​​atoms, causes the formation of hydrogen bonds between water molecules, as a result of which they combine into associates. The existence of these associates explains the anomalously high values ​​of mp. etc. kip. water.

Along with the formation of hydrogen bonds, the result of the mutual influence of H 2 O molecules on each other is their self-ionization:
in one molecule, a heterolytic break of the polar O-H bond occurs, and the released proton joins the oxygen atom of another molecule. The resulting hydroxonium ion H 3 O + is essentially a hydrated hydrogen ion H + H 2 O, therefore, the water self-ionization equation is simplified as follows:


H 2 O ↔ H + + OH -


The dissociation constant of water is extremely small:



This indicates that water very slightly dissociates into ions, and therefore the concentration of undissociated H 2 O molecules is almost constant:




In pure water, [H + ] = [OH - ] = 10 -7 mol / l. This means that water is a very weak amphoteric electrolyte that exhibits neither acidic nor basic properties to a noticeable degree.
However, water has a strong ionizing effect on the electrolytes dissolved in it. Under the action of water dipoles, polar covalent bonds in the molecules of solutes turn into ionic ones, the ions are hydrated, the bonds between them are weakened, resulting in electrolytic dissociation. For example:
HCl + H 2 O - H 3 O + + Cl -

(strong electrolyte)


(or excluding hydration: HCl → H + + Cl -)


CH 3 COOH + H 2 O ↔ CH 3 COO - + H + (weak electrolyte)


(or CH 3 COOH ↔ CH 3 COO - + H +)


According to the Bronsted-Lowry theory of acids and bases, in these processes, water exhibits the properties of a base (proton acceptor). According to the same theory, water acts as an acid (proton donor) in reactions, for example, with ammonia and amines:


NH 3 + H 2 O ↔ NH 4 + + OH -


CH 3 NH 2 + H 2 O ↔ CH 3 NH 3 + + OH -

Redox reactions involving water

I. Reactions in which water plays the role of an oxidizing agent

These reactions are possible only with strong reducing agents, which are able to reduce the hydrogen ions that are part of the water molecules to free hydrogen.


1) Interaction with metals


a) Under normal conditions, H 2 O interacts only with alkali. and alkali-earth. metals:


2Na + 2H + 2 O \u003d 2NaOH + H 0 2


Ca + 2H + 2 O \u003d Ca (OH) 2 + H 0 2


b) At high temperatures, H 2 O also reacts with some other metals, for example:


Mg + 2H + 2 O \u003d Mg (OH) 2 + H 0 2


3Fe + 4H + 2 O \u003d Fe 2 O 4 + 4H 0 2


c) Al and Zn displace H 2 from water in the presence of alkalis:


2Al + 6H + 2 O + 2NaOH \u003d 2Na + 3H 0 2


2) Interaction with non-metals having low EO (reactions occur under harsh conditions)


C + H + 2 O \u003d CO + H 0 2 ("water gas")


2P + 6H + 2 O \u003d 2HPO 3 + 5H 0 2


In the presence of alkalis, silicon displaces hydrogen from water:


Si + H + 2 O + 2NaOH \u003d Na 2 SiO 3 + 2H 0 2


3) Interaction with metal hydrides


NaH + H + 2 O \u003d NaOH + H 0 2


CaH 2 + 2H + 2 O \u003d Ca (OH) 2 + 2H 0 2


4) Interaction with carbon monoxide and methane


CO + H + 2 O \u003d CO 2 + H 0 2


2CH 4 + O 2 + 2H + 2 O \u003d 2CO 2 + 6H 0 2


Reactions are used in industry to produce hydrogen.

II. Reactions in which water acts as a reducing agent

These reactions are possible only with very strong oxidizing agents that are capable of oxidizing oxygen CO CO -2, which is part of water, to free oxygen O 2 or to peroxide anions 2-. In an exceptional case (in reaction with F 2), oxygen is formed with c o. +2.


1) Interaction with fluorine


2F 2 + 2H 2 O -2 \u003d O 0 2 + 4HF



2F 2 + H 2 O -2 \u003d O +2 F 2 + 2HF


2) Interaction with atomic oxygen


H 2 O -2 + O \u003d H 2 O - 2


3) Interaction with chlorine


At high T, a reversible reaction occurs


2Cl 2 + 2H 2 O -2 \u003d O 0 2 + 4HCl

III. Reactions of intramolecular oxidation - reduction of water.

Under the action of an electric current or high temperature, water can be decomposed into hydrogen and oxygen:


2H + 2 O -2 \u003d 2H 0 2 + O 0 2


Thermal decomposition is a reversible process; the degree of thermal decomposition of water is low.

Hydration reactions

I. Hydration of ions. The ions formed during the dissociation of electrolytes in aqueous solutions attach a certain number of water molecules and exist in the form of hydrated ions. Some ions form such strong bonds with water molecules that their hydrates can exist not only in solution, but also in the solid state. This explains the formation of crystalline hydrates such as CuSO4 5H 2 O, FeSO 4 7H 2 O, etc., as well as aqua complexes: CI 3 , Br 4 , etc.

II. Hydration of oxides

III. Hydration of organic compounds containing multiple bonds

Hydrolysis reactions

I. Hydrolysis of salts


Reversible hydrolysis:


a) according to the salt cation


Fe 3+ + H 2 O \u003d FeOH 2+ + H +; (acidic environment. pH

b) by salt anion


CO 3 2- + H 2 O \u003d HCO 3 - + OH -; (alkaline environment. pH > 7)


c) by the cation and by the anion of the salt


NH 4 + + CH 3 COO - + H 2 O \u003d NH 4 OH + CH 3 COOH (environment close to neutral)


Irreversible hydrolysis:


Al 2 S 3 + 6H 2 O \u003d 2Al (OH) 3 ↓ + 3H 2 S


II. Hydrolysis of metal carbides


Al 4 C 3 + 12H 2 O \u003d 4Al (OH) 3 ↓ + 3CH 4 netane


CaC 2 + 2H 2 O \u003d Ca (OH) 2 + C 2 H 2 acetylene


III. Hydrolysis of silicides, nitrides, phosphides


Mg 2 Si + 4H 2 O \u003d 2Mg (OH) 2 ↓ + SiH 4 silane


Ca 3 N 2 + 6H 2 O \u003d ZCa (OH) 2 + 2NH 3 ammonia


Cu 3 P 2 + 6H 2 O \u003d ZCu (OH) 2 + 2PH 3 phosphine


IV. Hydrolysis of halogens


Cl 2 + H 2 O \u003d HCl + HClO


Br 2 + H 2 O \u003d HBr + HBrO


V. Hydrolysis of organic compounds


Classes of organic substances

Hydrolysis products (organic)

Halogenalkanes (alkyl halides)

Aryl halides

Dihaloalkanes

Aldehydes or ketones

Metal alcoholates

Carboxylic acid halides

carboxylic acids

Anhydrides of carboxylic acids

carboxylic acids

Esters of carboxylic acids

Carboxylic acids and alcohols

Glycerin and higher carboxylic acids

Di- and polysaccharides

Monosaccharides

Peptides and proteins

α-Amino acids

Nucleic acids

§3. Reaction equation and how to write it

Interaction hydrogen With oxygen, as Sir Henry Cavendish established, leads to the formation of water. Let's use this simple example to learn how to write equations of chemical reactions.
What comes from hydrogen and oxygen, we already know:

H 2 + O 2 → H 2 O

Now we take into account that the atoms of chemical elements in chemical reactions do not disappear and do not appear from nothing, do not turn into each other, but combine in new combinations to form new molecules. This means that in the equation of the chemical reaction of atoms of each type there must be the same number before reactions ( left from the equal sign) and after the end of the reaction ( on right from the equal sign), like this:

2H 2 + O 2 \u003d 2H 2 O

That's what it is reaction equation - conditional record of an ongoing chemical reaction using formulas of substances and coefficients.

This means that in the above reaction two moles hydrogen should react with by one mole oxygen, and the result will be two moles water.

Interaction hydrogen With oxygen- not a simple process at all. It leads to a change in the oxidation states of these elements. To select coefficients in such equations, one usually uses the method " electronic balance".

When water is formed from hydrogen and oxygen, this means that hydrogen changed its oxidation state from 0 before +I, a oxygen- from 0 before −II. At the same time, several (n) electrons:

Hydrogen donating electrons serves here reducing agent, and oxygen accepting electrons - oxidizing agent.

Oxidizing and reducing agents


Now let's see how the processes of giving and receiving electrons look like separately. Hydrogen, having met with the "robber" - oxygen, loses all its property - two electrons, and its oxidation state becomes equal to +I:

H 2 0 − 2 e− = 2Н + I

Happened oxidation half-reaction equation hydrogen.

And the bandit oxygen About 2, having taken the last electrons from the unfortunate hydrogen, is very pleased with his new oxidation state -II:

O 2 + 4 e− = 2O − II

it reduction half-reaction equation oxygen.

It remains to add that both the "bandit" and his "victim" have lost their chemical identity and from simple substances - gases with diatomic molecules H 2 and About 2 turned into components of a new chemical substance - water H 2 O.

Further, we will argue as follows: how many electrons the reductant gave to the oxidizing bandit, that is how much he received. The number of electrons donated by the reducing agent must be equal to the number of electrons accepted by the oxidizing agent..

So you need equalize the number of electrons in the first and second half-reactions. In chemistry, the following conditional form of writing the equations of half-reactions is accepted:

2 H 2 0 − 2 e− = 2Н + I

1 O 2 0 + 4 e− = 2O − II

Here, the numbers 2 and 1 to the left of the curly bracket are factors that will help ensure that the number of given and received electrons is equal. We take into account that in the equations of half-reactions 2 electrons are given away, and 4 are accepted. To equalize the number of received and given electrons, the least common multiple and additional factors are found. In our case, the least common multiple is 4. Additional factors will be 2 for hydrogen (4: 2 = 2), and for oxygen - 1 (4: 4 = 1)
The resulting multipliers will serve as the coefficients of the future reaction equation:

2H 2 0 + O 2 0 \u003d 2H 2 + I O -II

Hydrogen oxidized not only when meeting oxygen. Approximately the same effect on hydrogen and fluorine F2, halogen and the famous "robber", and seemingly harmless nitrogen N 2:

H 2 0 + F 2 0 = 2H + I F −I

3H 2 0 + N 2 0 \u003d 2N -III H 3 + I

This results in hydrogen fluoride HF or ammonia NH3.

In both compounds, the oxidation state hydrogen becomes equal +I, because he gets partners in the molecule "greedy" for someone else's electronic good, with high electronegativity - fluorine F and nitrogen N. At nitrogen the value of electronegativity is considered equal to three conventional units, and y fluorine in general, the highest electronegativity among all chemical elements is four units. So it's no wonder they leave the poor hydrogen atom without any electronic environment.

But hydrogen maybe restore- accept electrons. This happens if alkali metals or calcium, in which the electronegativity is less than that of hydrogen, participate in the reaction with it.

– (old name hydrogen peroxide), a compound of hydrogen and oxygen H 2 O 2 containing a record amount of oxygen 94% by mass. In H molecules 2 O 2 contains peroxide groups OO ( cm. PEROXIDES), which largely determine the properties of this compound.For the first time, hydrogen peroxide was obtained in 1818 by the French chemist Louis Jacques Tenard (1777 1857), acting on barium peroxide with highly chilled hydrochloric acid: BaO 2 + 2HCl ® BaCl 2 + H 2 O 2 . Barium peroxide, in turn, was obtained by burning metallic barium. To isolate from a solution of H 2 O 2 Tenard removed the resulting barium chloride from it: BaCl 2 + Ag 2 SO 4 ® 2AgCl + BaSO 4 . In order not to use expensive silver salt in the future to obtain H 2 O 2 used sulfuric acid: BaO 2 + H 2 SO 4 ® BaSO 4 + H 2 O 2 , because the barium sulfate remains in the sediment. Sometimes another method was used: carbon dioxide was passed into the suspension of BaO 2 in water: BaO 2 + H 2 O + CO 2 ® BaCO 3 + H 2 O 2 , since barium carbonate is also insoluble. This method was proposed by the French chemist Antoine Jerome Balard (18021876), who became famous for the discovery of the new chemical element bromine (1826). More exotic methods were also used, for example, the action of an electric discharge on a mixture of 97% oxygen and 3% hydrogen at liquid air temperature (about 190 ° C), so an 87% solution of H 2 O 2 . Concentrated H 2 O 2 by careful evaporation of very pure solutions in a water bath at a temperature not exceeding 7075 ° C; so you can get about 50% solution. It is impossible to heat up more strongly H 2 O 2 , so the distillation of water was carried out under reduced pressure, using a strong difference in vapor pressure (and, therefore, in boiling point) N 2 O and H 2 O 2 . So, at a pressure of 15 mm Hg. first, mostly water is distilled off, and at 28 mm Hg. and a temperature of 69.7 ° C, pure hydrogen peroxide is distilled off. Another method of concentration is freezing, since when weak solutions freeze, ice contains almost no H 2 O 2 . Finally, it can be dehydrated by absorbing water vapor with sulfuric acid in the cold under a glass bell.

Many researchers of the 19th century, who received pure hydrogen peroxide, noted the danger of this compound. So when they tried to separate

 2 O 2 from water by extraction from dilute solutions with diethyl ether, followed by distillation of the volatile ether, the resulting substance sometimes exploded for no apparent reason. In one of these experiments, the German chemist Yu.V. Brühl obtained anhydrous H 2 O 2 , which had the smell of ozone and exploded at the touch of an unmelted glass rod. Despite the small amount of H 2 O 2 (only 12 ml) the explosion was so strong that it punched a round hole in the table board, destroyed the contents of his drawer, as well as flasks and instruments standing on the table and nearby.physical properties. Pure hydrogen peroxide is very different from the familiar 3% solution of H 2 O 2 , which is in the home first aid kit. First of all, it is almost one and a half times heavier than water (density at 20 ° C is 1.45 g/cm 3). H 2 O 2 freezes at a temperature slightly lower than the freezing point of water at minus 0.41 ° C, but if a pure liquid is quickly cooled, it usually does not freeze, but supercools, turning into a transparent glassy mass. Solutions H 2 O 2 freeze at a much lower temperature: a 30% solution at minus 30 ° C, and a 60% solution at minus 53 ° C. Boils H 2 O 2 at a temperature higher than ordinary water, at 150.2 ° C. Wet glass H 2 O 2 worse than water, and this leads to an interesting phenomenon in the slow distillation of aqueous solutions: while water is distilled from the solution, it, as usual, enters from the refrigerator into the receiver in the form of drops; when does it start to overtake 2 O 2 , the liquid comes out of the refrigerator in the form of a continuous thin stream. On the skin, pure hydrogen peroxide and its concentrated solutions leave white spots and cause a sensation of burning pain due to a severe chemical burn.

In an article on the production of hydrogen peroxide, Tenar did not very well compare this substance with syrup, perhaps he meant that pure H

 2 O 2 , like sugar syrup, strongly refracts light. Indeed, the refractive index of anhydrous H 2 O 2 (1.41) is much larger than that of water (1.33). However, either as a result of a misinterpretation, or because of a poor translation from French, almost all textbooks still write that pure hydrogen peroxide is a “thick syrupy liquid”, and even explain this theoretically by the formation of hydrogen bonds. But water also forms hydrogen bonds. In fact, the viscosity of N 2 O 2 the same as that of slightly chilled (up to about 13 ° C) water, but it cannot be said that cool water is thick, like syrup.decomposition reaction. Pure hydrogen peroxide is a very dangerous substance, since under certain conditions its explosive decomposition is possible: H 2 O 2 ® H 2 O + 1/2 O 2 with a release of 98 kJ per mol N 2 O 2 (34 g). This is a very large energy: it is more than that which is released during the formation of 1 mole of HCl in the explosion of a mixture of hydrogen and chlorine; it is enough to completely evaporate 2.5 times more water than is formed in this reaction. Dangerous and concentrated aqueous solutions H 2 O 2 , in their presence, many organic compounds easily ignite spontaneously, and upon impact, such mixtures can explode. To store concentrated solutions, vessels made of extra pure aluminum or waxed glass vessels are used.

More often you have to meet with a less concentrated 30% solution of H

 2 O 2 , which is called perhydrol, but such a solution is also dangerous: it causes burns on the skin (during its action, the skin immediately turns white due to the discoloration of dyes), explosive boiling is possible if impurities enter. Decomposition H 2 O 2 and its solutions, including explosive ones, cause many substances, for example, heavy metal ions, which in this case play the role of a catalyst, and even dust particles. 2 O 2 are explained by the strong exothermicity of the reaction, the chain nature of the process, and a significant decrease in the activation energy of the decomposition of H 2 O 2 in the presence of various substances, as can be judged from the following data:The enzyme catalase is found in the blood; it is thanks to her that the pharmaceutical “hydrogen peroxide” “boils” from the release of oxygen when it is used to disinfect a cut finger. The decomposition reaction of a concentrated solution of H 2 O 2 under the influence of catalase, not only a person uses; it is this reaction that helps the bombardier beetle fight enemies by releasing a hot jet at them ( cm . EXPLOSIVES). Another enzyme, peroxidase, acts differently: it does not decompose H 2 O 2 , but in its presence, other substances are oxidized by hydrogen peroxide.

Enzymes that affect the reactions of hydrogen peroxide play an important role in the life of the cell. Energy is supplied to the body by oxidation reactions with the participation of oxygen coming from the lungs. In these reactions, H is formed intermediately.

 2 O 2 , which is harmful to the cell, as it causes irreversible damage to various biomolecules. Catalase and peroxidase work together to convert H 2 O 2 into water and oxygen.

Decomposition reaction H

 2 O 2 often proceeds by a radical chain mechanism ( cm. CHAIN ​​REACTIONS), while the role of the catalyst is to initiate free radicals. So, in a mixture of aqueous solutions of H 2 O 2 and Fe 2+ (the so-called Fenton's reagent) there is an electron transfer reaction from the Fe ion 2+ per H 2 O 2 molecule with the formation of the Fe ion 3+ and a very unstable radical anion . – , which immediately decays into the OH anion– and the free hydroxyl radical OH. ( cm. FREE RADICALS). Radical OH. very active. If there are organic compounds in the system, then their various reactions with hydroxyl radicals are possible. So, aromatic compounds and hydroxy acids are oxidized (benzene, for example, turns into phenol), unsaturated compounds can add hydroxyl groups to the double bond: CH 2 \u003d CHCH 2 OH + 2OH. ® HOCH 2 CH(OH)CH 2 OH, but can enter into a polymerization reaction. In the absence of suitable reagents, OH. reacts with H 2 O 2 with the formation of a less active radical HO 2 . , which is capable of reducing Fe ions 2+ , which closes the catalytic cycle: H 2 O 2 + Fe 2+ ® Fe 3+ + OH . + OH OH . + H 2 O 2 ® H 2 O + HO 2 .

HO2 . + Fe3+

® Fe 2+ + O 2 + H + ® H 2 O. Under certain conditions, a chain decomposition of H 2 O 2 , whose simplified mechanism can be represented by the diagram. + H 2 O 2 ® H 2 O + HO 2 . 2 . + H 2 O 2® H 2 O + O 2 + OH . etc.

Decomposition reactions H

 2 O 2 go in the presence of various metals of variable valency. Bound into complex compounds, they often greatly enhance their activity. For example, copper ions are less active than iron ions, but bound to ammonia complexes 2+ , they cause rapid decomposition of H 2 O 2 . Mn ions have a similar effect 2+ associated in complexes with some organic compounds. In the presence of these ions, it was possible to measure the length of the reaction chain. To do this, the reaction rate was first measured by the rate of oxygen evolution from the solution. Then the solution was introduced in a very low concentration (about 10 5 mol / l) inhibitor - a substance that effectively reacts with free radicals and thus terminates the chain. The release of oxygen immediately stopped, but after about 10 minutes, when all the inhibitor was used up, it resumed again at the same rate. Knowing the reaction rate and chain termination rate, it is easy to calculate the chain length, which turned out to be 10 3 links. The long chain length determines the high efficiency of the decomposition of H 2 O 2 in the presence of the most efficient catalysts that generate free radicals at a high rate. At the specified chain length, the decomposition rate H 2 O 2 actually increased by a thousand times.

Sometimes noticeable decomposition of H

 2 O 2 cause even traces of impurities, which are almost not detected analytically. So, one of the most effective catalysts turned out to be a metal osmium sol: its strong catalytic effect was observed even at a dilution of 1:10 9 , i.e. 1 g Os per 1000 tons of water. Active catalysts are colloidal solutions of palladium, platinum, iridium, gold, silver, as well as solid oxides of some metals MnO 2 , Co 2 O 3 , PbO 2 etc., which themselves do not change. Decomposition can go very rapidly. So, if a small pinch of MnO 2 throw into a test tube with a 30% solution of H 2 O 2 , a column of vapor escapes from the test tube with splashes of liquid. With more concentrated solutions, an explosion occurs. The decomposition proceeds more smoothly on the platinum surface. In this case, the state of the surface has a strong influence on the reaction rate. The German chemist Walter Spring conducted at the end of the 19th century. such an experience. In a thoroughly cleaned and polished platinum cup, the decomposition of a 38% solution of H 2 O 2 did not go even when heated to 60 ° C. If, however, a barely noticeable scratch is made with a needle at the bottom of the cup, then the already cold (at 12 ° C) solution begins to release oxygen bubbles at the site of the scratch, and when heated, decomposition along this place noticeably increases. If spongy platinum, which has a very large surface, is introduced into such a solution, explosive decomposition is possible.

Rapid decomposition H

 2 O 2 can be used for a spectacular lecture experience if a surfactant (soap, shampoo) is added to the solution before the catalyst is added. The released oxygen creates a rich white foam, which has been called "elephant toothpaste".

Some catalysts initiate non-chain decomposition of H

 2 O 2, for example: H 2 O 2 + 2I + 2H + ® 2H 2 O + I 2 ® 2I + 2H + + O 2. A non-chain reaction also occurs in the case of oxidation of Fe ions 2+ in acid solutions: 2FeSO 4 + H 2 O 2 + H 2 SO 4 ® Fe 2 (SO 4) 3 + 2H 2 O. Since aqueous solutions almost always contain traces of various catalysts (decomposition can also be catalyzed by metal ions contained in glass), H 2 O 2 , even diluted, during their long-term storage, inhibitors and stabilizers that bind metal ions are added. In this case, the solutions are slightly acidified, since under the action of pure water on the glass, a weakly alkaline solution is obtained, which contributes to the decomposition of H 2 O 2 . All these features of the decomposition of H 2 O 2 allow the conflict to be resolved. To get pure H 2 O 2 it is necessary to carry out distillation under reduced pressure, since the substance decomposes when heated above 70 ° C and even, although very slowly, at room temperature (as stated in the Chemical Encyclopedia, at a rate of 0.5% per year). In this case, how was the boiling point at atmospheric pressure, which appears in the same encyclopedia, equal to 150.2 ° C, obtained? Usually, in such cases, a physicochemical regularity is used: the logarithm of the vapor pressure of a liquid depends linearly on the reciprocal temperature (on the Kelvin scale), so if you accurately measure the vapor pressure H 2 O 2 at several (low) temperatures, it is easy to calculate at what temperature this pressure will reach 760 mm Hg. And this is the boiling point under normal conditions.

Theoretically, OH radicals

. can also be formed in the absence of initiators, as a result of breaking the weaker OO bond, but this requires a rather high temperature. Despite the relatively low breaking energy of this bond in the H 2 O 2 (it is equal to 214 kJ / mol, which is 2.3 times less than for the HOH bond in a water molecule), the OO bond is still strong enough for hydrogen peroxide to be absolutely stable at room temperature. And even at the boiling point (150°C), it must decompose very slowly. The calculation shows that atAt this temperature, decomposition by 0.5% should also occur quite slowly, even if the chain length is 1000 links. The discrepancy between calculations and experimental data is explained by catalytic decomposition caused by both the smallest impurities in the liquid and the walls of the reaction vessel. Therefore, the decomposition activation energy H measured by many authors 2 O 2 always significantly less than 214 kJ/mol even "in the absence of a catalyst". In fact, there is always a decomposition catalyst both in the form of insignificant impurities in the solution and in the form of vessel walls, which is why heating anhydrous H 2 O 2 to boiling at atmospheric pressure repeatedly caused explosions.

Under certain conditions, the decomposition of H

 2 O 2 occurs very unusually, for example, if you heat a solution of H acidified with sulfuric acid 2 O 2 in the presence of potassium iodate KIO 3 , then at certain concentrations of reagents, an oscillatory reaction is observed, while the release of oxygen periodically stops, and then resumes with a period of 40 to 800 seconds.Chemical properties H 2 O 2 . Hydrogen peroxide is an acid, but a very weak one. Dissociation constant H 2 O 2 H + + HO 2 at 25 ° C is 2.4 10 12 , which is 5 orders of magnitude less than for H 2 S. Medium salts H 2 O 2 alkali and alkaline earth metals are commonly referred to as peroxides ( cm. PEROXIDES). When dissolved in water, they are almost completely hydrolyzed: Na 2 O 2 + 2H 2 O ® 2NaOH + H 2 O 2 . Hydrolysis is promoted by acidification of solutions. Like an acid 2 O 2 also forms acidic salts, for example, Ba (HO 2) 2 , NaHO 2 and others. Acid salts are less susceptible to hydrolysis, but easily decompose when heated to release oxygen: 2NaHO 2 ® 2NaOH + O 2 . The liberated alkali, as in the case of H 2 O 2 promotes decomposition.

Solutions H

 2 O 2 , especially concentrated, have a strong oxidizing effect. So, under the action of a 65% solution of H 2 O 2 on paper, sawdust and other combustible substances, they ignite. Less concentrated solutions decolorize many organic compounds, such as indigo. Formaldehyde is oxidized unusually: H 2 O 2 is reduced not to water (as usual), but to free hydrogen: 2HCHO + H 2 O 2 ® 2HCOOH + H 2 . If we take a 30% solution of H 2 O 2 and a 40% solution of HCHO, then after a slight heating, a violent reaction begins, the liquid boils and foams. Oxidative effect of dilute solutions of H 2 O 2 is most pronounced in an acidic environment, for example, H 2 O 2 + H 2 C 2 O 4 ® 2H 2 O + 2CO 2 , but oxidation is also possible in an alkaline environment:Na + H 2 O 2 + NaOH® Na 2 ; 2K 3 + 3H 2 O 2® 2KCrO 4 + 2KOH + 8H 2 O. Oxidation of black lead sulfide to white sulfate PbS+ 4H 2 O 2 ® PbSO 4 + 4H 2 O can be used to restore tarnished white lead in old paintings. Under the action of light, hydrochloric acid is also oxidized: H 2 O 2 + 2HCl ® 2H 2 O + Cl 2. Addition of H 2 O 2 to acids greatly increases their effect on metals. So, in a mixture of H 2O2 and dilute H 2 SO 4 copper, silver and mercury dissolve; iodine in an acidic environment is oxidized to iodic acid HIO 3 , sulfur dioxide to sulfuric acid, etc.

Unusually, the potassium-sodium salt of tartaric acid (Rochelle salt) is oxidized in the presence of cobalt chloride as a catalyst. During the reaction KOOC(CHOH)

 2 COONa + 5H 2 O 2 ® KHCO 3 + NaHCO 3 + 6H 2 O + 2CO 2 pink CoCl 2 changes color to green due to the formation of a complex compound with tartrate tartaric acid anion. As the reaction proceeds and the tartrate is oxidized, the complex is destroyed and the catalyst turns pink again. If copper sulphate is used instead of cobalt chloride as a catalyst, then the intermediate compound, depending on the ratio of the initial reagents, will be colored orange or green. After the end of the reaction, the blue color of copper sulphate is restored.

Hydrogen peroxide reacts completely differently in the presence of strong oxidizing agents, as well as substances that readily give off oxygen. In such cases, N

 2 O 2 can also act as a reducing agent with the simultaneous evolution of oxygen (the so-called reductive decomposition of H 2 O 2 ), for example: 2KMnO 4 + 5H 2 O 2 + 3H 2 SO 4® K 2 SO 4 + 2MnSO 4 + 5O 2 + 8H 2 O;

Ag 2 O + H 2 O 2

® 2Ag + H 2 O + O 2; O 3 + H 2 O 2 ® H 2 O + 2O 2; ® NaCl + H 2 O + O 2 . The latter reaction is interesting in that it produces excited oxygen molecules that emit orange fluorescence ( cm. CHLORINE ACTIVE). Similarly, metallic gold is isolated from solutions of gold salts, metallic mercury is obtained from mercury oxide, etc. Such an unusual property 2 O 2 allows, for example, to carry out the oxidation of potassium hexacyanoferrate(II), and then, by changing the conditions, to restore the reaction product to the starting compound using the same reagent. The first reaction takes place in an acidic environment, the second - in an alkaline one:2K 4 + H 2 O 2 + H 2 SO 4® 2K 3 + K 2 SO 4 + 2H 2 O;

2K 3 + H 2 O 2 + 2KOH

® 2K 4 + 2H 2 O + O 2 .("Dual character" H 2 O 2 allowed one chemistry teacher to compare hydrogen peroxide with the hero of the story of the famous English writer Stevenson The Strange Case of Dr. Jekyll and Mr. Hyde, under the influence of the composition invented by him, he could drastically change his character, turning from a respectable gentleman into a bloodthirsty maniac.)Obtaining H 2 O 2. H 2 O 2 molecules always obtained in small quantities during combustion and oxidation of various compounds. When burning H 2 O 2 is formed either when hydrogen atoms are abstracted from the starting compounds by intermediate hydroperoxide radicals, for example: HO 2 . + CH 4 ® H 2 O 2 + CH 3 . , or as a result of recombination of active free radicals: 2OH. ® H 2 O 2, N. + NO 2 . ® H 2 O 2 . For example, if an oxy-hydrogen flame is directed to a piece of ice, then the melted water will contain appreciable amounts of H 2 O 2 , formed as a result of the recombination of free radicals (in the flame of the H molecule 2 O 2 disintegrate immediately). A similar result is obtained with the combustion of other gases. Education H 2 O 2 can also occur at low temperatures as a result of various redox processes.

In industry, hydrogen peroxide is no longer obtained by the Tenard method from barium peroxide, but more modern methods are used. One of them is the electrolysis of sulfuric acid solutions. At the same time, sulfate ions are oxidized at the anode to oversulfate ions: 2SO

 4 2 2e ® S 2 O 8 2 . Persulfuric acid is then hydrolyzed: H 2 S 2 O 8 + 2H 2 O ® H 2 O 2 + 2H 2 SO 4. At the cathode, as usual, hydrogen is released, so that the overall reaction is described by the equation 2H 2 O ® H 2 O 2 + H 2 . But the main modern method (over 80% of world production) is the oxidation of some organic compounds, for example, ethylanthrahydroquinone, with atmospheric oxygen in an organic solvent, while H 2 O 2 and the corresponding anthraquinone, which is then again reduced with hydrogen on a catalyst to anthrahydroquinone. Hydrogen peroxide is removed from the mixture with water and concentrated by distillation. A similar reaction also occurs when isopropyl alcohol is used (it proceeds with the intermediate formation of hydroperoxide): (CH 3) 2 CHOH + O 2 ® (CH 3) 2 C (UN) OH ® (CH 3) 2 CO + H 2 O 2 . If necessary, the resulting acetone can also be reduced to isopropyl alcohol.The use of H 2 O 2. Hydrogen peroxide is widely used, and its world production amounts to hundreds of thousands of tons per year. It is used to obtain inorganic peroxides, as an oxidizer for rocket fuels, in organic synthesis, for bleaching oils, fats, fabrics, paper, for cleaning semiconductor materials, for extracting valuable metals from ores (for example, uranium by converting its insoluble form into a soluble one), for waste water treatment. In medicine, solutions of H 2 O 2 used for rinsing and lubrication in inflammatory diseases of the mucous membranes (stomatitis, tonsillitis), for the treatment of purulent wounds. In contact lens cases, a very small amount of platinum catalyst is sometimes placed in the lid. Lenses for their disinfection are poured in a pencil case with a 3% solution of H 2 O 2 , but since this solution is harmful to the eyes, the pencil case is turned over after a while. In this case, the catalyst in the lid quickly decomposes H 2 O 2 for clean water and oxygen.

Once it was fashionable to bleach hair with “peroxide”, now there are safer formulations for hair coloring.

In the presence of some salts, hydrogen peroxide forms a kind of solid "concentrate", which is more convenient to transport and use. So, if H

 2 O 2 in the presence, large transparent crystals of sodium peroxoborate Na are gradually formed 2 [(BO 2) 2 (OH) 4 ]. This substance is widely used for bleaching fabrics and as a component of detergents. H molecules 2 O 2 , like water molecules, are able to penetrate into the crystal structure of salts, forming a kind of crystalline hydrates peroxohydrates, for example, K 2 CO 3 3H 2 O 2, Na 2 CO 3 1.5H 2 O; the latter compound is commonly known as "persol".

The so-called "hydroperite" CO(NH

 2) 2 H 2 O 2 is a clathrate inclusion compound of H molecules 2 O 2 into the voids of the urea crystal lattice.

In analytical chemistry, certain metals can be determined using hydrogen peroxide. For example, if hydrogen peroxide is added to a solution of a titanium (IV) salt of titanyl sulfate, the solution becomes bright orange due to the formation of pertitanic acid:

 TiOSO 4 + H 2 SO 4 + H 2 O 2 ® H 2 + H 2 O.Colorless molybdate ion MoO 4 2 oxidized H 2 O 2 into an intensely orange-colored peroxide anion. Acidified potassium dichromate solution in the presence of H 2 O 2 forms perchromic acid: K2 Cr 2 O 7 + H 2 SO 4 + 5H 2 O 2® H 2 Cr 2 O 12 + K 2 SO 4 + 5H 2O, which decomposes fairly quickly: H 2 Cr 2 O 12 + 3H 2 SO 4 ® Cr 2 (SO 4) 3 + 4H 2 O + 4O 2. If you add these two equations, you get the reaction for the reduction of potassium dichromate with hydrogen peroxide:K 2 Cr 2 O 7 + 4H 2 SO 4 + 5H 2 O 2® Cr 2 (SO 4) 3 + K 2 SO 4 + 9H 2 O + 4O 2.Perchromic acid can be extracted from an aqueous solution with ether (it is much more stable in an ether solution than in water). The ethereal layer is colored in an intense blue color.

Ilya Leenson

LITERATURE Dolgoplosk B.A., Tinyakova E.I. Free radical generation and reactions. M., Chemistry, 1982
Chemistry and technology of hydrogen peroxide. L., Chemistry, 1984