Interaction of oxygen with hydrogen equation. Why does water not burn, although it consists of flammable substances (hydrogen and oxygen). Reactions of hydrogen with simple substances

Oxygen- one of the most common elements on Earth. It makes up about half the weight of the Earth's crust, the planet's outer shell. When combined with hydrogen, it forms water, which covers more than two-thirds of the earth's surface.

We cannot see oxygen, nor can we taste or smell it. However, it makes up one fifth of air and is essential for life. To live, we, just like animals and plants, need to breathe.

Oxygen is an indispensable participant in the chemical reactions that take place inside any microscopic cell of a living organism, as a result of which nutrients are broken down and the energy necessary for life is released. This is why oxygen is so necessary for every living creature (with the exception of a few types of microbes).

When burning, substances combine with oxygen, releasing energy in the form of heat and light.

Hydrogen

The most abundant element in the Universe is hydrogen. It accounts for the bulk of most stars. On Earth, most hydrogen (chemical symbol H) combines with oxygen (O) to form water (H20). Hydrogen is the simplest and lightest chemical element, since each of its atoms consists of only one proton and one electron.

At the beginning of the 20th century, airships and large aircraft were filled with hydrogen. However, hydrogen is very flammable. After several disasters caused by fires, hydrogen was no longer used in airships. Today, another light gas is used in aeronautics - non-flammable helium.

Hydrogen combines with carbon to form substances called hydrocarbons. These include products derived from natural gas and crude oil, such as propane and butane gases, or liquid gasoline. Hydrogen also combines with carbon and oxygen to form carbohydrates. Starch in potatoes and rice, sugar in beets are carbohydrates.

The Sun and other stars are mostly made of hydrogen. At the center of the star, monstrous temperatures and pressures force hydrogen atoms to fuse with each other and turn into another gas - helium. This releases a huge amount of energy in the form of heat and light.

• Designation - H (Hydrogen);
• Latin name - Hydrogenium;
• Period - I;
• Group - 1 (Ia);
• Atomic mass - 1.00794;
• Atomic number - 1;
• Atomic radius = 53 pm;
• Covalent radius = 32 pm;
• Electron distribution - 1s 1;
• melting temperature = -259.14°C;
• boiling point = -252.87°C;
• Electronegativity (according to Pauling/according to Alpred and Rochow) = 2.02/-;
• Oxidation state: +1; 0; -1;
• Density (no.) = 0.0000899 g/cm 3 ;
• Molar volume = 14.1 cm 3 /mol.

Binary compounds of hydrogen with oxygen:

Hydrogen (“giving birth to water”) was discovered by the English scientist G. Cavendish in 1766. It is the simplest element in nature - a hydrogen atom has a nucleus and one electron, which is probably why hydrogen is the most abundant element in the Universe (accounting for more than half the mass of most stars).

About hydrogen we can say that “the spool is small, but expensive.” Despite its “simplicity,” hydrogen provides energy to all living beings on Earth - a continuous thermonuclear reaction takes place on the Sun during which one helium atom is formed from four hydrogen atoms, this process is accompanied by the release of a colossal amount of energy (for more details, see Nuclear fusion).

In the earth's crust, the mass fraction of hydrogen is only 0.15%. Meanwhile, the overwhelming majority (95%) of all chemical substances known on Earth contain one or more hydrogen atoms.

In compounds with non-metals (HCl, H 2 O, CH 4 ...), hydrogen gives up its only electron to more electronegative elements, exhibiting an oxidation state of +1 (more often), forming only covalent bonds (see Covalent bond).

In compounds with metals (NaH, CaH 2 ...), hydrogen, on the contrary, accepts another electron into its only s-orbital, thus trying to complete its electronic layer, exhibiting an oxidation state of -1 (less often), often forming an ionic bond (see Ionic bond), because the difference in electronegativity of the hydrogen atom and the metal atom can be quite large.

H 2

In the gaseous state, hydrogen exists in the form of diatomic molecules, forming a nonpolar covalent bond.

Hydrogen molecules have:

• great mobility;
• great strength;
• low polarizability;
• small size and weight.

Properties of hydrogen gas:

• the lightest gas in nature, colorless and odorless;
• poorly soluble in water and organic solvents;
• dissolves in small amounts in liquid and solid metals (especially platinum and palladium);
• difficult to liquefy (due to its low polarizability);
• has the highest thermal conductivity of all known gases;
• when heated, it reacts with many non-metals, exhibiting the properties of a reducing agent;
• at room temperature it reacts with fluorine (an explosion occurs): H 2 + F 2 = 2HF;
• reacts with metals to form hydrides, exhibiting oxidizing properties: H 2 + Ca = CaH 2 ;

In compounds, hydrogen exhibits its reducing properties much more strongly than its oxidizing properties. Hydrogen is the most powerful reducing agent after coal, aluminum and calcium. The reducing properties of hydrogen are widely used in industry to obtain metals and nonmetals (simple substances) from oxides and gallides.

Fe 2 O 3 + 3H 2 = 2Fe + 3H 2 O

Reactions of hydrogen with simple substances

Hydrogen accepts an electron, playing a role reducing agent, in reactions:

• With oxygen(when ignited or in the presence of a catalyst), in a ratio of 2:1 (hydrogen:oxygen) an explosive detonating gas is formed: 2H 2 0 +O 2 = 2H 2 +1 O+572 kJ
• With gray(when heated to 150°C-300°C): H 2 0 +S ↔ H 2 +1 S
• With chlorine(when ignited or irradiated with UV rays): H 2 0 +Cl 2 = 2H +1 Cl
• With fluorine: H 2 0 +F 2 = 2H +1 F
• With nitrogen(when heated in the presence of catalysts or at high pressure): 3H 2 0 +N 2 ↔ 2NH 3 +1

Hydrogen donates an electron, playing a role oxidizing agent, in reactions with alkaline And alkaline earth metals with the formation of metal hydrides - salt-like ionic compounds containing hydride ions H - these are unstable white crystalline substances.

Ca+H 2 = CaH 2 -1 2Na+H 2 0 = 2NaH -1

It is not typical for hydrogen to exhibit an oxidation state of -1. When reacting with water, the hydrides decompose, reducing water to hydrogen. The reaction of calcium hydride with water is as follows:

CaH 2 -1 +2H 2 +1 0 = 2H 2 0 +Ca(OH) 2

Reactions of hydrogen with complex substances

• at high temperatures, hydrogen reduces many metal oxides: ZnO+H 2 = Zn+H 2 O
• methyl alcohol is obtained by the reaction of hydrogen with carbon monoxide (II): 2H 2 +CO → CH 3 OH
• In hydrogenation reactions, hydrogen reacts with many organic substances.

The equations of chemical reactions of hydrogen and its compounds are discussed in more detail on the page “Hydrogen and its compounds - equations of chemical reactions involving hydrogen.”

Applications of hydrogen

• in nuclear energy, hydrogen isotopes are used - deuterium and tritium;
• in the chemical industry, hydrogen is used for the synthesis of many organic substances, ammonia, hydrogen chloride;
• in the food industry, hydrogen is used in the production of solid fats through the hydrogenation of vegetable oils;
• for welding and cutting metals, the high combustion temperature of hydrogen in oxygen (2600°C) is used;
• in the production of some metals, hydrogen is used as a reducing agent (see above);
• since hydrogen is a light gas, it is used in aeronautics as a filler for balloons, aerostats, and airships;
• Hydrogen is used as a fuel mixed with CO.

Recently, scientists have been paying a lot of attention to the search for alternative sources of renewable energy. One of the promising areas is “hydrogen” energy, in which hydrogen is used as fuel, the combustion product of which is ordinary water.

Methods for producing hydrogen

Industrial methods for producing hydrogen:

• methane conversion (catalytic reduction of water vapor) with water vapor at high temperature (800°C) on a nickel catalyst: CH 4 + 2H 2 O = 4H 2 + CO 2 ;
• conversion of carbon monoxide with water vapor (t=500°C) on a Fe 2 O 3 catalyst: CO + H 2 O = CO 2 + H 2 ;
• thermal decomposition of methane: CH 4 = C + 2H 2;
• gasification of solid fuels (t=1000°C): C + H 2 O = CO + H 2 ;
• electrolysis of water (a very expensive method that produces very pure hydrogen): 2H 2 O → 2H 2 + O 2.

Laboratory methods for producing hydrogen:

• action on metals (usually zinc) with hydrochloric or dilute sulfuric acid: Zn + 2HCl = ZCl 2 + H 2 ; Zn + H 2 SO 4 = ZnSO 4 + H 2;
• interaction of water vapor with hot iron filings: 4H 2 O + 3Fe = Fe 3 O 4 + 4H 2.

Chemical properties of hydrogen

Under ordinary conditions, molecular Hydrogen is relatively little active, directly combining only with the most active of non-metals (with fluorine, and in the light with chlorine). However, when heated, it reacts with many elements.

Hydrogen reacts with simple and complex substances:

- Interaction of hydrogen with metals leads to the formation of complex substances - hydrides, in the chemical formulas of which the metal atom always comes first:

At high temperature, Hydrogen reacts directly with some metals(alkaline, alkaline earth and others), forming white crystalline substances - metal hydrides (Li H, Na H, KH, CaH 2, etc.):

H 2 + 2Li = 2LiH

Metal hydrides are easily decomposed by water to form the corresponding alkali and hydrogen:

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

- When hydrogen interacts with non-metals volatile hydrogen compounds are formed. In the chemical formula of a volatile hydrogen compound, the hydrogen atom can be in either the first or second place, depending on its location in the PSHE (see plate in the slide):

1). With oxygen Hydrogen forms water:

Video "Hydrogen combustion"

2H 2 + O 2 = 2H 2 O + Q

At normal temperatures the reaction proceeds extremely slowly, above 550°C - with explosion (a mixture of 2 volumes of H 2 and 1 volume of O 2 is called explosive gas) .

Video "Explosion of detonating gas"

Video "Preparation and explosion of an explosive mixture"

2). With halogens Hydrogen forms hydrogen halides, for example:

H 2 + Cl 2 = 2HCl

At the same time, Hydrogen explodes with fluorine (even in the dark and at - 252°C), reacts with chlorine and bromine only when illuminated or heated, and with iodine only when heated.

3). With nitrogen Hydrogen reacts to form ammonia:

ZN 2 + N 2 = 2NH 3

only on a catalyst and at elevated temperatures and pressures.

4). When heated, Hydrogen reacts vigorously with sulfur:

H 2 + S = H 2 S (hydrogen sulfide),

much more difficult with selenium and tellurium.

5). With pure carbon Hydrogen can react without a catalyst only at high temperatures:

2H 2 + C (amorphous) = CH 4 (methane)

- Hydrogen undergoes a substitution reaction with metal oxides , in this case water is formed in the products and the metal is reduced. Hydrogen - exhibits the properties of a reducing agent:

Hydrogen is used for the recovery of many metals, since it takes oxygen away from their oxides:

Fe 3 O 4 + 4H 2 = 3Fe + 4H 2 O, etc.

Applications of hydrogen

Video "Using Hydrogen"

Currently, hydrogen is produced in huge quantities. A very large part of it is used in the synthesis of ammonia, hydrogenation of fats and in the hydrogenation of coal, oils and hydrocarbons. In addition, hydrogen is used for the synthesis of hydrochloric acid, methyl alcohol, hydrocyanic acid, in welding and forging metals, as well as in the manufacture of incandescent lamps and precious stones. Hydrogen is sold in cylinders under a pressure of over 150 atm. They are painted dark green and have a red inscription "Hydrogen".

Hydrogen is used to convert liquid fats into solid fats (hydrogenation), producing liquid fuel by hydrogenating coal and fuel oil. In metallurgy, hydrogen is used as a reducing agent for oxides or chlorides to produce metals and non-metals (germanium, silicon, gallium, zirconium, hafnium, molybdenum, tungsten, etc.).

The practical uses of hydrogen are varied: it is usually used to fill probe balloons, in the chemical industry it serves as a raw material for the production of many very important products (ammonia, etc.), in the food industry - for the production of solid fats from vegetable oils, etc. High temperature (up to 2600 °C), obtained by burning hydrogen in oxygen, is used for melting refractory metals, quartz, etc. Liquid hydrogen is one of the most efficient jet fuels. Annual global consumption of hydrogen exceeds 1 million tons.

SIMULATORS

No. 2. Hydrogen

ASSIGNMENT TASKS

Task No. 1
Write down reaction equations for the interaction of hydrogen with the following substances: F 2, Ca, Al 2 O 3, mercury (II) oxide, tungsten (VI) oxide. Name the reaction products, indicate the types of reactions.

Task No. 2
Carry out transformations according to the scheme:
H 2 O -> H 2 -> H 2 S -> SO 2

Task No. 3.
Calculate the mass of water that can be obtained by burning 8 g of hydrogen?

In the periodic table, hydrogen is located in two groups of elements that are completely opposite in their properties. This feature makes it completely unique. Hydrogen is not just an element or substance, but is also an integral part of many complex compounds, an organogenic and biogenic element. Therefore, let's look at its properties and characteristics in more detail.

The release of flammable gas during the interaction of metals and acids was observed back in the 16th century, that is, during the formation of chemistry as a science. The famous English scientist Henry Cavendish studied the substance starting in 1766 and gave it the name “combustible air”. When burned, this gas produced water. Unfortunately, the scientist’s adherence to the theory of phlogiston (hypothetical “ultrafine matter”) prevented him from coming to the right conclusions.

The French chemist and naturalist A. Lavoisier, together with the engineer J. Meunier and with the help of special gasometers, synthesized water in 1783, and then analyzed it through the decomposition of water vapor with hot iron. Thus, scientists were able to come to the right conclusions. They found that “combustible air” is not only part of water, but can also be obtained from it.

In 1787, Lavoisier suggested that the gas under study was a simple substance and, accordingly, was one of the primary chemical elements. He called it hydrogene (from the Greek words hydor - water + gennao - I give birth), i.e. “giving birth to water.”

The Russian name “hydrogen” was proposed in 1824 by the chemist M. Soloviev. The determination of the composition of water marked the end of the “phlogiston theory.” At the turn of the 18th and 19th centuries, it was established that the hydrogen atom is very light (compared to the atoms of other elements) and its mass was taken as the basic unit for comparing atomic masses, receiving a value equal to 1.

Physical properties

Hydrogen is the lightest substance known to science (it is 14.4 times lighter than air), its density is 0.0899 g/l (1 atm, 0 °C). This material melts (solidifies) and boils (liquefies), respectively, at -259.1 ° C and -252.8 ° C (only helium has lower boiling and melting temperatures).

The critical temperature of hydrogen is extremely low (-240 °C). For this reason, its liquefaction is a rather complex and costly process. The critical pressure of the substance is 12.8 kgf/cm², and the critical density is 0.0312 g/cm³. Among all gases, hydrogen has the highest thermal conductivity: at 1 atm and 0 °C it is equal to 0.174 W/(mxK).

The specific heat capacity of the substance under the same conditions is 14.208 kJ/(kgxK) or 3.394 cal/(rx°C). This element is slightly soluble in water (about 0.0182 ml/g at 1 atm and 20 °C), but well soluble in most metals (Ni, Pt, Pa and others), especially in palladium (about 850 volumes per volume of Pd ).

The latter property is associated with its ability to diffuse, and diffusion through a carbon alloy (for example, steel) can be accompanied by the destruction of the alloy due to the interaction of hydrogen with carbon (this process is called decarbonization). In the liquid state, the substance is very light (density - 0.0708 g/cm³ at t° = -253 °C) and fluid (viscosity - 13.8 spoise under the same conditions).

In many compounds, this element exhibits a +1 valency (oxidation state), like sodium and other alkali metals. It is usually considered as an analogue of these metals. Accordingly, he heads group I of the periodic system. In metal hydrides, the hydrogen ion exhibits a negative charge (the oxidation state is -1), that is, Na+H- has a structure similar to Na+Cl- chloride. In accordance with this and some other facts (the similarity of the physical properties of the element “H” and halogens, the ability to replace it with halogens in organic compounds), Hydrogene is classified in group VII of the periodic system.

Under normal conditions, molecular hydrogen has low activity, directly combining only with the most active of non-metals (with fluorine and chlorine, with the latter in the light). In turn, when heated, it interacts with many chemical elements.

Atomic hydrogen has increased chemical activity (compared to molecular hydrogen). With oxygen it forms water according to the formula:

Н₂ + ½О₂ = Н₂О,

releasing 285.937 kJ/mol of heat or 68.3174 kcal/mol (25 °C, 1 atm). Under normal temperature conditions, the reaction proceeds rather slowly, and at t° >= 550 °C it is uncontrollable. The explosive limits of a hydrogen + oxygen mixture by volume are 4–94% H₂, and a hydrogen + air mixture is 4–74% H₂ (a mixture of two volumes of H₂ and one volume of O₂ is called detonating gas).

This element is used to reduce most metals, as it removes oxygen from oxides:

Fe₃O₄ + 4H₂ = 3Fe + 4H₂O,

CuO + H₂ = Cu + H₂O, etc.

Hydrogen forms hydrogen halides with different halogens, for example:

H₂ + Cl₂ = 2HCl.

However, when reacting with fluorine, hydrogen explodes (this also happens in the dark, at -252 ° C), with bromine and chlorine it reacts only when heated or illuminated, and with iodine - only when heated. When interacting with nitrogen, ammonia is formed, but only on a catalyst, at elevated pressures and temperatures:

ЗН₂ + N₂ = 2NN₃.

When heated, hydrogen reacts actively with sulfur:

H₂ + S = H₂S (hydrogen sulfide),

and much more difficult with tellurium or selenium. Hydrogen reacts with pure carbon without a catalyst, but at high temperatures:

2H₂ + C (amorphous) = CH₄ (methane).

This substance reacts directly with some of the metals (alkali, alkaline earth and others), forming hydrides, for example:

H₂ + 2Li = 2LiH.

The interactions between hydrogen and carbon monoxide (II) are of considerable practical importance. In this case, depending on the pressure, temperature and catalyst, different organic compounds are formed: HCHO, CH₃OH, etc. Unsaturated hydrocarbons during the reaction become saturated, for example:

С n Н₂ n + Н₂ = С n Н₂ n ₊₂.

Hydrogen and its compounds play an exceptional role in chemistry. It determines the acidic properties of the so-called. protic acids, tends to form hydrogen bonds with various elements, which have a significant impact on the properties of many inorganic and organic compounds.

Hydrogen production

The main types of raw materials for the industrial production of this element are oil refining gases, natural combustible and coke oven gases. It is also obtained from water through electrolysis (in places where electricity is available). One of the most important methods for producing material from natural gas is the catalytic interaction of hydrocarbons, mainly methane, with water vapor (so-called conversion). For example:

CH₄ + H₂O = CO + ZN₂.

Incomplete oxidation of hydrocarbons with oxygen:

CH₄ + ½O₂ = CO + 2H₂.

The synthesized carbon monoxide (II) undergoes conversion:

CO + H₂O = CO₂ + H₂.

Hydrogen produced from natural gas is the cheapest.

For the electrolysis of water, direct current is used, which is passed through a solution of NaOH or KOH (acids are not used to avoid corrosion of the equipment). In laboratory conditions, the material is obtained by electrolysis of water or as a result of the reaction between hydrochloric acid and zinc. However, ready-made factory material in cylinders is more often used.

This element is isolated from oil refining gases and coke oven gas by removing all other components of the gas mixture, since they liquefy more easily during deep cooling.

This material began to be produced industrially at the end of the 18th century. Back then it was used to fill balloons. At the moment, hydrogen is widely used in industry, mainly in the chemical industry, for the production of ammonia.

Mass consumers of the substance are producers of methyl and other alcohols, synthetic gasoline and many other products. They are obtained by synthesis from carbon monoxide (II) and hydrogen. Hydrogene is used for the hydrogenation of heavy and solid liquid fuels, fats, etc., for the synthesis of HCl, hydrotreating of petroleum products, as well as in metal cutting/welding. The most important elements for nuclear energy are its isotopes - tritium and deuterium.

Biological role of hydrogen

About 10% of the mass of living organisms (on average) comes from this element. It is part of water and the most important groups of natural compounds, including proteins, nucleic acids, lipids, and carbohydrates. What is it used for?

This material plays a decisive role: in maintaining the spatial structure of proteins (quaternary), in implementing the principle of complementarity of nucleic acids (i.e., in the implementation and storage of genetic information), and in general in “recognition” at the molecular level.

The hydrogen ion H+ takes part in important dynamic reactions/processes in the body. Including: in biological oxidation, which provides living cells with energy, in biosynthesis reactions, in photosynthesis in plants, in bacterial photosynthesis and nitrogen fixation, in maintaining acid-base balance and homeostasis, in membrane transport processes. Along with carbon and oxygen, it forms the functional and structural basis of life phenomena.

Oxygen is the most abundant element on Earth. Together with nitrogen and a small amount of other gases, free oxygen forms the Earth's atmosphere. Its content in the air is 20.95% by volume or 23.15% by mass. In the earth's crust, 58% of the atoms are bound oxygen atoms (47% by mass). Oxygen is part of water (the reserves of bound oxygen in the hydrosphere are extremely large), rocks, many minerals and salts, and is found in fats, proteins and carbohydrates that make up living organisms. Almost all of the Earth's free oxygen is created and preserved as a result of the process of photosynthesis.

Physical properties.

Oxygen is a colorless, tasteless, and odorless gas, slightly heavier than air. It is slightly soluble in water (31 ml of oxygen dissolves in 1 liter of water at 20 degrees), but it is still better than other atmospheric gases, so water is enriched with oxygen. Oxygen density under normal conditions is 1.429 g/l. At a temperature of -183 0 C and a pressure of 101.325 kPa, oxygen turns into a liquid state. Liquid oxygen has a bluish color, is drawn into a magnetic field, and at -218.7 ° C, forms blue crystals.

Natural oxygen has three isotopes O 16, O 17, O 18.

Allotropy- the ability of a chemical element to exist in the form of two or more simple substances that differ only in the number of atoms in the molecule or in structure.

Ozone O 3 – exists in the upper layers of the atmosphere at an altitude of 20-25 km from the Earth’s surface and forms the so-called “ozone layer”, which protects the Earth from the harmful ultraviolet radiation of the Sun; a pale purple, poisonous gas in large quantities with a specific, pungent but pleasant odor. The melting point is -192.7 0 C, the boiling point is 111.9 0 C. We dissolve oxygen better in water.

Ozone is a strong oxidizing agent. Its oxidative activity is based on the ability of the molecule to decompose with the release of atomic oxygen:

It oxidizes many simple and complex substances. With some metals it forms ozonides, for example potassium ozonide:

K + O 3 = KO 3

Ozone is produced in special devices - ozonizers. In them, under the influence of an electric discharge, molecular oxygen is converted into ozone:

A similar reaction occurs under the influence of lightning discharges.

The use of ozone is due to its strong oxidizing properties: it is used to bleach fabrics, disinfect drinking water, and in medicine as a disinfectant.

Inhaling ozone in large quantities is harmful: it irritates the mucous membranes of the eyes and respiratory organs.

Chemical properties.

In chemical reactions with atoms of other elements (except fluorine), oxygen exhibits exclusively oxidizing properties

The most important chemical property is the ability to form oxides with almost all elements. At the same time, oxygen reacts directly with most substances, especially when heated.

As a result of these reactions, as a rule, oxides are formed, less often peroxides:

2Ca + O 2 = 2CaO

2Ba + O 2 = 2BaO

2Na + O 2 = Na 2 O 2

Oxygen does not interact directly with halogens, gold, and platinum; their oxides are obtained indirectly. When heated, sulfur, carbon, and phosphorus burn in oxygen.

The interaction of oxygen with nitrogen begins only at a temperature of 1200 0 C or in an electrical discharge:

N 2 + O 2 = 2NO

With hydrogen, oxygen forms water:

2H 2 + O 2 = 2H 2 O

During this reaction, a significant amount of heat is released.

A mixture of two volumes of hydrogen with one volume of oxygen explodes when ignited; it is called detonating gas.

Many metals upon contact with atmospheric oxygen are subject to destruction - corrosion. Some metals under normal conditions are oxidized only from the surface (for example, aluminum, chromium). The resulting oxide film prevents further interaction.

4Al + 3O 2 = 2Al 2 O 3

Under certain conditions, complex substances also interact with oxygen. In this case, oxides are formed, and in some cases, oxides and simple substances.

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

H 2 S + O 2 = 2SO 2 + 2H 2 O

4NН 3 +ЗО 2 =2N 2 +6Н 2 О

4CH 3 NH 2 + 9O 2 = 4CO 2 + 2N 2 + 10H 2 O

When interacting with complex substances, oxygen acts as an oxidizing agent. Its important property, the ability to maintain combustion substances.

Oxygen also forms a compound with hydrogen - hydrogen peroxide H 2 O 2 - a colorless transparent liquid with a pungent astringent taste, highly soluble in water. Chemically, hydrogen peroxide is a very interesting compound. Its low stability is characteristic: when standing, it slowly decomposes into water and oxygen:

H 2 O 2 = H 2 O + O 2

Light, heat, the presence of alkalis, and contact with oxidizing or reducing agents accelerate the decomposition process. The oxidation state of oxygen in hydrogen peroxide = - 1, i.e. has an intermediate value between the oxidation state of oxygen in water (-2) and in molecular oxygen (0), so hydrogen peroxide exhibits redox duality. The oxidizing properties of hydrogen peroxide are much more pronounced than the reducing properties, and they manifest themselves in acidic, alkaline and neutral environments.

H 2 O 2 + 2KI + H 2 SO 4 = K 2 SO 4 + I 2 + 2H 2 O