§3. Reaction equation and how to write it

Interaction hydrogen With oxygen, as Sir Henry Cavendish established, leads to the formation of water. Let's get on with it simple example let's learn how to compose chemical reaction equations.
What comes out of hydrogen And oxygen, we already know:

H 2 + O 2 → H 2 O

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

2H 2 + O 2 = 2H 2 O

This is it reaction equation - conditional recording of an ongoing chemical reaction using formulas of substances and coefficients.

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

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

When water is formed from hydrogen and oxygen, it means that hydrogen changed its oxidation state from 0 to +I, A oxygen- from 0 to −II. In this case, several passed from hydrogen atoms to oxygen atoms. (n) electrons:

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

Oxidizing agents and reducing agents

Let's now see what the processes of giving and receiving electrons look like separately. Hydrogen, having met with the “robber” oxygen, loses all its assets - two electrons, and its oxidation state becomes equal +I:

N 2 0 − 2 e− = 2Н +I

It worked oxidation half-reaction equation hydrogen.

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

O2+4 e− = 2O −II

This reduction half-reaction equation oxygen.

It remains to add that both the “bandit” and his “victim” have lost their chemical individuality and are made from simple substances - gases with diatomic molecules H 2 And O 2 turned into components of a new chemical substance - water H 2 O.

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

So it's necessary equalize the number of electrons in the first and second half-reactions. In chemistry, the following conventional form of writing half-reaction equations is accepted:

	2 N 2 0 − 2 e− = 2Н +I
	1 O 2 0 + 4 e− = 2O −II

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

2H 2 0 + O 2 0 = 2H 2 +I O −II

Hydrogen oxidizes not only when meeting with oxygen. They act on hydrogen in approximately the same way. fluorine F 2, a halogen and a known "robber", and seemingly harmless nitrogen N 2:

H 2 0 + F 2 0 = 2H +I F −I

3H 2 0 + N 2 0 = 2N −III H 3 +I

In this case it turns out hydrogen fluoride HF or ammonia NH 3.

In both compounds the oxidation state is hydrogen becomes equal +I, because he gets molecule partners who are “greedy” for other people’s electronic goods, with high electronegativity - fluorine F And nitrogen N. U nitrogen the electronegativity value is considered equal to three conventional units, and fluoride In general, the highest electronegativity among all chemical elements is four units. So it’s no wonder they left the poor hydrogen atom without any electronic environment.

But hydrogen maybe restore- accept electrons. This happens if alkali metals or calcium, which have a lower electronegativity than hydrogen, participate in the reaction with it.

Hydrogen has a special position in Periodic table chemical elements D.I. Mendeleev. In terms of the number of valence electrons and the ability to form the hydration ion H + in solutions, it is similar to alkali metals, and should be placed in group I. According to the number of electrons required to complete the outer electron shell, the value of ionization energy, the ability to exhibit a negative oxidation state, small atomic radius Hydrogen should be placed in group VII of the periodic table. Thus, the placement of hydrogen in one group or another of the periodic table is largely arbitrary, but in most cases it is placed in group VII.

Electronic formula of hydrogen 1 s 1. The only valence electron is directly in the range of action atomic nucleus. Simplicity electronic configuration hydrogen does not mean that chemical properties of this element are simple. In contrast, the chemistry of hydrogen differs in many ways from the chemistry of other elements. Hydrogen in its compounds is capable of exhibiting oxidation states of +1 and –1.

There are a large number of methods for producing hydrogen. In the laboratory it is obtained by reacting certain metals with acids, for example:

Hydrogen can be obtained by electrolysis of aqueous solutions of sulfuric acid or alkalis. In this case, the process of hydrogen evolution occurs at the cathode and oxygen at the anode.

In industry, hydrogen is produced mainly from natural and associated gases, fuel gasification products and coke oven gas.

Simple substance hydrogen (H 2) is a flammable gas, colorless and odorless. Boiling point –252.8 °C. Hydrogen is 14.5 times lighter than air and slightly soluble in water.

The hydrogen molecule is stable and has great strength. Due to the high dissociation energy (435 kJ/mol), the decomposition of H 2 molecules into atoms occurs to a noticeable extent only at temperatures above 2000 °C.

Positive and negative oxidation states are possible for hydrogen, so in chemical reactions hydrogen can exhibit both oxidizing and reducing properties. In cases where hydrogen acts as an oxidizing agent, it behaves like halogens, forming hydrides similar to halides ( hydrides name a group of chemical compounds of hydrogen with metals and elements less electronegative than it):

In terms of oxidative activity, hydrogen is significantly inferior to halogens. Therefore, only hydrides of alkali and alkaline earth metals exhibit ionic character. Ionic as well as complex hydrides, for example, are strong reducing agents. They are widely used in chemical syntheses.

In most reactions, hydrogen behaves as a reducing agent. Under normal conditions, hydrogen does not react with oxygen, but when ignited, the reaction occurs explosively:

A mixture of two volumes of hydrogen with one volume of oxygen is called detonating gas. During controlled combustion, release occurs large quantity heat, and the temperature of the hydrogen-oxygen flame reaches 3000 °C.

The reaction with halogens proceeds in different ways, depending on the nature of the halogen:

With fluorine, this reaction occurs explosively even with low temperatures. With chlorine in the light, the reaction also occurs explosively. With bromine the reaction is much slower, and with iodine it does not reach completion even with high temperature. The mechanism of these reactions is radical.

At elevated temperatures, hydrogen interacts with elements of group VI - sulfur, selenium, tellurium, for example:

The reaction of hydrogen with nitrogen is very important. This reaction is reversible. To shift the equilibrium towards the formation of ammonia, use high blood pressure. In industry, this process is carried out at a temperature of 450–500 °C in the presence of various catalysts:

Hydrogen reduces many metals from oxides, for example:

This reaction is used to obtain some pure metals.

Hydrogenation reactions play a huge role organic compounds, which are widely used both in laboratory practice and in industrial organic synthesis.

Reduction natural sources hydrocarbons, pollution environment fuel combustion products are increasing interest in hydrogen as an environmentally friendly fuel. Hydrogen is likely to play an important role in the energy sector of the future.

Currently, hydrogen is widely used in industry for the synthesis of ammonia, methanol, hydrogenation of solid and liquid fuels, in organic synthesis, for welding and cutting metals, etc.

Water H 2 O, hydrogen oxide, is the most important chemical compound. Under normal conditions, water is a colorless liquid, odorless and tasteless. Water is the most abundant substance on the surface of the Earth. IN human body contains 63–68% water.

The physical properties of water are in many ways anomalous. At normal atmospheric pressure, water boils at 100 °C. The freezing point of pure water is 0 °C. Unlike other liquids, the density of water does not increase monotonically when cooled, but has a maximum at +4 °C. The heat capacity of water is very high and amounts to 418 kJ/mol·K. The heat capacity of ice at 0 °C is 2.038 kJ/mol·K. The heat of melting of ice is abnormally high. The electrical conductivity of water is very low. Abnormal physical properties waters explain its structure. The H–O–H bond angle is 104.5°. The water molecule is a distorted tetrahedron, at two vertices of which hydrogen atoms are located, and the other two are occupied by the orbitals of lone pairs of electrons of the oxygen atom, which are not involved in the formation of chemical bonds.

Water is a stable compound; its decomposition into oxygen and hydrogen occurs only under the influence of direct electric current or at a temperature of about 2000 °C:

Water directly interacts with metals in the range of standard electronic potentials up to hydrogen. Depending on the nature of the metal, the reaction products can be the corresponding hydroxides and oxides. The reaction rate, depending on the nature of the metal, also varies within wide limits. Thus, sodium reacts with water already at room temperature, the reaction is accompanied by the release of a large amount of heat; iron reacts with water at a temperature of 800 °C:

Hydrogen is a gas; it is in first place in the Periodic Table. The name of this element, widespread in nature, is translated from Latin as “generating water.” So what physical and chemical properties of hydrogen do we know?

Hydrogen: general information

At normal conditions Hydrogen has no taste, no smell, no color.

Rice. 1. Formula of hydrogen.

Since an atom has one electronic energy level, which can contain a maximum of two electrons, then for a stable state the atom can either accept one electron (oxidation state -1) or give up one electron (oxidation state +1), exhibiting a constant valence I This is why the symbol of the element hydrogen is placed not only in group IA (the main subgroup of group I) together with the alkali metals, but also in group VIIA (the main subgroup of group VII) together with the halogens. Halogen atoms also lack one electron to fill the outer level, and they, like hydrogen, are nonmetals. Hydrogen exhibits a positive oxidation state in compounds where it is associated with more electronegative nonmetal elements, and a negative oxidation state in compounds with metals.

Rice. 2. The location of hydrogen in the periodic table.

Hydrogen has three isotopes, each of which has its own name: protium, deuterium, tritium. The amount of the latter on Earth is negligible.

Chemical properties of hydrogen

In the simple substance H2, the bond between the atoms is strong (bond energy 436 kJ/mol), therefore the activity of molecular hydrogen is low. Under normal conditions, it reacts only with very reactive metals, and the only non-metal with which hydrogen reacts is fluorine:

F 2 +H 2 =2HF (hydrogen fluoride)

Hydrogen reacts with other simple (metals and non-metals) and complex (oxides, unspecified organic compounds) substances either upon irradiation and increased temperature, or in the presence of a catalyst.

Hydrogen burns in oxygen, releasing a significant amount of heat:

2H 2 +O 2 =2H 2 O

A mixture of hydrogen and oxygen (2 volumes of hydrogen and 1 volume of oxygen) explodes violently when ignited and is therefore called detonating gas. When working with hydrogen, safety regulations must be followed.

Rice. 3. Explosive gas.

In the presence of catalysts, the gas can react with nitrogen:

3H 2 +N 2 =2NH 3

– this reaction at elevated temperatures and pressures produces ammonia in industry.

At high temperatures, hydrogen is able to react with sulfur, selenium, and tellurium. and when interacting with alkali and alkaline earth metals, the formation of hydrides occurs:

- V in this case hydrogen plays the role of an oxidizing agent.

Hydrogen has the ability to reduce the oxides of many metals when the temperature increases, resulting in the formation of water. For example:

CuO+H 2 =H 2 O+Cu

- V this process hydrogen is a reducing agent4.3. Total ratings received: 70.

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

Hydrogen is one of the most common elements on Earth. The mole fraction of hydrogen atoms in the atmosphere, hydrosphere and lithosphere of the Earth (collectively called the earth's crust) is 0.17. It is found in water, many minerals, oil, natural gas, plants and animals. The average human body contains 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, so it is practically never found in nature. Deuterium is stable, but there is 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 Dn (1.00794 Dn).

b) Hydrogen atom

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

The valence capabilities of a hydrogen atom are determined by the presence of one electron in a single valence orbital. A high ionization energy makes a hydrogen atom not inclined to give up an electron, and a not too high electron affinity energy leads to a slight tendency to accept one. Consequently, in chemical systems the formation of the H cation is impossible, and compounds with the H anion are not very stable. Thus, the hydrogen atom is most likely to form a covalent bond with other atoms due to its one unpaired electron. Both 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 the hydrides of the least electronegative elements does hydrogen have an oxidation state of –I.
Information about the valence capabilities of the hydrogen atom is given in Table 28. The valence state of a hydrogen atom bound by one covalent bond to any atom is indicated in the table by the symbol “H-”.

Table 28.Valence possibilities 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 H2 is formed when hydrogen atoms are bonded with the only covalent bond possible for them. The connection is formed by an exchange mechanism. According to the way electron clouds overlap, this is an s-bond (Fig. 10.1 A). Since the atoms are the same, the bond is non-polar.

Interatomic distance (more precisely, equilibrium interatomic distance, because atoms vibrate) in a hydrogen molecule r(H–H) = 0.74 A (Fig. 10.1 V), which is significantly less than the sum of the orbital radii (1.06 A). Consequently, the electron clouds of bonded atoms overlap deeply (Fig. 10.1 b), and the bond in the hydrogen molecule is strong. This is pretty much the same thing great value 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) ball (Fig. 10.1 G).

d) Hydrogen (substance)

Under normal conditions, hydrogen is a colorless and odorless gas. In small 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 the molar heats of fusion (0.117 kJ/mol) and vaporization (0.903 kJ/mol) indicate that intermolecular bonds in hydrogen very weak.
Hydrogen density r(H 2) = (2 g/mol): (22.4 l/mol) = 0.0893 g/l. For comparison: the average air density 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 either increase or decrease their oxidation state: H 2 + 2 e– = 2Н –I, Н 2 – 2 e– = 2Н +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 produced as a result of the reaction

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

Instead of zinc, you can use iron, aluminum and some other metals, and instead of sulfuric acid, you can use some other dilute acids. The resulting hydrogen is collected in a test tube by displacing water (see Fig. 10.2 b) or simply into an inverted flask (Fig. 10.2 A).

In industry, hydrogen is produced in large quantities from natural gas (mainly methane) by reacting it 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 treat coal 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 of groups IVA – VIIA and boron form molecular hydrides. Of these, only the hydrides of elements forming nonmetals are stable:

B 2 H 6 ; CH 4 ; NH3; H2O; HF
SiH 4 ;PH 3 ; H2S; HCl
AsH3; H2Se; HBr
H2Te; HI
With the exception of water, all these compounds are gaseous substances at room temperature, hence their name - “volatile hydrides”.
Some of the elements that form nonmetals are also found in more complex hydrides. For example, carbon forms compounds with the 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 in organic chemistry).
Ionic hydrides include hydrides of alkali, alkaline earth elements and magnesium. The 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 element system).

LiH
NaH	MgH 2
KH	CaH2
RbH	SrH 2
CsH	BaH 2

Both ionic and almost all molecular hydrides (except 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 become familiar with the properties of these more complex hydrogen compounds in the following chapters.
The main consumers of hydrogen produced in industry are plants for the production of ammonia and nitrogen fertilizers, where ammonia is obtained directly from nitrogen and hydrogen:

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

Hydrogen is used in large quantities to produce 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 do the nuclei of a) protium, b) deuterium, c) tritium consist of?
2.Compare the ionization energy of the hydrogen atom with the ionization energy of atoms of other elements. Which element is hydrogen closest to in terms of this characteristic?
3.Do the same for electron affinity energy
4. Compare the direction of polarization of the covalent bond and the degree of oxidation 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 often used?
6. They often say: “Hydrogen is lighter than air.” What does this mean? In what cases can this expression be taken literally, and in what cases can it not?
7.Make up the structural formulas of potassium and calcium hydrides, as well as ammonia, hydrogen sulfide and hydrogen bromide.
8.Knowing the molar heats of melting and vaporization of hydrogen, determine the values ​​of the corresponding specific quantities.
9.For each of the four reactions illustrating the basic chemical properties of hydrogen, create an electronic balance. Label the oxidizing and reducing agents.
10. Determine the mass of zinc required to produce 4.48 liters of hydrogen using a laboratory method.
11. Determine the mass and volume of hydrogen that can be obtained from 30 m 3 of a mixture of methane and water vapor, taken in a volume ratio of 1:2, with a yield of 80%.
12. Make up equations for the reactions that occur during the interaction of hydrogen a) with fluorine, b) with sulfur.
13.The reaction schemes below illustrate the basic 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 down the equations for the corresponding reactions if M is sodium. Illustrate the chemical properties of calcium hydride using reaction equations.
14.Using the electron balance method, create equations for the following reactions illustrating 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 with hydrogen, the word "oxygen" is the name of both a chemical element and simple substance. Apart from simple matter" oxygen"(dioxygen) chemical element oxygen forms another simple substance called " ozone"(trioxygen). This allotropic modifications oxygen. The substance oxygen consists of oxygen molecules O 2 , and the substance ozone consists of ozone molecules O 3 .

a) Chemical element oxygen

In the natural series of elements, the serial 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 earth's crust every second atom is an oxygen atom, that is, the mole 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 found in water, many minerals, and plants and animals. The human body contains on average 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 common. Therefore, the atomic mass of oxygen is close to 16 Dn (15.9994 Dn).

b) Oxygen atom

You know the following characteristics of the oxygen atom.

Table 29.Valence possibilities of the oxygen atom

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 O2** O2F2 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 as ionic compounds.
** The oxygen atoms in the molecule are not in this 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 that of hydrogen) prevents 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 propensity for the oxygen atom to gain electrons and the ability to form O 2A anions. But the electron affinity energy of the oxygen atom is still lower than that of halogen atoms and even other elements of the VIA group. Therefore, oxygen anions ( oxide ions) exist only in compounds of oxygen with elements whose atoms give up electrons very easily.
By sharing two unpaired electrons, an oxygen atom can form two covalent bonds. Two lone pairs of electrons, due to the impossibility of excitation, can only enter into donor-acceptor interaction. Thus, without taking into account the bond multiplicity and hybridization, the oxygen atom can be in one of five valence states (Table 29).
The most typical valence state for the oxygen atom is 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 oxidation states, some of them are given 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 valence bond method, this electronic structure of this molecule cannot be explained. However, the properties of the bond in the oxygen molecule are close to that of a covalent bond. The oxygen molecule is non-polar. Interatomic distance ( r o–o = 1.21 A = 121 nm) is less than the distance between atoms connected by a single 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 larger temperature range of existence of liquid oxygen (36 °C) and larger molar heats of fusion (0.446 kJ/mol) and vaporization (6. 83 kJ/mol).
Oxygen is slightly soluble in water: at 0 °C, only 5 volumes of oxygen (gas!) dissolve in 100 volumes of water (liquid!).
The high propensity of oxygen atoms to gain 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, when burned, turn into peroxides:

2Na + O 2 = Na 2 O 2.

The structural formula of the resulting sodium peroxide is (Na) 2 (O-O).
A smoldering splinter placed in oxygen bursts into flames. This 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 certain 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 the 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 rocket technology, 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 in the air to breathe. But in conditions where there is not enough air, or there is no air at all (in airplanes, during diving work, in spaceships, etc.), special ones are prepared for breathing gas mixtures containing oxygen. Oxygen is also used in medicine for diseases that cause difficulty breathing.

e) Ozone and its molecules

Ozone O 3 is the second allotropic modification of oxygen.
The triatomic ozone molecule has a corner structure intermediate between the two structures represented by the following formulas:

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

3O 2 = 2O 3 ().

At normal temperatures, ozone slowly turns into oxygen, and when heated, this process occurs explosively.
Ozone is contained in the so-called "ozone layer" earth's atmosphere, protecting all life on Earth from harmful effects solar radiation.
In some cities, ozone is used instead of chlorine to disinfect (disinfect) 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 oxygen atom.
2. Make up equations for the combustion reactions of lithium, magnesium, aluminum, silicon, red phosphorus and selenium in oxygen (selenium atoms are oxidized to the oxidation state +IV, atoms of other elements are oxidized 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 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 it forms the Earth’s polar caps and high-mountain 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 all three states of aggregation have their own special names.

Electronic structure of a water molecule (Fig. 10.4 A) we 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 just " dipole moment".

In chemistry, the dipole moment is measured in debyes: 1 D = 3.34. 10 –30 Class. m

In a water molecule there are two polar covalent bonds, that is, two electric dipoles, each of which has its own dipole moment (u). The total dipole moment of a 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, and , 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 represented as follows:
Pure water is a colorless liquid without taste or smell. Some basic physical characteristics of water are given in the table.

Table 31.Some physical characteristics of water

The large values ​​of the molar heats of melting and vaporization (an order of magnitude greater 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, BOND POLARITY, MOLECULE POLARITY.
How many valence electrons of an oxygen atom take part in the formation of bonds in a water molecule?
2. When what orbitals overlap, bonds are formed between hydrogen and oxygen in a water molecule?
3.Make a diagram of the formation of bonds in a molecule of hydrogen peroxide H 2 O 2. What can you say about the spatial structure of this molecule?
4. Interatomic distances in HF, HCl and HBr molecules are equal to 0.92, respectively; 1.28 and 1.41. Using the table of dipole moments, 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 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 the hybridization of the valence orbitals of the sulfur atom?

10.4. Hydrogen bond

As you already know, due to the significant difference in electronegativity of hydrogen and oxygen (2.10 and 3.50), the hydrogen atom in the water molecule acquires 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, a special kind intermolecular bonds - hydrogen bond.
In the case of water, hydrogen bond formation can be represented schematically as follows:

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

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

In liquid hydrogen fluoride, its molecules are linked by hydrogen bonds into rather long chains, and in liquid and solid ammonia three-dimensional networks are formed.
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 (i.e., 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 ice more “loose” compared to liquid water, where some of the hydrogen bonds are broken, and the molecules are able to “pack” a little more tightly. This feature of the structure of ice explains why, unlike most other substances, water in the solid state has a lower density than in the liquid state. Water reaches its maximum density at 4 °C - at this temperature quite a lot of hydrogen bonds are broken, and thermal expansion does not yet have a very strong effect on the density.
Hydrogen bonds are very important in our lives. Let's imagine for a moment that hydrogen bonds have stopped forming. Here are some consequences:

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

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

HYDROGEN BOND, 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? Write structural formulas illustrating their formation.
2. Hydrogen bonds exist not only in individual substances, but also in solutions. Show using structural formulas how hydrogen bonds are formed in aqueous solution a) ammonia, b) hydrogen fluoride, c) ethanol (ethyl alcohol). = 2H 2 O.
Both of these reactions occur in water constantly and at the same speed, therefore, there is an equilibrium in water: 2H 2 O AN 3 O + OH.
This equilibrium 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 mass action

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

AUTOPROTOLYSIS OF WATER, AUTOPROTOLYSIS CONSTANT (IONIC PRODUCT) OF WATER.
The ionic product of liquid ammonia (boiling point –33 °C) is 2·10 –28. Write an equation for the autoprotolysis of ammonia. Determine the concentration of ammonium ions in pure liquid ammonia. Which substance has greater electrical conductivity, water or liquid ammonia?

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

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:

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

- When hydrogen interacts with non-metals volatile hydrogen compounds are formed. IN chemical formula volatile hydrogen compound, the hydrogen atom can be in either the first or second place, depending on its location in the PSCE (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 synthesis 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: they are usually used to fill balloons, chemical industry it serves as a raw material for the production of many very important products (ammonia, etc.), in food - for the production of vegetable oils solid fats, etc. The high temperature (up to 2600 °C) resulting from the combustion of hydrogen in oxygen is used to melt 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?