Oxygen and its properties. Compounds of oxygen with hydrogen. Physical and chemical properties of hydrogen

The hydrogen atom has electronic formula external (and only) electronic level 1 s 1 . On the one hand, by the presence of one electron at the outer electronic level, the hydrogen atom is similar to the atoms of alkali metals. However, he, like halogens, lacks only one electron to fill the external electronic level, since no more than 2 electrons can be located at the first electronic level. It turns out that hydrogen can be placed simultaneously in both the first and the penultimate (seventh) groups of the periodic table, which is sometimes done in different options periodic system:

In terms of the properties of hydrogen as a simple substance, it still has more in common with halogens. Hydrogen, like halogens, is a non-metal and forms, like them, diatomic molecules (H 2).

Under normal conditions, hydrogen is a gaseous, low-activity substance. The low activity of hydrogen is explained by the high strength of the bond between the hydrogen atoms in the molecule, which requires either strong heating, or the use of catalysts, or both at the same time to break it.

Interaction of hydrogen with simple substances

with metals

Of metals, hydrogen reacts only with alkali and alkaline earth! Alkali metals include metals of the main subgroup Group I(Li, Na, K, Rb, Cs, Fr), and to alkaline earth metals - metals of the main subgroup of group II, except for beryllium and magnesium (Ca, Sr, Ba, Ra)

When interacting with active metals, hydrogen exhibits oxidizing properties, i.e. lowers its oxidation state. In this case, hydrides of alkali and alkaline earth metals are formed, which have an ionic structure. This reaction takes place by heating:

It should be noted that the interaction with active metals is the only case when molecular hydrogen H 2 is an oxidizing agent.

with non-metals

Of non-metals, hydrogen reacts only with carbon, nitrogen, oxygen, sulfur, selenium and halogens!

Carbon should be understood as graphite or amorphous carbon, since diamond is an extremely inert allotropic modification of carbon.

When interacting with non-metals, hydrogen can only perform the function of a reducing agent, that is, only increase its oxidation state:

Interaction of hydrogen with complex substances

with metal oxides

Hydrogen does not react with metal oxides that are in the range of metal activity up to aluminum (inclusive), however, it is able to reduce many metal oxides to the right of aluminum when heated:

with oxides of non-metals

Of the oxides of non-metals, hydrogen reacts when heated with oxides of nitrogen, halogens and carbon. Of all the interactions of hydrogen with oxides of non-metals, its reaction with carbon monoxide CO should be especially noted.

The mixture of CO and H 2 even has its own name - "synthesis gas", since, depending on the conditions, such popular industrial products as methanol, formaldehyde and even synthetic hydrocarbons can be obtained from it:

with acids

Hydrogen does not react with inorganic acids!

Of organic acids, hydrogen reacts only with unsaturated ones, as well as with acids containing functional groups capable of being reduced by hydrogen, in particular aldehyde, keto or nitro groups.

with salts

In the case of aqueous solutions of salts, their interaction with hydrogen does not occur. However, when hydrogen is passed over solid salts of some metals of medium and low activity, their partial or complete reduction is possible, for example:

Chemical properties of halogens

Chemical elements of group VIIA (F, Cl, Br, I, At), as well as the simple substances formed by them, are called halogens. Here and further in the text, unless otherwise stated, under halogens we mean just simple substances.

All halogens have a molecular structure, which leads to low melting and boiling points of these substances. Halogen molecules are diatomic, i.e. their formula can be written in general view like Hal 2.

It should be noted such a specific physical property of iodine as its ability to sublimation or, in other words, sublimation. Sublimation, is called the phenomenon in which a substance in a solid state does not melt when heated, but, bypassing the liquid phase, immediately passes into a gaseous state.

The electronic structure of the external energy level of an atom of any halogen has the form ns 2 np 5, where n is the number of the period of the periodic table in which the halogen is located. As you can see, up to the eight-electron outer shell, halogen atoms lack only one electron. From this, it is logical to assume the predominantly oxidizing properties of free halogens, which is also confirmed in practice. As you know, the electronegativity of non-metals decreases when moving down the subgroup, and therefore the activity of halogens decreases in the following order:

F 2> Cl 2> Br 2> I 2

Interaction of halogens with simple substances

All halogens are highly reactive and react with most simple substances. However, it should be noted that fluorine, due to its extremely high reactivity, can react even with those simple substances with which other halogens cannot react. These simple substances include oxygen, carbon (diamond), nitrogen, platinum, gold, and some noble gases (xenon and krypton). Those. actually, fluorine does not react only with some noble gases.

The rest of the halogens, i.e. chlorine, bromine and iodine are also active substances, but less active than fluorine. They react with almost all simple substances except oxygen, nitrogen, carbon in the form of diamond, platinum, gold and noble gases.

Interaction of halogens with non-metals

hydrogen

When all halogens react with hydrogen, hydrogen halides with general formula HHal. At the same time, the reaction of fluorine with hydrogen begins spontaneously even in the dark and proceeds with an explosion in accordance with the equation:

The reaction of chlorine with hydrogen can be initiated by intense ultraviolet irradiation or heating. Also proceeds with an explosion:

Bromine and iodine react with hydrogen only when heated, and at the same time, the reaction with iodine is reversible:

phosphorus

The interaction of fluorine with phosphorus leads to the oxidation of phosphorus to the highest oxidation state (+5). In this case, the formation of phosphorus pentafluoride occurs:

When chlorine and bromine interacts with phosphorus, it is possible to obtain phosphorus halides both in the + 3 oxidation state and in the +5 oxidation state, which depends on the proportions of the reactants:

In this case, in the case of white phosphorus in an atmosphere of fluorine, chlorine or liquid bromine, the reaction starts spontaneously.

The interaction of phosphorus with iodine can lead to the formation of only phosphorus triodide due to the significantly lower oxidizing ability than that of other halogens:

gray

Fluorine oxidizes sulfur to the highest oxidation state +6, forming sulfur hexafluoride:

Chlorine and bromine react with sulfur, forming compounds containing sulfur in the extremely unusual oxidation states of +1 and +2. These interactions are very specific, and for passing the exam in chemistry, the ability to write down the equations of these interactions is not necessary. Therefore, the following three equations are given rather for informational purposes:

Interaction of halogens with metals

As mentioned above, fluorine is capable of reacting with all metals, even such inactive ones as platinum and gold:

The rest of the halogens react with all metals except platinum and gold:

Reactions of halogens with complex substances

Substitution reactions with halogens

More active halogens, i.e. the chemical elements of which are located higher in the periodic table are able to displace less active halogens from the hydrohalic acids and metal halides they form:

Similarly, bromine and iodine displace sulfur from sulfide and or hydrogen sulfide solutions:

Chlorine is a stronger oxidizing agent and oxidizes hydrogen sulfide in its aqueous solution not to sulfur, but to sulfuric acid:

Interaction of halogens with water

Water burns in fluorine with a blue flame in accordance with the reaction equation:

Bromine and chlorine react with water differently than fluorine. If fluorine acted as an oxidizing agent, then chlorine and bromine disproportionate in water, forming a mixture of acids. In this case, the reactions are reversible:

The interaction of iodine with water occurs to such an insignificant extent that it can be neglected and it can be assumed that the reaction does not proceed at all.

Interaction of halogens with alkali solutions

Fluorine, when interacting with an aqueous solution of alkali, again acts as an oxidizing agent:

The ability to write this equation is not required to pass the exam. It is enough to know the fact about the possibility of such interaction and the oxidative role of fluorine in this reaction.

Unlike fluorine, other halogens in alkali solutions disproportionate, that is, they simultaneously increase and decrease their oxidation state. In this case, in the case of chlorine and bromine, depending on the temperature, flow in two different directions is possible. In particular, in the cold, reactions proceed as follows:

and when heated:

Iodine reacts with alkalis exclusively according to the second option, i.e. with the formation of iodate, because hypoioditis is not stable not only when heated, but also at normal temperatures and even in the cold.

Hydrogen H is the most abundant element in the Universe (about 75% by mass), on Earth - the ninth most abundant. The most important natural compound of hydrogen is water.
Hydrogen ranks first in the periodic table (Z = 1). It has the simplest atomic structure: the nucleus of an atom - 1 proton, surrounded by an electron cloud, consisting of 1 electron.
Under some conditions, hydrogen exhibits metallic properties (gives up an electron), in others - non-metallic (accepts an electron).
Hydrogen isotopes occur in nature: 1H - protium (the nucleus consists of one proton), 2H - deuterium (D - the nucleus consists of one proton and one neutron), 3H - tritium (T - the nucleus consists of one proton and two neutrons).

Simple substance hydrogen

A hydrogen molecule consists of two atoms linked together by a covalent non-polar bond.
Physical properties. Hydrogen is a colorless, odorless, tasteless, non-toxic gas. The hydrogen molecule is not polar. Therefore, the forces of intermolecular interaction in gaseous hydrogen are small. This manifests itself in low temperatures boiling (-252.6 0С) and melting (-259.2 0С).
Hydrogen is lighter than air, D (by air) = 0.069; slightly soluble in water (100 volumes of H2O dissolves 2 volumes of H2). Therefore, hydrogen, when produced in a laboratory, can be collected by air or water displacement methods.

Hydrogen production

In the laboratory:

1.The action of dilute acids on metals:
Zn + 2HCl → ZnCl 2 + H 2

2. Interaction of alkaline and u-z metals with water:
Ca + 2H 2 O → Ca (OH) 2 + H 2

3. Hydrolysis of hydrides: metal hydrides are readily decomposed by water to form the corresponding alkali and hydrogen:
NaH + H 2 O → NaOH + H 2
CaH 2 + 2H 2 O = Ca (OH) 2 + 2H 2

4. The action of alkalis on zinc or aluminum or silicon:
2Al + 2NaOH + 6H 2 O → 2Na + 3H 2
Zn + 2KOH + 2H 2 O → K 2 + H 2
Si + 2NaOH + H 2 O → Na 2 SiO 3 + 2H 2

5. Water electrolysis. To increase the electrical conductivity of water, an electrolyte is added to it, for example, NaOH, H 2 SO 4 or Na 2 SO 4. At the cathode, 2 volumes of hydrogen are formed, at the anode - 1 volume of oxygen.
2H 2 O → 2H 2 + O 2

Industrial production of hydrogen

1. Conversion of methane with steam, Ni 800 ° C (the cheapest):
CH 4 + H 2 O → CO + 3 H 2
CO + H 2 O → CO 2 + H 2

In total:
CH 4 + 2 H 2 O → 4 H 2 + CO 2

2. Water vapor through red-hot coke at 1000 о С:
C + H 2 O → CO + H 2
CO + H 2 O → CO 2 + H 2

The resulting carbon monoxide (IV) is absorbed by water, in this way 50% of industrial hydrogen is obtained.

3. Heating methane to 350 ° C in the presence of an iron or nickel catalyst:
CH 4 → C + 2H 2

4. By electrolysis of aqueous solutions of KCl or NaCl, as a by-product:
2Н 2 О + 2NaCl → Cl 2 + H 2 + 2NaOH

Chemical properties of hydrogen

• In compounds, hydrogen is always monovalent. It is characterized by an oxidation state of +1, but in metal hydrides it is -1.
• The hydrogen molecule consists of two atoms. The emergence of a bond between them is explained by the formation of a generalized pair of electrons H: H or H 2
• Due to this generalization of electrons, the H2 molecule is more energetically stable than its individual atoms. To break a molecule into atoms in 1 mole of hydrogen, it is necessary to expend energy of 436 kJ: Н 2 = 2Н, ∆H ° = 436 kJ / mol
• This explains the relatively low activity of molecular hydrogen at ordinary temperatures.
• With many non-metals, hydrogen forms gaseous compounds such as RH 4, RH 3, RH 2, RH.

1) Forms hydrogen halides with halogens:
H 2 + Cl 2 → 2HCl.
At the same time, it explodes with fluorine, reacts with chlorine and bromine only when illuminated or heated, and with iodine only when heated.

2) With oxygen:
2H 2 + O 2 → 2H 2 O
with the release of heat. At ordinary temperatures, the reaction proceeds slowly, above 550 ° C - with an explosion. A mixture of 2 volumes of H 2 and 1 volume of O 2 is called an explosive gas.

3) When heated, it reacts vigorously with sulfur (much more difficult with selenium and tellurium):
H 2 + S → H 2 S (hydrogen sulfide),

4) With nitrogen with the formation of ammonia only on the catalyst and at elevated temperatures and pressures:
ЗН 2 + N 2 → 2NН 3

5) With carbon at high temperatures:
2H 2 + C → CH 4 (methane)

6) Forms hydrides with alkali and alkaline earth metals (hydrogen is an oxidizing agent):
Н 2 + 2Li → 2LiH
in metal hydrides, the hydrogen ion is negatively charged (oxidation state -1), that is, the hydride Na + H - is built like the chloride Na + Cl -

With complex substances:

7) With metal oxides (used for the reduction of metals):
CuO + H 2 → Cu + H 2 O
Fe 3 O 4 + 4H 2 → 3Fe + 4H 2 O

8) with carbon monoxide (II):
CO + 2H 2 → CH 3 OH
Synthesis - gas (mixture of hydrogen and carbon monoxide) has an important practical significance, mk, depending on temperature, pressure and catalyst, various organic compounds are formed, for example HCHO, CH 3 OH and others.

9) Unsaturated hydrocarbons react with hydrogen, turning into saturated ones:
C n H 2n + H 2 → C n H 2n + 2.

10.1 Hydrogen

The name "hydrogen" refers to both a chemical element and a simple substance. Element hydrogen consists of hydrogen atoms. Simple substance hydrogen consists of hydrogen molecules.

a) Chemical element hydrogen

In the natural series of elements, the ordinal number of hydrogen is 1. In the system of elements, hydrogen is in the first period in the IA or VIIA group.

Hydrogen is one of the most abundant elements on Earth. The molar fraction of hydrogen atoms in the atmosphere, hydrosphere and lithosphere of the Earth (all together this is called the earth's crust) is 0.17. It is found in water, many minerals, petroleum, natural gas, plants and animals. The human body contains on average about 7 kilograms of hydrogen.

There are three isotopes of hydrogen:
a) light hydrogen - protium,
b) heavy hydrogen - deuterium(D),
c) superheavy hydrogen - tritium(T).

Tritium is an unstable (radioactive) isotope; therefore, it practically does not occur in nature. Deuterium is stable, but very little of it: w D = 0.015% (of the mass of all terrestrial hydrogen). Therefore, the atomic mass of hydrogen differs very little from 1 D (1.00794 D).

b) Hydrogen atom

From previous chapters chemistry course, you already know the following characteristics of the hydrogen atom:

The valence capabilities of the hydrogen atom are determined by the presence of one electron in a single valence orbital. A high ionization energy makes a hydrogen atom not prone to give up an electron, and a not too high energy of affinity for an electron leads to a slight tendency to accept it. Consequently, in chemical systems the formation of the H cation is impossible, and the compounds with the H anion are not very stable. Thus, for the hydrogen atom, the most characteristic is the formation of a covalent bond with other atoms due to its one unpaired electron. And in the case of the formation of an anion, and in the case of the formation of a covalent bond, the hydrogen atom is monovalent.
In a simple substance, the oxidation state of hydrogen atoms is zero, in most compounds, hydrogen exhibits an oxidation state of + I, and only in hydrides of the least electronegative elements does hydrogen have an oxidation state of –I.
Information on the valence capabilities of the hydrogen atom is given in table 28. The valence state of the hydrogen atom bound by one covalent bond to any atom is indicated in the table by the symbol "H-".

Table 28.The valence capabilities of the hydrogen atom

Valence state				Examples of chemicals
			I 0 –I	HCl, H 2 O, H 2 S, NH 3, CH 4, C 2 H 6, NH 4 Cl, H 2 SO 4, NaHCO 3, KOH H 2 B 2 H 6, SiH 4, GeH 4
				NaH, KH, CaH 2, BaH 2

c) Hydrogen molecule

The diatomic hydrogen molecule H 2 is formed when hydrogen atoms are bound by the only covalent bond possible for them. The bond is formed by the exchange mechanism. By the way the electron clouds overlap, this is s-bond (Fig.10.1 a). Since the atoms are the same, the bond is non-polar.

Interatomic distance (more precisely, the equilibrium interatomic distance, because atoms vibrate) in a hydrogen molecule r(H – H) = 0.74 A (fig.10.1 v), which is much less than the sum of the orbital radii (1.06 A). Consequently, the electron clouds of the bonded atoms overlap deeply (Fig.10.1 b), and the bond in the hydrogen molecule is strong. This is also evidenced by the rather large value of the binding energy (454 kJ / mol).
If we characterize the shape of the molecule by the boundary surface (similar to the boundary surface of the electron cloud), then we can say that the hydrogen molecule has the shape of a slightly deformed (elongated) sphere (Fig.10.1 G).

d) Hydrogen (substance)

Under normal conditions, hydrogen is a colorless and odorless gas. In not large quantities it is non-toxic. Solid hydrogen melts at 14 K (–259 ° C), and liquid hydrogen boils at 20 K (–253 ° C). Low melting and boiling points, a very small temperature range for the existence of liquid hydrogen (only 6 ° C), as well as small values ​​of molar heats of fusion (0.117 kJ / mol) and vaporization (0.903 kJ / mol) indicate that intermolecular bonds in hydrogen very weak.
The density of hydrogen r (H 2) = (2 g / mol) :( 22.4 l / mol) = 0.0893 g / l. For comparison: the average density of air is 1.29 g / l. That is, hydrogen is 14.5 times lighter than air. It is practically insoluble in water.
At room temperature hydrogen is inactive, but when heated it reacts with many substances. In these reactions, hydrogen atoms can both increase and decrease their oxidation state: Н 2 + 2 e- = 2Н -I, Н 2 - 2 e- = 2H + I.
In the first case, hydrogen is an oxidizing agent, for example, in reactions with sodium or calcium: 2Na + H 2 = 2NaH, ( t) Ca + H 2 = CaH 2. ( t)
But the reducing properties of hydrogen are more characteristic: O 2 + 2H 2 = 2H 2 O, ( t)
CuO + H 2 = Cu + H 2 O. ( t)
When heated, hydrogen is oxidized not only by oxygen, but also by some other non-metals, for example, fluorine, chlorine, sulfur, and even nitrogen.
In the laboratory, hydrogen is obtained as a result of the reaction

Zn + H 2 SO 4 = ZnSO 4 + H 2.

Iron, aluminum and some other metals can be used instead of zinc, and some other dilute acids can be used instead of sulfuric acid. The resulting hydrogen is collected in a test tube by the method of displacement of water (see Fig.10.2 b) or simply into an inverted flask (fig.10.2 a).

In industry, hydrogen is obtained in large quantities from natural gas (mainly methane) by its interaction with water vapor at 800 ° C in the presence of a nickel catalyst:

CH 4 + 2H 2 O = 4H 2 + CO 2 ( t, Ni)

or coal is treated at high temperature with water vapor:

2H 2 O + C = 2H 2 + CO 2. ( t)

Pure hydrogen is obtained from water by decomposing it electric shock(subjecting to electrolysis):

2H 2 O = 2H 2 + O 2 (electrolysis).

e) Hydrogen compounds

Hydrides (binary compounds containing hydrogen) are divided into two main types:
a) volatile (molecular) hydrides,
b) salt-like (ionic) hydrides.
Elements IVA - VIIA of groups and boron form molecular hydrides. Of these, only hydrides of elements that form non-metals are stable:

B 2 H 6; CH 4; NH 3; H 2 O; HF
SiH 4; PH 3; H 2 S; HCl
AsH 3; H 2 Se; HBr
H 2 Te; HI
With the exception of water, all these compounds at room temperature are gaseous substances, hence their name - "volatile hydrides".
Some of the elements that make up non-metals are also found in more complex hydrides. For example, carbon forms compounds with general formulas C n H 2 n+2, C n H 2 n, C n H 2 n–2 and others, where n can be very large (these compounds are studied by organic chemistry).
Ionic hydrides include hydrides of alkali, alkaline earth elements and magnesium. Crystals of these hydrides consist of H anions and metal cations in the highest oxidation state Me or Me 2 (depending on the group of the system of elements).

LiH
NaH	MgH 2
KH	CaH 2
RbH	SrH 2
CsH	BaH 2

Both ionic and almost all molecular hydrides (except for H 2 O and HF) are reducing agents, but ionic hydrides exhibit reducing properties much stronger than molecular ones.
In addition to hydrides, hydrogen is part of hydroxides and some salts. You will learn about the properties of these more complex hydrogen compounds in the following chapters.
The main consumers of hydrogen produced in the industry are ammonia plants and nitrogen fertilizers where ammonia is obtained directly from nitrogen and hydrogen:

N 2 + 3H 2 2NH 3 ( R, t, Pt - catalyst).

In large quantities, hydrogen is used to obtain methyl alcohol (methanol) by the reaction 2H 2 + CO = CH 3 OH ( t, ZnO - catalyst), as well as in the production of hydrogen chloride, which is obtained directly from chlorine and hydrogen:

H 2 + Cl 2 = 2HCl.

Sometimes hydrogen is used in metallurgy as a reducing agent in the production of pure metals, for example: Fe 2 O 3 + 3H 2 = 2Fe + 3H 2 O.

1. What particles are the nuclei of a) protium, b) deuterium, c) tritium?
2. Compare the ionization energy of the hydrogen atom with the ionization energy of the atoms of other elements. To which element is hydrogen closest to this characteristic?
3. Do the same for the electron affinity energy
4. Compare the direction of polarization of the covalent bond and the oxidation state of hydrogen in the compounds: a) BeH 2, CH 4, NH 3, H 2 O, HF; b) CH 4, SiH 4, GeH 4.
5. Write down the simplest, molecular, structural and spatial formula of hydrogen. Which one is most commonly used?
6. It is often said: "Hydrogen is lighter than air." What does this mean? When can this expression be taken literally, and when not?
7. Make the structural formulas of potassium and calcium hydrides, as well as ammonia, hydrogen sulfide and hydrogen bromide.
8. Knowing the molar heats of fusion and vaporization of hydrogen, determine the values ​​of the corresponding specific quantities.
9.For each of the four reactions that illustrate the basic chemical properties of hydrogen, draw up an electronic balance. Note oxidants and reducing agents.
10. Determine the mass of zinc required to obtain 4.48 liters of hydrogen in the laboratory.
11. Determine the mass and volume of hydrogen that can be obtained from a 30 m 3 mixture of methane and water vapor, taken in a volume ratio of 1: 2, with a yield of 80%.
12. Make the equations of the reactions occurring in the interaction of hydrogen a) with fluorine, b) with sulfur.
13. The following reaction schemes illustrate the main chemical properties of ionic hydrides:

a) MH + O 2 MOH ( t); b) MH + Cl 2 MCl + HCl ( t);
c) MH + H 2 O MOH + H 2; d) MH + HCl (p) MCl + H 2
Here M is lithium, sodium, potassium, rubidium or cesium. Write the equations of the corresponding reactions if M is sodium. Illustrate the chemical properties of calcium hydride with reaction equations.
14. Using the electronic balance method, write the equations for the following reactions to illustrate the reducing properties of some molecular hydrides:
a) HI + Cl 2 HCl + I 2 ( t); b) NH 3 + O 2 H 2 O + N 2 ( t); c) CH 4 + O 2 H 2 O + CO 2 ( t).

10.2 Oxygen

As in the case of hydrogen, the word "oxygen" is the name of both a chemical element and a simple substance. In addition to a simple substance " oxygen"(dioxygen) chemical element oxygen forms another simple substance called " ozone"(trioxygen). These are allotropic modifications of oxygen. The substance oxygen consists of oxygen molecules O 2, and the substance ozone consists of molecules of ozone O 3.

a) Chemical element oxygen

In the natural row of elements, the ordinal number of oxygen is 8. In the system of elements, oxygen is in the second period in the VIA group.
Oxygen is the most abundant element on Earth. In the earth's crust, every second atom is an oxygen atom, that is, the molar fraction of oxygen in the atmosphere, hydrosphere and lithosphere of the Earth is about 50%. Oxygen (substance) - component air. The volume fraction of oxygen in the air is 21%. Oxygen (an element) is a part of water, many minerals, as well as plants and animals. The human body contains an average of 43 kg of oxygen.
Natural oxygen consists of three isotopes (16 O, 17 O and 18 O), of which the lightest isotope 16 O is the most abundant. Therefore, the atomic mass of oxygen is close to 16 D (15.9994 D).

b) Oxygen atom

You are familiar with the following characteristics of the oxygen atom.

Table 29.Oxygen atom valence

Valence state				Examples of chemicals
				Al 2 O 3, Fe 2 O 3, Cr 2 O 3 *
			–II –I 0 + I + II	H 2 O, SO 2, SO 3, CO 2, SiO 2, H 2 SO 4, HNO 2, HClO 4, COCl 2, H 2 O 2 O 2 ** O 2 F 2 OF 2
				NaOH, KOH, Ca (OH) 2, Ba (OH) 2 Na 2 O 2, K 2 O 2, CaO 2, BaO 2
				Li 2 O, Na 2 O, MgO, CaO, BaO, FeO, La 2 O 3

* These oxides can also be considered ionic compounds.
** Oxygen atoms in a molecule are not in a given valence state; this is just an example of a substance with an oxidation state of oxygen atoms equal to zero
The high ionization energy (like hydrogen) excludes the formation of a simple cation from the oxygen atom. The electron affinity energy is quite high (almost twice that of hydrogen), which provides a greater tendency for the oxygen atom to attach electrons and the ability to form O 2A anions. But the energy of electron affinity for the oxygen atom is still less than that of halogen atoms and even other elements of group VIA. Therefore, oxygen anions ( oxide ions) exist only in compounds of oxygen with elements, the atoms of which donate electrons very easily.
By socializing two unpaired electrons, an oxygen atom can form two covalent bonds. Due to the impossibility of excitation, two lone pairs of electrons can only enter into a donor-acceptor interaction. Thus, disregarding the multiplicity of the bond and hybridization, the oxygen atom can be in one of five valence states (Table 29).
The most characteristic of the oxygen atom is the valence state with W k = 2, that is, the formation of two covalent bonds due to two unpaired electrons.
The very high electronegativity of the oxygen atom (higher - only for fluorine) leads to the fact that in most of its compounds oxygen has an oxidation state of –II. There are substances in which oxygen exhibits other values ​​of the oxidation state, some of which are shown in Table 29 as examples, and the comparative stability is shown in Fig. 10.3.

c) Oxygen molecule

It has been experimentally established that the diatomic oxygen molecule O 2 contains two unpaired electrons. Using the method of valence bonds, such an electronic structure of this molecule cannot be explained. Nevertheless, the bond in the oxygen molecule is close in properties to covalent. The oxygen molecule is non-polar. Interatomic distance ( r o – o = 1.21 A = 121 nm) is less than the distance between atoms linked by a simple bond. The molar binding energy is quite high and amounts to 498 kJ / mol.

d) Oxygen (substance)

Under normal conditions oxygen is a colorless and odorless gas. Solid oxygen melts at 55 K (–218 ° C), and liquid oxygen boils at 90 K (–183 ° C).
Intermolecular bonds in solid and liquid oxygen are somewhat stronger than in hydrogen, as evidenced by the wider temperature range for the existence of liquid oxygen (36 ° C) and higher molar heats of fusion (0.446 kJ / mol) and vaporization (6, 83 kJ / mol).
Oxygen is insignificantly soluble in water: at 0 ° C, only 5 volumes of oxygen (gas!) Dissolve in 100 volumes of water (liquid!).
The high tendency of oxygen atoms to attach electrons and high electronegativity lead to the fact that oxygen exhibits only oxidizing properties. These properties are especially pronounced at high temperatures.
Oxygen reacts with many metals: 2Ca + O 2 = 2CaO, 3Fe + 2O 2 = Fe 3 O 4 ( t);
non-metals: C + O 2 = CO 2, P 4 + 5O 2 = P 4 O 10,
and complex substances: CH 4 + 2O 2 = CO 2 + 2H 2 O, 2H 2 S + 3O 2 = 2H 2 O + 2SO 2.

Most often, as a result of such reactions, various oxides are obtained (see Chapter II § 5), but active alkali metals, for example sodium, are converted into peroxides by combustion:

2Na + O 2 = Na 2 O 2.

Structural formula of the resulting sodium peroxide (Na) 2 (O-O).
A smoldering splinter, placed in oxygen, flares up. It is a convenient and easy way to detect pure oxygen.
In industry, oxygen is obtained from air by rectification (complex distillation), and in the laboratory, by subjecting some oxygen-containing compounds to thermal decomposition, for example:
2KMnO 4 = K 2 MnO 4 + MnO 2 + O 2 (200 ° C);
2KClO 3 = 2KCl + 3O 2 (150 ° C, MnO 2 - catalyst);
2KNO 3 = 2KNO 2 + 3O 2 (400 ° C)
and, in addition, by catalytic decomposition of hydrogen peroxide at room temperature: 2H 2 O 2 = 2H 2 O + O 2 (MnO 2 -catalyst).
Pure oxygen is used in industry to intensify those processes in which oxidation occurs and to create a high-temperature flame. In rocketry, liquid oxygen is used as an oxidizer.
Oxygen is of great importance for maintaining the life of plants, animals and humans. Under normal conditions, a person has enough oxygen to breathe. But in conditions when there is not enough air, or it is absent altogether (in airplanes, during diving operations, in spaceships etc.), special gas mixtures containing oxygen. Oxygen is also used in medicine for diseases that cause difficulty in breathing.

e) Ozone and its molecules

Ozone O 3 is the second allotropic modification of oxygen.
The triatomic ozone molecule has an angular structure midway between two structures, represented by the following formulas:

Ozone is a dark blue gas with a pungent odor. Due to its strong oxidative activity, it is poisonous. Ozone is one and a half times "heavier" than oxygen and slightly more than oxygen, we will dissolve in water.
Ozone is formed in the atmosphere from oxygen during lightning electrical discharges:

3O 2 = 2O 3 ().

At normal temperatures, ozone slowly converts to oxygen, and when heated, this process proceeds with an explosion.
Ozone is contained in the so-called "ozone layer" of the earth's atmosphere, protecting all life on earth from the harmful effects of solar radiation.
In some cities, ozone is used instead of chlorine to disinfect (decontaminate) drinking water.

Draw the structural formulas of the following substances: OF 2, H 2 O, H 2 O 2, H 3 PO 4, (H 3 O) 2 SO 4, BaO, BaO 2, Ba (OH) 2. Name these substances. Describe the valence states of oxygen atoms in these compounds.
Determine the valence and oxidation state of each of the oxygen atoms.
2. Make the equations of combustion reactions in oxygen of lithium, magnesium, aluminum, silicon, red phosphorus and selenium (selenium atoms are oxidized to the oxidation state + IV, the atoms of other elements - to the highest oxidation state). What classes of oxides do the products of these reactions belong to?
3. How many liters of ozone can be obtained (under normal conditions) a) from 9 liters of oxygen, b) from 8 g of oxygen?

Water is the most abundant substance in the earth's crust. The mass of the earth's water is estimated at 10 18 tons. Water is the basis of the hydrosphere of our planet, in addition, it is contained in the atmosphere, in the form of ice forms the polar caps of the Earth and alpine glaciers, and is also part of various rocks. The mass fraction of water in the human body is about 70%.
Water is the only substance that has its own special names in all three states of aggregation.

The electronic structure of a water molecule (Fig.10.4 a) we have studied in detail earlier (see § 7.10).
Due to the polarity of the O – H bonds and the angular shape, the water molecule is electric dipole.

To characterize the polarity of an electric dipole, a physical quantity called " electric moment of an electric dipole " or simply " dipole moment ".

In chemistry, the dipole moment is measured in Debyes: 1 D = 3.34. 10-30 Cl. m

In a water molecule there are two polar covalent bonds, that is, two electric dipoles, each of which has its own dipole moment (and). The total dipole moment of the molecule is equal to the vector sum of these two moments (Fig.10.5):

(H 2 O) = ,

where q 1 and q 2 - partial charges (+) on hydrogen atoms, and and - interatomic O - H distances in the molecule. Because q 1 = q 2 = q, a, then

The experimentally determined dipole moments of the water molecule and some other molecules are given in the table.

Table 30.Dipole moments of some polar molecules

Molecule		Molecule		Molecule

Given the dipole nature of the water molecule, it is often schematically depicted as follows:
Pure water- colorless liquid, tasteless and odorless. Some of the main physical characteristics of water are given in the table.

Table 31.Some physical characteristics of water

Large values ​​of the molar heats of fusion and vaporization (an order of magnitude higher than those of hydrogen and oxygen) indicate that water molecules, both in solid and liquid matter, are quite tightly bound together. These connections are called " hydrogen bonds ".

ELECTRIC DIPOLE, DIPOLE MOMENT, BONDING POLARITY, MOLECULE POLARITY.
How many valence electrons of an oxygen atom take part in the formation of bonds in a water molecule?
2.When overlapping of which orbitals are bonds formed between hydrogen and oxygen in a water molecule?
3. Make a diagram of the formation of bonds in the hydrogen peroxide molecule H 2 O 2. What can you say about the spatial structure of this molecule?
4. The interatomic distances in HF, HCl and HBr molecules are 0.92, respectively; 1.28 and 1.41. Using the dipole moment table, calculate and compare the partial charges on the hydrogen atoms in these molecules.
5. The interatomic distances S - H in the hydrogen sulfide molecule are equal to 1.34, and the angle between the bonds is 92 °. Determine the values ​​of the partial charges on the sulfur and hydrogen atoms. What can you say about hybridization of the valence orbitals of the sulfur atom?

10.4. Hydrogen bond

As you already know, due to the significant difference in the electronegativity of hydrogen and oxygen (2.10 and 3.50), the hydrogen atom in the water molecule has a large positive partial charge ( q h = 0.33 e), and the oxygen atom has an even greater negative partial charge ( q h = -0.66 e). Recall also that the oxygen atom has two lone pairs of electrons per sp 3-hybrid AO. The hydrogen atom of one water molecule is attracted to the oxygen atom of another molecule, and, in addition, the half-empty 1s-AO of the hydrogen atom partially accepts a pair of electrons of the oxygen atom. As a result of these interactions between molecules, special kind intermolecular bonds - hydrogen bond.
In the case of water, hydrogen bonding can be schematically represented as follows:

In the last structural formula, three dots (dotted line, not electrons!) Show the hydrogen bond.

The hydrogen bond exists not only between water molecules. It is formed if two conditions are met:
1) there is a strongly polar N – E bond in the molecule (E is the symbol of an atom of a sufficiently electronegative element),
2) there is an E atom in the molecule with a large negative partial charge and a lone pair of electrons.
The element E can be fluorine, oxygen and nitrogen. Hydrogen bonds are much weaker if E is chlorine or sulfur.
Examples of substances with a hydrogen bond between molecules: hydrogen fluoride, solid or liquid ammonia, ethyl alcohol, and many others.

In liquid hydrogen fluoride, its molecules are linked by hydrogen bonds in rather long chains, and three-dimensional networks are formed in liquid and solid ammonia.
In terms of strength, a hydrogen bond is intermediate between a chemical bond and other types of intermolecular bonds. The molar energy of a hydrogen bond usually ranges from 5 to 50 kJ / mol.
In solid water (that is, ice crystals), all hydrogen atoms are hydrogen bonded to oxygen atoms, with each oxygen atom forming two hydrogen bonds (using both lone pairs of electrons). This structure makes the ice "looser" in comparison with liquid water, where some of the hydrogen bonds are broken, and the molecules are able to "pack" somewhat more densely. This feature of the structure of ice explains why, unlike most other substances, water in a solid state has a lower density than in a liquid state. Water reaches its maximum density at 4 ° C - at this temperature, a lot of hydrogen bonds break, and thermal expansion does not have a very strong effect on the density.
Hydrogen bonds are very important in our life. Let's imagine for a moment that hydrogen bonds have ceased to form. Here are some of the consequences:

• water at room temperature would become gaseous, as its boiling point would drop to about –80 ° C;
• all reservoirs would freeze from the bottom, since the density of ice would be greater than the density of liquid water;
• the double helix of DNA would cease to exist and much more.

The examples given are enough to understand that in this case, nature on our planet would become completely different.

HYDROGEN BONDING, CONDITIONS OF ITS FORMATION.
The formula of ethyl alcohol is CH 3 –CH 2 –O – H. Between which atoms of different molecules of this substance are hydrogen bonds formed? Draw up structural formulas to illustrate their formation.
2. Hydrogen bonds exist not only in individual substances, but also in solutions. Show with the help of structural formulas how hydrogen bonds are formed in an aqueous solution of a) ammonia, b) hydrogen fluoride, c) ethanol (ethyl alcohol). = 2H 2 O.
Both of these reactions occur in water constantly and at an equal rate, therefore, there is an equilibrium in water: 2H 2 O AH 3 O + OH.
This balance is called equilibrium of autoprotolysis water.

The direct reaction of this reversible process is endothermic, therefore, when heated, autoprotolysis increases, but at room temperature the equilibrium is shifted to the left, that is, the concentration of H 3 O and OH ions is negligible. What are they equal to?
According to the law of the acting masses

But due to the fact that the number of reacted water molecules in comparison with the total number of water molecules is insignificant, it can be assumed that the water concentration during autoprotolysis practically does not change, and 2 = const Such a low concentration of oppositely charged ions in pure water explains why this liquid, although poorly, still conducts an electric current.

AUTOPROTOLYSIS OF WATER, CONSTANT OF AUTOPROTOLYSIS (IONIC PRODUCT) OF WATER.
The ionic product of liquid ammonia (boiling point –33 ° C) is 2 · 10 –28. Make the equation for the auto-protolysis of ammonia. Determine the concentration of ammonium ions in pure liquid ammonia. Which of the substances has the highest electrical conductivity, water or liquid ammonia?

1. Obtaining hydrogen and its combustion (reducing properties).
2. Obtaining oxygen and combustion of substances in it (oxidizing properties).

The best known and best studied oxygen compound is its oxide H 2 O - water. Pure water is a colorless transparent liquid, odorless and tasteless. In a thick layer, it has a bluish-greenish color.

Water exists in three states of aggregation: solid - ice, liquid and gaseous - water vapor.

Of all liquid and solid substances, water has the greatest specific heat... Due to this fact, water is a heat accumulator in various organisms.

At normal pressure, the melting point of ice is 0 0 C (273 0 K), the boiling point of water is +100 0 C (373 0 K). These are abnormally high values. At T 0 +4 0 C, water has a low density equal to 1 g / ml. Above or below this temperature, the density of water is less than 1 g / ml. This feature distinguishes water from all other substances, the density of which increases with decreasing t 0. With the transition of water from their liquid state to a solid state, an increase in volume occurs: for every 92 volumes of liquid water, 100 volumes of ice are formed. With an increase in volume, the density decreases, therefore, being lighter than water, ice always floats to the surface.

Studies of the structure of water have shown that the water molecule is built like a triangle, at the top of which there is an electronegative oxygen atom, and at the corners of the bases there is hydrogen. The bond angle is 104, 27. The water molecule is polar - the electron density is shifted to the oxygen atom. Such a polar molecule can interact with another molecule to form more complex aggregates both through the interaction of dipoles and through the formation of hydrogen bonds. This phenomenon is called water association. The association of water molecules is mainly determined by the formation of hydrogen bonds between them. The molecular weight of water in a vapor state is 18 and corresponds to its simplest formula - H 2 O. In other cases, the molecular weight of water is a multiple of eighteen (18).

The polarity and small size of the molecule lead to the fact that it has strong hydrating properties.

The dielectric constant of water is so high (81) that it has a powerful ionizing effect on substances dissolved in it, causing the dissociation of acids, salts and bases.

A water molecule is able to bind to various ions to form hydrates. These compounds are characterized by a specific structure, resembling complex compounds.

One of the most important addition products is the hydronium ion - H 3 O, which is formed as a result of the addition of the H + ion to the lone pair of electrons of the oxygen atom.

As a result of this addition, the resulting hydronium ion acquires a charge of +1.

H + + H 2 O H 3 O +

Such a process is possible in systems containing substances that split off a hydrogen ion.

Water, both in the cold and when heated, actively interacts with many metals, standing in the line of activity up to hydrogen. In these reactions, the corresponding oxides or hydroxides are formed and hydrogen is displaced:

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

2 Na + 2 HOH = 2 NaOH + H 2

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

Water quite actively joins the main and acid oxides, forming the corresponding hydroxides:

CaO + H 2 O = Ca (OH) 2 - base

P 2 O 5 + 3 H 2 O = 2 H 3 PO 4 - acid

Water, which is attached in these cases, is called constitutional (as opposed to crystallization in crystalline hydrates).

Water reacts with halogens, in this case a mixture of acids is formed:

H 2 + HOH HCl + HClO

The most important property of water is its dissolving power.

Water is the most common solvent in nature and technology. Most chemical reactions take place in water. But perhaps the most important are biological and biochemical processes occurring in plant and animal organisms with the participation of proteins, fats, carbohydrates and other substances in aquatic environment organism.

The second compound of hydrogen with oxygen is hydrogen peroxide H 2 O 2.

Structural formula H - O - O - H, molecular weight - 34.

Latin name Hydrogenii peroxydum.

This substance was discovered in 1818 by the French scientist Louis-Jacques Thénard, who studied the effect of various mineral acids on barium peroxide (BaO 2). In nature, hydrogen peroxide is formed during oxidation. The most convenient and in a modern way obtaining H 2 O 2 is an electrolytic method, which is used in industry. Sulfuric acid or ammonium sulfate are used as starting materials.

It has been established by modern physicochemical methods that both oxygen atoms in hydrogen peroxide are linked directly to each other by a non-polar covalent bond. the bonds between hydrogen and oxygen atoms (due to the displacement of common electrons towards oxygen) are polar. Therefore, the H 2 O 2 molecule is also polar. A hydrogen bond arises between the H 2 O 2 molecules, which leads to their association with the O - O bond energy of 210 kJ, which is significantly less than the H - O bond energy (470 kJ).

Hydrogen peroxide solution- a clear, colorless liquid, odorless or with a faint peculiar smell, slightly acidic reaction. It decomposes quickly on exposure to light, on heating, on contact with alkali, oxidizing and reducing substances, releasing oxygen. The reaction occurs: H 2 O 2 = H 2 O + O

The low stability of H 2 O 2 molecules is due to the fragility of the O - O bond.

Store it in a dark glass dish and in a cool place. When concentrated solutions of hydrogen peroxide act on the skin, burns are formed, and the burned area hurts.

APPLICATION: in medicine, a 3% solution of hydrogen peroxide is used as a hemostatic agent, disinfectant and deodorizing agent for rinsing and rinsing for stomatitis, sore throat, gynecological diseases, etc.

When in contact with the enzyme catalase (from blood, pus, tissues), atomic oxygen acts at the time of release. The action of H 2 O 2 is short-term. The value of the drug lies in the fact that its decomposition products are harmless to tissues.

HYDROPERIT is a complex compound of hydrogen peroxide with urea. The hydrogen peroxide content is about 35%. Apply as antiseptic instead of hydrogen peroxide.

One of the main chemical properties of H 2 O 2 is its redox properties. The oxidation state of oxygen in H 2 O 2 is -1, i.e. has an intermediate value between the oxidation state of oxygen in water (-2) and in molecular oxygen (0). Therefore, hydrogen peroxide has the properties of both an oxidizing agent and a reducing agent, i.e. exhibits redox duality. It should be noted that the oxidizing properties of H 2 O 2 are much more pronounced than the reducing ones and they are manifested in acidic, alkaline and neutral media. For example:

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

2 I - - 2ē → I 2 0 1 - v-l

H 2 O 2 + 2 H + + 2ē → 2 H 2 O 1 - ok-l

2 I - + H 2 O 2 + 2 H + → I 2 + 2 H 2 O

Under the influence of strong oxidants, H 2 O 2 exhibits reducing properties:

2 KMnO 4 + 5 H 2 O 2 + 3 H 2 SO 4 = 2 MnSO 4 + 5 O 2 + K 2 SO 4 + 8 H 2 O

MnO 4 - + 8H + + 5ē → Mn +2 + 4 H 2 O 2 - ok-l

H 2 O 2 - 2ē → O 2 + 2 H + 5 - v-l

2 MnO 4 - + 5 H 2 O 2 + 16 H + → 2 Mn +2 + 8 H 2 O + 5 O 2 + 10 H +

Conclusions:

1. Oxygen is the most abundant element on Earth.

In nature, oxygen occurs in two allotropic modifications: O 2 - dioxygen or "ordinary oxygen" and O 3 - trioxide (ozone).

2.Allotropy- the formation of different simple substances by one element.

3. Allotropic modifications of oxygen: oxygen and ozone.

4. Compounds of oxygen with hydrogen - water and hydrogen peroxide .

5. Water exists in three states of aggregation: in solid - ice, liquid and gaseous - water vapor.

6. At T 0 +4 0 C, water has a density equal to 1 g / ml.

7. The water molecule is built in the form of a triangle, at the apex of which there is an electronegative oxygen atom, and at the corners of the bases there is hydrogen.

8. The bond angle is 104, 27

9. The water molecule is polar - the electron density is shifted towards the oxygen atom.

12. Sulfur. Characterization of sulfur, based on its position in the periodic system, from the point of view of the theory of atomic structure, possible oxidation states, physical properties, distribution in nature, biological role, production methods, chemical properties. ... The use of sulfur and its compounds in medicine and the national economy.

SULFUR:

A) being in nature

B) biological role

C) use in medicine

Sulfur is widespread in nature and occurs both in a free state (native sulfur) and in the form of compounds - FeSe (pyrite), CuS, Ag 2 S, PbS, CaSO 4, etc. different connections contained in natural coals, oils and natural gases.

Sulfur is one of the elements that are important for life processes, because it is part of protein substances. The sulfur content in the human body is 0.25%. It is part of the amino acids: cysteine, glutathione, methionine, etc.

Especially a lot of sulfur is in the proteins of hair, horns, wool. In addition, sulfur is part of biologically active substances of the body: vitamins and hormones (eg, insulin).

Sulfur is found in the form of compounds in nervous tissue, cartilage, bones and bile. She participates in the redox processes of the body.

With a lack of sulfur in the body, there is fragility and fragility of bones, hair loss.

Sulfur is found in gooseberries, grapes, apples, cabbage, onions, rye, peas, barley, buckwheat, and wheat.

Record holders: 190 peas, 244% soy.

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 weight. In the earth's crust, 58% of the atoms are bound oxygen atoms (47% by mass). Oxygen is a part of water (the reserves of bound oxygen in the hydrosphere are extremely large), rocks, many minerals and salts, it is contained in fats, proteins and carbohydrates that make up living organisms. Almost all of the Earth's free oxygen arose and is preserved as a result of the process of photosynthesis.

Physical properties.

Oxygen is a colorless, tasteless and odorless gas, slightly heavier than air. We are slightly soluble in water (31 ml of oxygen dissolves in 1 liter of water at 20 degrees), but still better than other gases of the atmosphere, therefore, water is enriched with oxygen. The density of oxygen 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- ability chemical element exist in the form of two or more simple substances that differ only in the number of atoms in a molecule, or in structure.

Ozone O 3 - exists in upper layers 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 destructive ultraviolet radiation The sun; pale purple, poisonous in large quantities gas with a specific, pungent, but pleasant odor. The melting point is -192.7 0 С, the boiling point is -111.9 0 С. In water we will dissolve better than oxygen.

Ozone - 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. Forms ozonides with some metals, for example potassium ozonide:

K + O 3 = KO 3

Ozone is obtained in special devices - ozonizers. In them, under the action 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 for bleaching fabrics, disinfecting drinking water, in medicine as a disinfectant.

Inhalation of large quantities of ozone 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:

2Са + О 2 = 2СаО

2Ва + О 2 = 2ВаО

2Na + O 2 = Na 2 O 2

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

The interaction of oxygen with nitrogen begins only at a temperature of 1200 0 C or in an electric 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 oxygen explodes when ignited; it is called detonating gas.

Many metals, when in 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

Complex substances with certain conditions 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. The oxidative activity of oxygen is based on its important property - the ability to maintain combustion substances.

With hydrogen, oxygen also forms a compound - hydrogen peroxide Н 2 О 2 - a colorless transparent liquid with a burning astringent taste, readily 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, contact with oxidizing agents 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); therefore, hydrogen peroxide exhibits redox duality. The oxidizing properties of hydrogen peroxide are much more pronounced than the reducing ones, and they are manifested in acidic, alkaline and neutral media.

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