– (old name hydrogen peroxide), a compound of hydrogen and oxygen H 2 O 2 , containing a record amount of oxygen 94% by weight. In molecules H 2 O 2 contains peroxide groups ОО ( cm. PEROXIDES), which largely determine the properties of this compound.Hydrogen peroxide was first obtained in 1818 by the French chemist Louis Jacques Thénard (1777 1857) by treating barium peroxide with highly cooled hydrochloric acid: BaO 2 + 2HCl ® BaCl 2 + H 2 O 2 . Barium peroxide, in turn, was obtained by burning barium metal. To isolate H from solution 2 O 2 Tenar 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 , since barium sulfate remains in the sediment. Sometimes another method was used: carbon dioxide was passed into the BaO suspension 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), famous for the discovery of a 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 was obtained 2 O 2 . Concentrated H 2 O 2 by carefully evaporating very pure solutions in a water bath at a temperature not exceeding 70-75 ° C; this way you can get approximately a 50% solution. You can’t heat it up any more; decomposition of H will occur. 2 O 2 , therefore the distillation of water was carried out under reduced pressure, taking advantage of the strong difference in vapor pressure (and therefore boiling point) H 2 O and H 2 O 2 . So, at a pressure of 15 mm Hg. First, mainly 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 is possible to dehydrate by absorbing water vapor with sulfuric acid in the cold under a glass bell.

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

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. Bruhl obtained anhydrous H 2 O 2 , which smelled like ozone and exploded when touched by an unfused glass rod. Despite small amounts of H 2 O 2 (total 12 ml) the explosion was so powerful that it punched a round hole in the board of the table, destroyed the contents of its drawer, as well as the bottles 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 medicine cabinet. First of all, it is almost one and a half times heavier than water (density at 20°C is 1.45 g/cm 3). H2O2 freezes at a temperature slightly lower than the freezing point of water at minus 0.41 ° C, but if you quickly cool a pure liquid, it usually does not freeze, but is supercooled, 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. Wets glass H 2 O 2 worse than water, and this leads to an interesting phenomenon during the slow distillation of aqueous solutions: while water is distilled from the solution, it, as usual, flows from the refrigerator to the receiver in the form of drops; when does it start to distill 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 burning sensation due to a severe chemical burn.

In an article devoted to the production of hydrogen peroxide, Tenard did not very successfully 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 greater than that of water (1.33). However, either as a result of misinterpretation, or because of 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 cooled (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 releasing 98 kJ per mol H 2 O 2 (34 g). This is a very large energy: it is greater than that released when 1 mole of HCl is formed during 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. Concentrated aqueous solutions of H are also dangerous 2 O 2 , in their presence many organic compounds easily ignite spontaneously, and upon impact such mixtures can explode. To store concentrated solutions, use vessels made of especially pure aluminum or waxed glass vessels.

More often you encounter 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 (when it acts, the skin immediately turns white due to the discoloration of coloring substances), and if impurities enter, explosive boiling is possible. Decomposition H 2 O 2 and its solutions, including explosive ones, are caused by 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 H decomposition 2 O 2 in the presence of various substances, as can be judged by the following data:The enzyme catalase is found in the blood; It is thanks to it that pharmaceutical “hydrogen peroxide” “boils” from the release of oxygen when it is used to disinfect a cut finger. Decomposition reaction of a concentrated solution of H 2 O 2 not only humans use catalase; It is this reaction that helps the bombardier beetle fight enemies by releasing a hot stream at them ( cm . EXPLOSIVES). Another enzyme, peroxidase, acts differently: it does not decompose H 2 O 2 , but in its presence, oxidation of other substances with hydrogen peroxide occurs.

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

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

H decomposition reaction

2 O 2 often proceeds via a radical chain mechanism ( cm. CHAIN ​​REACTIONS), while the role of the catalyst is to initiate free radicals. Thus, in a mixture of aqueous solutions of H 2 O 2 and Fe 2+ (the so-called Fenton reagent) an electron transfer reaction occurs from the Fe ion 2+ per H 2 O 2 molecule with the formation of Fe ion 3+ and a very unstable radical anion . – , which immediately decays into the OH anion– and free hydroxyl radical OH. ( cm. FREE RADICALS). Radical HE. very active. If there are organic compounds in the system, then various reactions with hydroxyl radicals are possible. Thus, aromatic compounds and hydroxy acids are oxidized (benzene, for example, turns into phenol), unsaturated compounds can attach hydroxyl groups to the double bond: CH 2 =CHCH 2 OH + 2OH. ® NOCH 2 CH(OH)CH 2 OH, and 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 .

HO 2 . + Fe 3+

® Fe 2+ + O 2 + H + ® H 2 O. Under certain conditions, chain decomposition of H is possible 2 O 2 , a simplified mechanism of which can be represented by the diagram. + H 2 O 2 ® H 2 O + HO 2 . 2 . +H2O2® H 2 O + O 2 + OH . etc.

H decomposition reactions

2 O 2 go in the presence of various metals variable valence. When bound into complex compounds, they often significantly enhance their activity. For example, copper ions are less active than iron ions, but are bound in ammonia complexes 2+ , they cause rapid decomposition of H 2 O 2 . Mn ions have a similar effect 2+ bound in complexes with certain organic compounds. In the presence of these ions, it was possible to measure the length of the reaction chain. To do this, we first measured the reaction rate by the rate of release of oxygen from the solution. Then a very low concentration (about 10 5 mol/l) inhibitor a substance that effectively reacts with free radicals and thus breaks 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 the rate of chain termination, it is easy to calculate the chain length, which turned out to be equal to 10 3 links The large chain length determines the high efficiency of H decomposition 2 O 2 in the presence of the most effective catalysts that generate free radicals at a high rate. For the indicated chain length, the decomposition rate H 2 O 2 actually increases a thousand times.

Sometimes noticeable decomposition of H

2 O 2 cause even traces of impurities that are almost undetectable analytically. Thus, one of the most effective catalysts turned out to be a sol of metal osmium: 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 proceed very rapidly. So, if a small pinch of MnO 2 drop into a test tube with a 30% solution of H 2 O 2 , a column of steam bursts out of the test tube with a splash of liquid. With more concentrated solutions an explosion occurs. Decomposition occurs more quietly on the surface of platinum. In this case, the reaction rate is strongly influenced by the state of the surface. 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 reaction of a 38% solution of H 2 O 2 did not go even when heated to 60° C. If you make a barely noticeable scratch on the bottom of the cup with a needle, then the already cold (at 12° C) solution begins to release oxygen bubbles at the scratch site, and when heated, decomposition along this place noticeably intensifies. If spongy platinum, which has a very large surface area, is introduced into such a solution, then explosive decomposition is possible.

Rapid decomposition of H

2 O 2 can be used for an effective lecture experiment if a surfactant (soap, shampoo) is added to the solution before adding the catalyst. The oxygen released 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 acidic 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 (metal ions contained in glass can also catalyze decomposition), solutions of H 2 O 2 , even diluted, during long-term storage, inhibitors and stabilizers that bind metal ions are added. In this case, the solutions are slightly acidified, since the action of pure water on glass produces a weakly alkaline solution, which promotes the decomposition of H 2 O 2 . All these features of the decomposition of H 2 O 2 allow the contradiction to be resolved. To obtain 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 of 150.2° C, which appears in the same encyclopedia, obtained? Usually in such cases a physicochemical law is used: the logarithm of the vapor pressure of a liquid linearly depends on the inverse 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 form in the absence of initiators, as a result of the rupture of a weaker OO bond, but this requires a fairly high temperature. Despite the relatively low energy of breaking this bond in the H molecule 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 boiling point (150°C) it should decompose very slowly. The calculation shows that whenAt this temperature, decomposition of 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 the smallest impurities in the liquid and the walls of the reaction vessel. Therefore, the activation energy of H decomposition measured by many authors 2 O 2 always significantly less than 214 kJ/mol even “in the absence of a catalyst.” In fact, a decomposition catalyst is always present, both in the form of insignificant impurities in the solution and in the form of the walls of the vessel, which is why heating anhydrous H 2 O 2 to boiling at atmospheric pressure repeatedly caused explosions.

Under some conditions, the decomposition of H

2 O 2 occurs very unusually, for example, if you heat a solution of H 2 O 2 in the presence of potassium iodate KIO 3 , then at certain concentrations of reagents there is oscillatory reaction, while the release of oxygen periodically stops and then resumes with a period of 40 to 800 seconds.Chemical properties of 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 equal to 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 usually called 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 acid H 2 O 2 also forms acid salts, for example, Ba(HO 2) 2, NaHO 2 etc. Acid salts are less susceptible to hydrolysis, but easily decompose when heated, releasing oxygen: 2NaHO 2 ® 2NaOH + O 2 . Alkali released, as in the case of H 2 O 2 , promotes decomposition.

Solutions H

2 O 2 , especially concentrated ones, have a strong oxidizing effect. Thus, under the influence of a 65% solution of H 2 O 2 on paper, sawdust and other flammable substances they ignite. Less concentrated solutions decolorize many organic compounds, such as indigo. The oxidation of formaldehyde occurs unusually: H 2 O 2 is reduced not to water (as usual), but to free hydrogen: 2HCHO + H 2 O 2 ® 2НСООН + Н 2 . If you take a 30% solution of H 2 O 2 and a 40% solution of HCHO, then after 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 discolored lead white on old paintings. Under the influence of light, oxidation occurs and of hydrochloric acid: H 2 O 2 + 2HCl ® 2H 2 O + Cl 2 . Adding H 2 O 2 to acids greatly increases their effect on metals. Thus, in a mixture of H 2 O 2 and dilute H 2 SO 4 copper, silver and mercury dissolve; iodine in an acidic environment is oxidized to periodic acid HIO 3 , sulfur dioxide to sulfuric acid, etc.

Unusually, the oxidation of potassium sodium salt of tartaric acid (Rochelle salt) occurs 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, the tartaric acid anion. As the reaction proceeds and the tartrate is oxidized, the complex is destroyed and the catalyst turns pink again. If copper sulfate is used as a catalyst instead of cobalt chloride, the intermediate compound, depending on the ratio of the starting reagents, will be colored orange or green. After the end of the reaction it is restored Blue colour copper sulfate.

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

2 O 2 can also act as a reducing agent with the simultaneous release 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 last reaction is interesting because it produces excited oxygen molecules that emit orange fluorescence ( cm. CHLORINE ACTIVE). Similarly, metallic gold is released 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, restore the reaction product to the original compound using the same reagent. The first reaction occurs in an acidic environment, the second in an alkaline environment:2K 4 + H 2 O 2 + H 2 SO 4® 2K 3 + K 2 SO 4 + 2H 2 O;

2K3 + H2O2 + 2KOH

® 2K 4 + 2H 2 O + O 2.(“Dual character” N 2 O 2 allowed one chemistry teacher to compare hydrogen peroxide with the hero of the story by the famous English writer Stevenson The Strange Case of Dr Jekyll and Mr Hyde, under the influence of the composition he invented, he could dramatically change his character, turning from a respectable gentleman into a bloodthirsty maniac.)Obtaining H 2 O 2. Molecules H 2 O 2 are always obtained in small quantities during the combustion and oxidation of various compounds. When burning H 2 O 2 is formed either by the abstraction of hydrogen atoms 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. ® Н 2 О 2 , Н . + BUT 2 . ® H 2 O 2 . For example, if an oxygen-hydrogen flame is directed at a piece of ice, then the melted water will contain noticeable 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 when other gases burn. Education N 2 O 2 can also occur at low temperatures as a result of various redox processes.

In industry, hydrogen peroxide has long been no longer produced by the Tenara method from barium peroxide, but is used more modern methods. One of them is electrolysis of sulfuric acid solutions. In this case, at the anode, sulfate ions are oxidized to persulfate ions: 2SO

4 2 2e ® S 2 O 8 2 . The 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 evolution occurs, so 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 reduced again with hydrogen on the catalyst to anthrahydroquinone. Hydrogen peroxide is removed from the mixture with water and concentrated by distillation. A similar reaction occurs when using isopropyl alcohol (it occurs 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.Application of H 2 O 2. Hydrogen peroxide is widely used, and its global production amounts to hundreds of thousands of tons per year. It is used to produce inorganic peroxides, as an oxidizer for rocket fuels, in organic syntheses, for bleaching oils, fats, fabrics, paper, for purifying semiconductor materials, for extracting valuable metals from ores (for example, uranium by converting its insoluble form into a soluble one), for wastewater treatment. In medicine, solutions N 2 O 2 used for rinsing and lubricating in inflammatory diseases of the mucous membranes (stomatitis, sore throat), for the treatment of purulent wounds. In cases for storing contact lenses, very small parts are sometimes placed in the lid. a large number of platinum catalyst. For disinfection, lenses are filled 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 upon a time it was fashionable to bleach hair with “peroxide”; now there are safer hair coloring compounds.

In the presence of certain salts, hydrogen peroxide forms a kind of solid “concentrate”, which is more convenient to transport and use. So, if you add H to a very cooled saturated solution of sodium borate (borax)

2 O 2 in the presence, large transparent crystals of sodium peroxoborate Na 2 [(BO 2) 2 (OH) 4 ]. This substance is widely used to bleach fabrics and as a component of detergents. Molecules H 2 O 2 , like water molecules, are able to penetrate into the crystalline structure of salts, forming something like 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 compound of inclusion of H molecules 2 O 2 into the voids of the urea crystal lattice.

In analytical chemistry, hydrogen peroxide can be used to determine some metals. For example, if hydrogen peroxide is added to a solution of titanium(IV) salt 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 is oxidized by H 2 O 2 into an intensely orange-colored peroxide anion. Acidified solution of potassium dichromate 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 quite quickly: H 2 Cr 2 O 12 + 3H 2 SO 4 ® Cr 2 (SO 4) 3 + 4H 2 O + 4O 2. If we add these two equations, we get the reaction of 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 turns intense blue.

Ilya Leenson

LITERATURE Dolgoplosk B.A., Tinyakova E.I. Generation of free radicals and their reactions. M., Chemistry, 1982
Chemistry and technology of hydrogen peroxide. L., Chemistry, 1984

Water (hydrogen oxide) - binary inorganic compound with the chemical formula H 2 O. A water molecule consists of two hydrogen atoms and one oxygen atom, which are connected to each other covalent bond.

Hydrogen peroxide.

Physical and chemical properties

Physical and Chemical properties water is 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, +1 and -2, respectively; therefore, water does not exhibit pronounced oxidative 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 the H and O atoms, as well as lone electron pairs on the 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 m.p. values. 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 cleavage of the polar O-N connections, and the released proton attaches to the oxygen atom of another molecule. The resulting hydronium ion H 3 O + is essentially a hydrated hydrogen ion H + H 2 O, so the self-ionization equation for water 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, exhibiting neither acidic nor basic properties to a noticeable extent.
However, water has a strong ionizing effect on the electrolytes dissolved in it. Under the influence of water dipoles, polar covalent bonds in the molecules of dissolved substances 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 without taking into account 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 Brønsted-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 only possible with strong reducing agents that are capable of reducing the hydrogen ions contained in water molecules to free hydrogen.

1) Interaction with metals

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

2Na + 2H + 2 O = 2NaOH + H 0 2

Ca + 2H + 2 O = Ca(OH) 2 + H 0 2

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

Mg + 2H + 2 O = Mg(OH) 2 + H 0 2

3Fe + 4H + 2 O = Fe 2 O 4 + 4H 0 2

c) Al and Zn displace H2 from water in the presence of alkalis:

2Al + 6H + 2 O + 2NaOH = 2Na + 3H 0 2

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

C + H + 2 O = CO + H 0 2 (“water gas”)

2P + 6H + 2 O = 2HPO 3 + 5H 0 2

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

Si + H + 2 O + 2NaOH = Na 2 SiO 3 + 2H 0 2

3) Interaction with metal hydrides

NaH + H + 2 O = NaOH + H 0 2

CaH 2 + 2H + 2 O = Ca(OH) 2 + 2H 0 2

4) Interaction with carbon monoxide and methane

CO + H + 2 O = CO 2 + H 0 2

2CH 4 + O 2 + 2H + 2 O = 2CO 2 + 6H 0 2

The reactions are used industrially to produce hydrogen.

II. Reactions in which water plays the role of 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 a reaction with F 2), oxygen is formed with c o. +2.

1) Interaction with fluorine

2F 2 + 2H 2 O -2 = O 0 2 + 4HF

2F 2 + H 2 O -2 = O +2 F 2 + 2HF

2) Interaction with atomic oxygen

H 2 O -2 + O = H 2 O - 2

3) Interaction with chlorine

At high T a reversible reaction occurs

2Cl 2 + 2H 2 O -2 = O 0 2 + 4HCl

III. Reactions of intramolecular oxidation - reduction of water.

Under the influence electric current or high temperature, water can decompose into hydrogen and oxygen:

2H + 2 O -2 = 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. 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. Oxides hydration

III. Hydration of organic compounds containing multiple bonds

Hydrolysis reactions

I. Hydrolysis of salts

Reversible hydrolysis:

a) by salt cation

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

b) according to the salt anion

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

c) by cation and anion of the salt

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

Irreversible hydrolysis:

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

II. Hydrolysis of metal carbides

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

CaC 2 + 2H 2 O = Ca(OH) 2 + C 2 H 2 acetylene

III. Hydrolysis of silicides, nitrides, phosphides

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

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

Cu 3 P 2 + 6H 2 O = 3Сu(OH) 2 + 2РН 3 phosphine

IV. Hydrolysis of halogens

Cl 2 + H 2 O = HCl + HClO

Br 2 + H 2 O = HBr + HBrO

V. Hydrolysis of organic compounds

Classes organic matter	Hydrolysis products (organic)
Haloalkanes (alkyl halides)
Aryl halides
Dihaloalkanes	Aldehydes or ketones
Metal alcoholates
Carboxylic acid halides	Carboxylic acids
Carboxylic acid anhydrides	Carboxylic acids
Complex ethers of carboxylic acids	Carboxylic acids and alcohols
	Glycerol and higher carboxylic acids
Di- and polysaccharides	Monosaccharides
Peptides and proteins	α-Amino acids
Nucleic acids

2. Write down the kinetic equation for the reaction: 2H2 + O2 = 2H2O. 3. How many times will the reaction rate increase if the temperature coefficient is 3 and the temperature is increased by 30 degrees? 4. When the temperature increases by 40 degrees, the reaction rate increases 16 times. Determine the temperature coefficient.

Picture 12 from the presentation “Reaction Speed” for chemistry lessons on the topic “Reactions”

Dimensions: 960 x 720 pixels, format: jpg. To download a picture for free chemistry lesson, right-click on the image and click “Save Image As...”. To display pictures in the lesson, you can also download the entire presentation “Reaction Speed.ppt” with all the pictures in a zip archive for free. The archive size is 15 KB.

Download presentation

Reactions

“Reaction speed” - Factors affecting speed. What did we study? Effect of concentration of reactants (for homogeneous systems) 3rd row. Temperature. What determines the rate of reactions? 2. Write down the kinetic equation for the reaction: 2H2 + O2 = 2H2O. Presence of catalysts or inhibitors. Problem solving. Catalysts and catalysis.

“The law of conservation of mass of substances” - 1673. Law of conservation of mass of substances. Index. The index shows the number of atoms in the formula unit of a substance. Like Boyle, the Russian scientist experimented in sealed retorts. 1789 General high school No. 36 named after Kazybek bi. Robert Boyle. Coefficient. 5n2o. 1748 Chemical formula. Lesson objectives: Educational - experimentally prove the law of conservation of mass of substances.

“Radioactive transformations” - Milestones of history. No is the number of radioactive nuclei at the initial time. t–decay time. The law of radioactive decay. Experience. What is half-life? T-half-life. Rutherford's research. Conclusion from the rules. Atoms of a radioactive substance are subject to spontaneous modifications. Background to radioactivity research.

“Chemical reactions practical work” - PPG. H2 – Gas, colorless, odorless, lighter than air. 4) Black CuO turns red, H2O forms on the walls of the test tube. Test tubes. 2) Pure H2 explodes with a dull bang, H2 with impurities - a barking sound. 3kcns+feci3=3kci+fe(cns)3 exchange. AI+HCI. Cu. Zn+H2SO4 = ZnSO4+H2 Substitution. Alcohol lamp. Signs of chemical reactions were observed.

“Reactions” - Odor appearance. Give initial ideas about chemical reaction. Gas release. Equipment: Solutions - hydrochloric acid and lime water, a piece of marble. Examination homework. Give examples of complex substances? The role of chemistry in human life. Formation of sediment. Release or absorption of heat.

“Theory of electrolytic dissociation” - All simple substances, all oxides and some acids, bases and salts. Svante Arrhenius. Substances in solutions. Substances with ionic and covalent polar bonds. The theory of electrolytic dissociation (ED). II position of the TED. Substances with covalent bonds: Orientation of water dipoles? hydration? ionization? dissociation.

There are 28 presentations in total

The formula for the basis of life - water - is well known. 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 be obtained - H2O2. What is it and how will the resulting substance differ from its “relative” water?

H2O2 - what is this substance?

Let's look at it in more detail. H2O2 is the formula of hydrogen peroxide, Yes, the same one that is used to treat scratches, white. Hydrogen peroxide H2O2 - scientific.

For disinfection, use a three percent peroxide solution. In pure or concentrated form, it causes chemical burns to the skin. A thirty percent peroxide solution is otherwise called perhydrol; Previously, it was used in hairdressers to bleach hair. The skin burned by it also turns white.

Chemical properties of H2O2

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

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

H2O2 can be both an oxidizing agent and a reducing agent in chemical reactions. However, more often 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 producing oxygen

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

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

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

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

The decomposition of hydrogen peroxide can also occur at room temperature if a catalyst is used ( Chemical substance, accelerating the reaction).

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

History of the discovery of peroxide

The first steps towards the discovery of peroxide were taken 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, scientists Tenard and Gay-Lussac conducted an experiment on burning alkali metals with excess oxygen, resulting in sodium peroxide. But hydrogen peroxide was obtained later, only in 1818, when Louis Thénard studied the effect of acids on metals; a low amount of oxygen was necessary for their stable interaction. 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 discovered small frozen drops on the walls of the container with barium peroxide. It became clear that this was H2O2. Then they gave the resulting H2O2 the name “oxidized water.” This was hydrogen peroxide - a colorless, odorless, difficult-to-evaporate liquid that dissolves other substances well. The result of the interaction of H2O2 and H2O2 is a dissociation reaction, 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. Tenar himself, using peroxide, restored a painting by Raphael that had darkened with time.

Hydrogen peroxide in the 20th century

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

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

Where is H2O2 used now?

• in medicine - for treating wounds;
• in the pulp and paper industry the bleaching properties of the substance are used;
• in the textile industry, natural and synthetic fabrics, furs, and 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 hair is severely damaged by peroxide);

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

• To clean surfaces, you need to pour peroxide into a container with a spray bottle and spray it on contaminated areas.
• To disinfect objects, they need to be wiped with an undiluted H2O2 solution. This will help cleanse them of harmful microorganisms. Washing sponges can be soaked in water with peroxide (1:1 ratio).
• To bleach fabrics, add a glass of peroxide when washing white items. You can also rinse white fabrics in water mixed with a glass of H2O2. This method restores whiteness, protects fabrics from yellowing and helps remove stubborn stains.
• To combat mold and mildew, mix peroxide and water in a 1:2 ratio in a container with a spray bottle. Spray the resulting mixture onto contaminated surfaces and after 10 minutes clean them with a brush or sponge.
• You can renew darkened grout in tiles by spraying peroxide on the desired areas. After 30 minutes, you need to thoroughly 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 this solution will shine clean.
• To clean your toothbrush, you need to dip it in an undiluted three percent peroxide solution. Then rinse under strong running water. This method disinfects hygiene items well.
• To disinfect purchased vegetables and fruits, you should spray a solution of 1 part peroxide and 1 part water on them, then rinse them thoroughly with water (can be cold).
• At your summer cottage, using 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 the aeration is turned off, or for another reason, you can try placing them in water with hydrogen peroxide. You need to mix three percent peroxide with water at the rate of 30 ml per 100 liters and place lifeless fish in the resulting mixture 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 vigorously shaking a bottle of water, 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. When heating, cooking, frying and other processes with accompanying high temperatures, a large amount of oxygen is destroyed. This is why cooked foods are considered not so healthy, although some vitamins remain in them. Freshly squeezed juices or oxygen cocktails served in sanatoriums are useful for the same reason - due to saturation with oxygen, which gives the body new strength and cleanses it.

Danger of peroxide when ingested

After the above, it may seem that peroxide can be specifically taken orally, and this will benefit the body. But this is not true at all. In water or juices, the compound is found in minimum quantities and is closely related to other substances. Taking “unnatural” hydrogen peroxide internally (and all peroxide purchased in a store or produced as a result chemical experiments on its own, cannot in any way be considered natural; moreover, it has too high a concentration compared to natural) can lead to consequences dangerous to life and health. To understand why, we need to turn again to chemistry.

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 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 treated area with alcohol, you will feel a burning sensation, which confirms the presence of microscopic damage after peroxide. But when low concentration peroxide is used externally, there will be no noticeable harm to the body.

It’s another matter if you try to take it orally. That substance, which can damage even relatively thick skin on the outside, ends up on 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 happen with the cells of the food tract. If burns as a result of the action of the oxidizing agent are repeated, then atrophy of the mucous membranes is possible, and this is the first step on the path to cancer. The death of intestinal cells leads to the body's inability to absorb nutrients, which explains, for example, weight loss and the disappearance of constipation in some people who practice “treatment” with peroxide.

Separately, it is necessary to say about this method of using peroxide, such 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 strict dosage calculations, there are still risks. But in such extreme situation this will be a chance for recovery. Under no circumstances should you prescribe hydrogen peroxide injections to yourself. H2O2 poses a great danger to blood cells - red blood cells and platelets, since it destroys them when it enters the bloodstream. In addition, a fatal blockage of blood vessels by the released oxygen can occur - a gas embolism.

Safety precautions for handling H2O2

• Keep out of the reach of children and disabled persons. The lack of odor and distinct taste makes peroxide especially dangerous for them, since large doses can be taken. If the solution gets inside, the consequences of use can be unpredictable. You should consult a doctor immediately.
• Peroxide solutions with a concentration of more than three percent cause burns if they come into contact with the skin. The burn area should be washed with plenty of water.

• Do not allow the peroxide solution to get into your eyes, as this will cause swelling, redness, irritation, and sometimes pain. First aid before contacting a doctor is to wash the eyes generously 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 use for other purposes.
• Storage conditions that prolong its life are a dark, dry, cool place.
• Hydrogen peroxide should not be mixed with any liquids other than clean water, including chlorinated tap water.
• All of the above applies not only to H2O2, but also to all preparations containing it.

2H2 + O2 ––> 2H2O

the concentrations of hydrogen, oxygen and water change to varying degrees: ΔC(H2) = ΔC(H2O) = 2 ΔC(O2).

The rate of a chemical reaction depends on many factors: the nature of the reactants, their concentration, temperature, the nature of the solvent, etc.

2.1.1 Kinetic equation of a chemical reaction. Order of reaction.

One of the tasks facing chemical kinetics is determining the composition of the reaction mixture (i.e., the concentrations of all reagents) at any time, for which it is necessary to know the dependence of the reaction rate on concentrations. In general, the greater the concentration of reactants, the greater the rate of the chemical reaction. Chemical kinetics is based on the so-called. the basic postulate of chemical kinetics:

The rate of a chemical reaction is directly proportional to the product of the concentrations of the reacting substances, taken to certain powers.

That is, for the reaction

aA + bB + dD + . ––> eE + .

can be written:

The proportionality coefficient k is the rate constant of a chemical reaction. The rate constant is numerically equal to the reaction rate at concentrations of all reactants equal to 1 mol/l.

The dependence of the reaction rate on the concentrations of the reactants is determined experimentally and is called the kinetic equation of a chemical reaction. Obviously, in order to write the kinetic equation, it is necessary to experimentally determine the value of the rate constant and exponents at the concentrations of the reacting substances. The exponent at the concentration of each of the reactants in the kinetic equation of a chemical reaction (in equation (II.4) x, y and z, respectively) is the particular order of the reaction for this component. The sum of the exponents in the kinetic equation of a chemical reaction (x + y + z) represents the overall order of the reaction. It should be emphasized that the reaction order is determined only from experimental data and is not related to stoichiometric coefficients with the reactants in the reaction equation. The stoichiometric equation of a reaction is a material balance equation and in no way can determine the nature of the course of this reaction over time.

In chemical kinetics, it is customary to classify reactions according to their magnitude general order reactions. Let us consider the dependence of the concentration of reactants on time for irreversible (one-sided) reactions of zero, first and second orders.

2.1.2 Zero order reactions

For zero-order reactions, the kinetic equation has the following form:

The rate of a zero-order reaction is constant over time and does not depend on the concentrations of the reactants; This is typical for many heterogeneous reactions (taking place at the phase interface) in the case when the rate of diffusion of reagents to the surface is less than the rate of their chemical transformation.

2.1.3 First order reactions

Let us consider the time dependence of the concentration of the starting substance A for the case of a first-order reaction A ––> B. First-order reactions are characterized by a kinetic equation of the form (II.6). Let us substitute expression (II.2) into it:

(II.7)

After integrating expression (II.7) we obtain:

We determine the integration constant g from the initial conditions: at time t = 0, the concentration of C is equal to the initial concentration of Co. It follows that g = ln Co. We get:

Rice. 2.3 Dependence of the logarithm of concentration on time for first-order reactions

Thus, the logarithm of concentration for a first-order reaction depends linearly on time (Fig. 2.3) and the rate constant is numerically equal to the tangent of the angle of inclination of the straight line to the time axis.

From equation (II.9) it is easy to obtain an expression for the rate constant of a one-way first-order reaction:

Another kinetic characteristic of the reaction is the half-life t1/2 - the time during which the concentration of the starting substance decreases by half compared to the original. Let us express t1/2 for a first-order reaction, taking into account that C = ½Co:

(II.12)

As can be seen from the resulting expression, the half-life of the first-order reaction does not depend on the initial concentration of the starting substance.

2.1.4 Second order reactions

For second-order reactions, the kinetic equation has the following form:

Let us consider the simplest case when the kinetic equation has the form (II.14) or, what is the same, in an equation of the form (II.15) the concentrations of the starting substances are the same; equation (II.14) in this case can be rewritten as follows:

(II.16)

After separation of variables and integration we get:

The integration constant g, as in the previous case, is determined from the initial conditions. We get:

Thus, for second-order reactions with a kinetic equation of the form (II.14), a linear dependence of the inverse concentration on time is characteristic (Fig. 2.4) and the rate constant is equal to the tangent of the angle of inclination of the straight line to the time axis:

(II.20)

Rice. 2.4 Dependence of inverse concentration on time for second order reactions

If the initial concentrations of the reactants Co, A and Co, B are different, then the reaction rate constant is found by integrating equation (II.21), in which CA and CB are the concentrations of the reactants at time t from the start of the reaction:

(II.21)

In this case, for the rate constant we obtain the expression