Melting point of hydrogen. Chemical properties of hydrogen: features and applications

Hydrogen was discovered in the second half of the 18th century by the English scientist in the field of physics and chemistry G. Cavendish. He managed to isolate the substance in its pure state, began studying it and described its properties.

This is the story of the discovery of hydrogen. During the experiments, the researcher determined that it is a flammable gas, the combustion of which in the air produces water. This led to the definition quality composition water.

What is hydrogen

The French chemist A. Lavoisier first announced hydrogen as a simple substance in 1784, since he determined that its molecule contains atoms of the same type.

The name of the chemical element in Latin sounds like hydrogenium (read “hydrogenium”), which means “water-giving.” The name refers to the combustion reaction that produces water.

Characteristics of hydrogen

Designation of hydrogen N. Mendeleev assigned the first atomic number to this chemical element, placing it in the main subgroup of the first group and the first period and conditionally in the main subgroup of the seventh group.

Atomic weight ( atomic mass) of hydrogen is 1.00797. The molecular weight of H2 is 2 a. e. Molar mass numerically equal to it.

It is represented by three isotopes that have a special name: the most common protium (H), heavy deuterium (D), radioactive tritium (T).

It is the first element that can be completely separated into isotopes in a simple way. It is based on the high difference in mass of isotopes. The process was first carried out in 1933. This is explained by the fact that only in 1932 an isotope with mass 2 was discovered.

Physical properties

IN normal conditions The simple substance hydrogen in the form of diatomic molecules is a gas, colorless, tasteless and odorless. Slightly soluble in water and other solvents.

Crystallization temperature - 259.2 o C, boiling point - 252.8 o C. The diameter of hydrogen molecules is so small that they have the ability to slowly diffuse through a number of materials (rubber, glass, metals). This property is used when it is necessary to purify hydrogen from gaseous impurities. When n. u. hydrogen has a density of 0.09 kg/m3.

Is it possible to transform hydrogen into a metal by analogy with the elements located in the first group? Scientists have found that hydrogen, under conditions when the pressure approaches 2 million atmospheres, begins to absorb infrared rays, which indicates the polarization of the molecules of the substance. Perhaps with even more high pressures, hydrogen will become a metal.

This is interesting: there is an assumption that on the giant planets, Jupiter and Saturn, hydrogen is found in the form of a metal. It is assumed that metallic solid hydrogen is also present in the earth's core, due to the ultra-high pressure created by the earth's mantle.

Chemical properties

Both simple and complex substances enter into chemical interaction with hydrogen. But the low activity of hydrogen needs to be increased by creating appropriate conditions - increasing the temperature, using catalysts, etc.

When heated, simple substances such as oxygen (O 2), chlorine (Cl 2), nitrogen (N 2), sulfur (S) react with hydrogen.

If you ignite pure hydrogen at the end of a gas outlet tube in air, it will burn evenly, but barely noticeably. If you place the gas outlet tube in an atmosphere of pure oxygen, then combustion will continue with the formation of water droplets on the walls of the vessel, as a result of the reaction:

The combustion of water is accompanied by the release of a large amount of heat. It is an exothermic compound reaction in which hydrogen is oxidized by oxygen to form the oxide H 2 O. It is also a redox reaction in which hydrogen is oxidized and oxygen is reduced.

The reaction with Cl 2 occurs similarly to form hydrogen chloride.

The interaction of nitrogen with hydrogen requires high temperature and high pressure, as well as the presence of a catalyst. The result is ammonia.

As a result of the reaction with sulfur, hydrogen sulfide is formed, the recognition of which is facilitated by the characteristic smell of rotten eggs.

The oxidation state of hydrogen in these reactions is +1, and in the hydrides described below – 1.

When reacting with some metals, hydrides are formed, for example, sodium hydride - NaH. Some of these complex compounds are used as fuel for rockets and also in thermonuclear power.

Hydrogen also reacts with substances from the complex category. For example, with copper (II) oxide, formula CuO. To carry out the reaction, copper hydrogen is passed over heated powdered copper (II) oxide. During the interaction, the reagent changes its color and becomes red-brown, and droplets of water settle on the cold walls of the test tube.

Hydrogen is oxidized during the reaction, forming water, and copper is reduced from oxide to a simple substance (Cu).

Areas of use

Hydrogen has great importance for humans and is used in a variety of areas:

1. In chemical production it is raw materials, in other industries it is fuel. Petrochemical and oil refining enterprises cannot do without hydrogen.
2. In the electric power industry, this simple substance acts as a cooling agent.
3. In ferrous and non-ferrous metallurgy, hydrogen plays the role of a reducing agent.
4. This helps create an inert environment when packaging products.
5. Pharmaceutical industry - uses hydrogen as a reagent in the production of hydrogen peroxide.
6. Weather balloons are filled with this light gas.
7. This element is also known as a fuel reducer for rocket engines.

Scientists unanimously predict that hydrogen fuel will take the lead in the energy sector.

Receipt in industry

In industry, hydrogen is produced by electrolysis, which is subjected to chlorides or hydroxides of alkali metals dissolved in water. It is also possible to obtain hydrogen directly from water using this method.

The conversion of coke or methane with water vapor is used for these purposes. The decomposition of methane at elevated temperatures also produces hydrogen. Liquefaction of coke oven gas by fractional method is also used for industrial production hydrogen.

Obtained in the laboratory

In the laboratory, a Kipp apparatus is used to produce hydrogen.

The reagents are hydrochloric or sulfuric acid and zinc. The reaction produces hydrogen.

Finding hydrogen in nature

Hydrogen is more common than any other element in the Universe. The bulk of stars, including the Sun, and other cosmic bodies is hydrogen.

IN earth's crust it is only 0.15%. It is present in many minerals, in all organic substances, as well as in water, which covers 3/4 of the surface of our planet.

In the upper atmosphere, traces of hydrogen can be found in pure form. It is also found in a number of flammable natural gases.

Gaseous hydrogen is the least dense, and liquid hydrogen is the densest substance on our planet. With the help of hydrogen, you can change the timbre of your voice if you inhale it and speak as you exhale.

The most powerful hydrogen bomb is based on the splitting of the lightest atom.

Hydrogen(lat. Hydrogenium), H, chemical element, first by serial number in Mendeleev’s periodic system; atomic mass 1.0079. At normal conditions Hydrogen is a gas; has no color, smell or taste.

Distribution of Hydrogen in nature. Hydrogen is widespread in nature; its content in the earth's crust (lithosphere and hydrosphere) is 1% by mass and 16% by number of atoms. Hydrogen is part of the most common substance on Earth - water (11.19% of Hydrogen by mass), in the composition of compounds that make up coal, oil, natural gases, clays, as well as animal and plant organisms (that is, in the composition of proteins, nucleic acids, fats, carbohydrates and others). Hydrogen is extremely rare in its free state; it is found in small quantities in volcanic and other natural gases. Minor amounts of free Hydrogen (0.0001% by number of atoms) are present in the atmosphere. In near-Earth space, Hydrogen in the form of a flow of protons forms the internal (“proton”) radiation belt of the Earth. In space, Hydrogen is the most abundant element. In the form of plasma, it makes up about half the mass of the Sun and most stars, the bulk of the gases of the interstellar medium and gaseous nebulae. Hydrogen is present in the atmosphere of a number of planets and in comets in the form of free H 2, methane CH 4, ammonia NH 3, water H 2 O, radicals such as CH, NH, OH, SiH, PH, etc. Hydrogen enters in the form of a flow of protons into the corpuscular radiation of the Sun and cosmic rays.

Isotopes, atom and molecule of Hydrogen. Ordinary Hydrogen consists of a mixture of 2 stable isotopes: light Hydrogen, or protium (1 H), and heavy Hydrogen, or deuterium (2 H, or D). In natural compounds of Hydrogen, there are on average 6800 atoms of 1 H per 1 atom of 2 H. A radioactive isotope with a mass number of 3 is called superheavy Hydrogen, or tritium (3 H, or T), with soft β-radiation and a half-life T ½ = 12.262 years . In nature, tritium is formed, for example, from atmospheric nitrogen under the influence of cosmic ray neutrons; in the atmosphere it is negligible (4·10 -15% of total number hydrogen atoms). An extremely unstable isotope 4 H was obtained. The mass numbers of the isotopes 1 H, 2 H, 3 H and 4 H, respectively 1, 2, 3 and 4, indicate that the nucleus of the protium atom contains only one proton, and that of deuterium - one proton and one neutron, tritium - one proton and 2 neutrons, 4 H - one proton and 3 neutrons. The large difference in the masses of Hydrogen isotopes causes a more noticeable difference in their physical and chemical properties than in the case of isotopes of other elements.

The Hydrogen atom has the simplest structure among the atoms of all other elements: it consists of a nucleus and one electron. The binding energy of an electron with a nucleus (ionization potential) is 13.595 eV. The neutral hydrogen atom can also attach a second electron, forming a negative H ion - in this case, the binding energy of the second electron with a neutral atom (electron affinity) is 0.78 eV. Quantum mechanics allows us to calculate all possible energy levels of the Hydrogen atom, and therefore give a complete interpretation of its atomic spectrum. The Hydrogen atom is used as a model atom in quantum mechanical calculations of the energy levels of other, more complex atoms.

The Hydrogen H2 molecule consists of two atoms connected by a covalent chemical bond. The energy of dissociation (that is, decay into atoms) is 4.776 eV. The interatomic distance at the equilibrium position of the nuclei is 0.7414 Å. At high temperatures, molecular Hydrogen dissociates into atoms (the degree of dissociation at 2000°C is 0.0013, at 5000°C 0.95). Atomic Hydrogen is also formed in various chemical reactions (for example, by the action of Zn on hydrochloric acid). However, the existence of Hydrogen in the atomic state lasts only a short time, the atoms recombine into H 2 molecules.

Physical properties of Hydrogen. Hydrogen is the lightest of all known substances (14.4 times lighter than air), density 0.0899 g/l at 0°C and 1 atm. Hydrogen boils (liquefies) and melts (solidifies) at -252.8°C and -259.1°C, respectively (only helium has lower melting and boiling points). Critical temperature Hydrogen is very low (-240°C), so its liquefaction is fraught with great difficulties; critical pressure 12.8 kgf/cm 2 (12.8 atm), critical density 0.0312 g/cm 3. Of all gases, Hydrogen has the highest thermal conductivity, equal to 0.174 W/(m·K) at 0°C and 1 atm, that is, 4.16·10 -4 cal/(s·cm·°C). The specific heat of Hydrogen at 0°C and 1 atm C p 14.208 kJ/(kg·K), that is, 3.394 cal/(g·°С). Hydrogen is slightly soluble in water (0.0182 ml/g at 20°C and 1 atm), but well soluble in many metals (Ni, Pt, Pa and others), especially in palladium (850 volumes per 1 volume of Pd). The solubility of Hydrogen in metals is related to its ability to diffuse through them; Diffusion through a carbon alloy (for example, steel) is sometimes accompanied by destruction of the alloy due to the interaction of Hydrogen with carbon (so-called decarbonization). Liquid Hydrogen is very light (density at -253°C 0.0708 g/cm 3) and fluid (viscosity at -253°C 13.8 spuaz).

Chemical properties Hydrogen. In most compounds, Hydrogen exhibits a valence (more precisely, oxidation state) +1, like sodium and other alkali metals; it is usually considered as an analogue of these metals, heading group I of the periodic system. However, in metal hydrides, the Hydrogen ion is negatively charged (oxidation state -1), that is, the hydride Na + H - is built similar to the chloride Na + Cl -. This and some other facts (the similarity of the physical properties of Hydrogen and halogens, the ability of halogens to replace Hydrogen in organic compounds) give grounds to classify Hydrogen also in Group VII of the periodic table. 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. Atomic Hydrogen has increased chemical activity compared to molecular Hydrogen. With oxygen, hydrogen forms water:

H 2 + 1/2 O 2 = H 2 O

with the release of 285.937 kJ/mol, that is, 68.3174 kcal/mol of heat (at 25°C and 1 atm). At normal temperatures the reaction proceeds extremely slowly, above 550°C it explodes. The explosive limits of a hydrogen-oxygen mixture are (by volume) from 4 to 94% H2, and of a hydrogen-air mixture - from 4 to 74% H2 (a mixture of 2 volumes of H2 and 1 volume of O2 is called detonating gas). Hydrogen is used to reduce many metals, as it removes oxygen from their oxides:

CuO + H 2 = Cu + H 2 O,

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

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. Hydrogen reacts with nitrogen to form ammonia:

ZN 2 + N 2 = 2NH 3

only on a catalyst and at elevated temperatures and pressures. When heated, Hydrogen reacts vigorously with sulfur:

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

much more difficult with selenium and tellurium. Hydrogen can react with pure carbon without a catalyst only at high temperatures:

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

Hydrogen reacts directly with some metals (alkali, alkaline earth and others), forming hydrides:

H 2 + 2Li = 2LiH.

Of great practical importance are the reactions of Hydrogen with carbon monoxide (II), in which various organic compounds are formed, depending on temperature, pressure and catalyst, for example HCHO, CH 3 OH and others. Unsaturated hydrocarbons react with Hydrogen, becoming saturated, for example:

C n H 2n + H 2 = C n H 2n+2.

The role of Hydrogen and its compounds in chemistry is exceptionally great. Hydrogen determines the acidic properties of so-called protic acids. Hydrogen tends to form a so-called hydrogen bond with some elements, which has a decisive influence on the properties of many organic and inorganic compounds.

Obtaining Hydrogen. The main types of raw materials for the industrial production of Hydrogen are natural flammable gases, coke oven gas and oil refining gases. Hydrogen is also obtained from water by electrolysis (in places with cheap electricity). The most important methods for producing Hydrogen from natural gas are the catalytic interaction of hydrocarbons, mainly methane, with water vapor (conversion):

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

and incomplete oxidation of hydrocarbons with oxygen:

CH 4 + 1/2 O 2 = CO + 2H 2

The resulting carbon monoxide (II) also undergoes conversion:

CO + H 2 O = CO 2 + H 2.

Hydrogen produced from natural gas is the cheapest.

Hydrogen is isolated from coke oven gas and oil refining gases by removing the remaining components of the gas mixture, which liquefy more easily than Hydrogen during deep cooling. Electrolysis of water is carried out with direct current, passing it through a solution of KOH or NaOH (acids are not used to avoid corrosion of steel equipment). In laboratories, Hydrogen is obtained by electrolysis of water, as well as by the reaction between zinc and hydrochloric acid. However, more often they use ready-made hydrogen in cylinders.

Application of Hydrogen. Hydrogen began to be produced on an industrial scale at the end of the 18th century for filling balloons. Currently, Hydrogen is widely used in chemical industry, mainly for the production of ammonia. A major consumer of Hydrogen is also the production of methyl and other alcohols, synthetic gasoline and other products obtained by synthesis from Hydrogen and carbon monoxide (II). Hydrogen is used for the hydrogenation of solid and heavy liquid fuels, fats and others, for the synthesis of HCl, for the hydrotreatment of petroleum products, in welding and cutting of metals with an oxygen-hydrogen flame (temperature up to 2800°C) and in atomic-hydrogen welding (up to 4000°C) . Hydrogen isotopes - deuterium and tritium - have found very important applications in nuclear energy.

Phenols

Structure
The hydroxyl group in molecules of organic compounds can be associated with the aromatic ring directly, or can be separated from it by one or more carbon atoms. It can be expected that, depending on this, the properties of substances will differ significantly from each other due to the mutual influence of groups of atoms (remember one of the provisions of Butlerov’s theory). Indeed, organic compounds containing the aromatic radical phenyl C 6 H 5 -, directly bonded to the hydroxyl group, exhibit special properties that differ from the properties of alcohols. Such compounds are called phenols.

Phenols - organic matter, the molecules of which contain a phenyl radical associated with one or more hydroxy groups.
Just like alcohols, phenols are classified by atomicity, i.e., by the number of hydroxyl groups. Monohydric phenols contain one hydroxyl group in the molecule:

There are other polyatomic phenols containing three or more hydroxyl groups in the benzene ring.
Let's take a closer look at the structure and properties of the simplest representative of this class - phenol C6H50H. The name of this substance formed the basis for the name of the entire class - phenols.

Physical properties
Solid colorless crystalline substance, tºmel = 43 °C, tº boil = °C, with a sharp characteristic odor. Poisonous. Phenol is slightly soluble in water at room temperature. An aqueous solution of phenol is called carbolic acid. It causes burns if it comes into contact with the skin, so phenol must be handled with care.
The structure of the phenol molecule
In the phenol molecule, the hydroxyl is directly bonded to the carbon atom of the benzene aromatic ring.
Let us recall the structure of the groups of atoms that form the phenol molecule.
The aromatic ring consists of six carbon atoms forming a regular hexagon due to the sp 2 hybridization of the electronic orbitals of the six carbon atoms. These atoms are connected by Þ bonds. The p-electrons of each carbon atom that do not participate in the formation of st-bonds, overlapping in different sides planes of Þ-bonds, form two parts of a single six-electron P-cloud covering the entire benzene ring (aromatic core). In the C6H6 benzene molecule, the aromatic ring is absolutely symmetrical, with a single electronic P-the cloud evenly covers the ring of carbon atoms below and above the plane of the molecule (Fig. 24). The covalent bond between the oxygen and hydrogen atoms of the hydroxyl radical is highly polar, the general electron cloud of the O-H bond is shifted towards the oxygen atom, on which a partial negative charge arises, and on the hydrogen atom - a partial positive charge. In addition, the oxygen atom in the hydroxyl group has two lone electron pairs that belong only to it.

In a phenol molecule, the hydroxyl radical interacts with the aromatic ring, while the lone electron pairs of the oxygen atom interact with the single TC cloud of the benzene ring, forming a single electronic system. This interaction of lone electron pairs and clouds of π bonds is called conjugation. As a result of the conjugation of the lone electron pair of the oxygen atom of the hydroxy group with the electron system of the benzene ring, the electron density on the oxygen atom decreases. This decrease is compensated by greater polarization of the O-H bond, which, in turn, leads to an increase positive charge on the hydrogen atom. Consequently, the hydrogen of the hydroxyl group in the phenol molecule has an “acidic” character.
It is logical to assume that the conjugation of electrons of the benzene ring and the hydroxyl group affects not only its properties, but also the reactivity of the benzene ring.
In fact, as you remember, the conjugation of lone pairs of the oxygen atom with the l-cloud of the benzene ring leads to a redistribution of the electron density in it. It decreases at the carbon atom associated with the OH group (due to the influence of the electron pairs of the oxygen atom) and increases at its neighboring carbon atoms (i.e., positions 2 and 6, or ortho positions). It is obvious that an increase in the electron density of these carbon atoms of the benzene ring leads to the localization (concentration) of a negative charge on them. Under the influence of this charge, a further redistribution of the electron density in the aromatic nucleus occurs - its displacement from the 3rd and 5th atoms (meta position) to the 4th (ortho position). These processes can be expressed by the diagram:

Thus, the presence of a hydroxyl radical in a phenol molecule leads to a change in the l-cloud of the benzene ring, an increase in the electron density at the 2, 4 and 6th carbon atoms (ortho-, dara-position) and a decrease in the electron density at the 3rd and 5- th carbon atoms (meta positions).
The localization of electron density in ortho and para positions makes them most likely to be attacked by electrophilic species when interacting with other substances.
Consequently, the influence of the radicals that make up the phenol molecule is mutual, and it determines its characteristic properties.
Chemical properties of phenol
Acid properties
As already mentioned, the hydrogen atom of the hydroxyl group of phenol is acidic in nature. The acidic properties of phenol are more pronounced than those of water and alcohols. Unlike alcohols and water, phenol reacts not only with alkali metals, but also with alkalis to form phenolates.
However, the acidic properties of phenols are less pronounced than those of inorganic and carboxylic acids. For example, the acidic properties of phenol are approximately 3000 times less than those of carbonic acid. Therefore, by passing carbon dioxide through an aqueous solution of sodium phenolate, free phenol can be isolated:

Adding hydrochloric or sulfuric acid to an aqueous solution of sodium phenolate also leads to the formation of phenol.
Qualitative reaction to phenol
Phenol reacts with iron(III) chloride to form an intensely colored purple complex compound.
This reaction allows it to be detected even in very small quantities. Other phenols containing one or more hydroxyl groups on the benzene ring also give a bright blue-violet color when reacted with iron(III) chloride.
Benzene ring reactions
The presence of a hydroxyl substituent greatly facilitates the occurrence of electrophilic substitution reactions in the benzene ring.
1. Bromination of phenol. Unlike benzene, the bromination of phenol does not require the addition of a catalyst (iron(III) bromide).
In addition, the interaction with phenol proceeds selectively: bromine atoms are directed to the ortho and para positions, replacing the hydrogen atoms located there. The selectivity of substitution is explained by the features of the electronic structure of the phenol molecule discussed above. Thus, when phenol reacts with bromine water, a white precipitate of 2,4,6-tribromophenol is formed.
This reaction, like the reaction with iron(III) chloride, serves for the qualitative detection of phenol.

2. Nitration of phenol also occurs more easily than nitration of benzene. The reaction with dilute nitric acid occurs at room temperature. As a result, a mixture of ortho- and para-isomers of nitrophenol is formed:

3. Hydrogenation of the aromatic ring of phenol in the presence of a catalyst occurs easily.
4. Polycondensation of phenol with aldehydes, in particular with formaldehyde, occurs with the formation of reaction products - phenol-formaldehyde resins and solid polymers.
The interaction of phenol with formaldehyde can be described by the following scheme:

You probably noticed that “mobile” hydrogen atoms are retained in the dimer molecule, which means that the reaction can continue further if there is a sufficient amount of reagents.
The polycondensation reaction, i.e., the reaction of producing a polymer that occurs with the release of a low-molecular-weight by-product (water), can continue further (until one of the reagents is completely consumed) with the formation of huge macromolecules. The process can be described by the summary equation:

The formation of linear molecules occurs at ordinary temperatures. Carrying out this reaction when heated leads to the fact that the resulting product has a branched structure, it is solid and insoluble in water. As a result of heating a linear phenol-formaldehyde resin with an excess of aldehyde, hard plastic masses with unique properties are obtained. Polymers based on phenol-formaldehyde resins are used for the manufacture of varnishes and paints, plastic products that are resistant to heating, cooling, water, alkalis and acids; they have high dielectric properties. The most responsible and important details electrical appliances, power unit housings and machine parts, polymer base printed circuit boards for radio devices.

Adhesives based on phenol-formaldehyde resins are capable of reliably connecting parts of a wide variety of natures, maintaining the highest joint strength over a very wide temperature range. This adhesive is used to attach the metal base of lighting lamps to glass flask. Now you understand why phenol and products based on it are widely used (Scheme 8).

DEFINITION

Hydrogen– the first element of the Periodic Table of Chemical Elements D.I. Mendeleev. Symbol - N.

Atomic mass – 1 amu. The hydrogen molecule is diatomic – H2.

Electronic configuration hydrogen atom – 1s 1. Hydrogen belongs to the s-element family. In its compounds it exhibits oxidation states -1, 0, +1. Natural hydrogen consists of two stable isotopes - protium 1H (99.98%) and deuterium 2H (D) (0.015%) - and the radioactive isotope tritium 3H (T) (trace amounts, half-life - 12.5 years) .

Chemical properties of hydrogen

Under normal conditions, molecular hydrogen exhibits relatively low reactivity, which is explained by the high strength of bonds in the molecule. When heated, it interacts with almost all simple substances, formed by elements of the main subgroups (except for noble gases, B, Si, P, Al). In chemical reactions it can act both as a reducing agent (more often) and an oxidizing agent (less often).

Hydrogen exhibits properties of the reducing agent(H 2 0 -2e → 2H +) in the following reactions:

1. Reactions of interaction with simple substances - non-metals. Hydrogen reacts with halogens, moreover, the reaction of interaction with fluorine under normal conditions, in the dark, with an explosion, with chlorine - under illumination (or UV irradiation) according to a chain mechanism, with bromine and iodine only when heated; oxygen(a mixture of oxygen and hydrogen in a volume ratio of 2:1 is called “explosive gas”), gray, nitrogen And carbon:

H 2 + Hal 2 = 2HHal;

2H 2 + O 2 = 2H 2 O + Q (t);

H 2 + S = H 2 S (t = 150 – 300C);

3H 2 + N 2 ↔ 2NH 3 (t = 500C, p, kat = Fe, Pt);

2H 2 + C ↔ CH 4 (t, p, kat).

2. Reactions of interaction with complex substances. Hydrogen reacts with oxides of low-active metals, and it is capable of reducing only metals that are in the activity series to the right of zinc:

CuO + H 2 = Cu + H 2 O (t);

Fe 2 O 3 + 3H 2 = 2Fe + 3H 2 O (t);

WO 3 + 3H 2 = W + 3H 2 O (t).

Hydrogen reacts with non-metal oxides:

H 2 + CO 2 ↔ CO + H 2 O (t);

2H 2 + CO ↔ CH 3 OH (t = 300C, p = 250 – 300 atm., kat = ZnO, Cr 2 O 3).

Hydrogen enters into hydrogenation reactions with organic compounds of the class of cycloalkanes, alkenes, arenes, aldehydes and ketones, etc. All these reactions are carried out with heating, under pressure, using platinum or nickel as catalysts:

CH 2 = CH 2 + H 2 ↔ CH 3 -CH 3 ;

C 6 H 6 + 3H 2 ↔ C 6 H 12 ;

C 3 H 6 + H 2 ↔ C 3 H 8;

CH 3 CHO + H 2 ↔ CH 3 -CH 2 -OH;

CH 3 -CO-CH 3 + H 2 ↔ CH 3 -CH(OH)-CH 3.

Hydrogen as an oxidizing agent(H 2 +2e → 2H -) appears in reactions with alkali and alkaline earth metals. In this case, hydrides are formed - crystalline ionic compounds in which hydrogen exhibits an oxidation state of -1.

2Na +H 2 ↔ 2NaH (t, p).

Ca + H 2 ↔ CaH 2 (t, p).

Physical properties of hydrogen

Hydrogen is a light, colorless, odorless gas, density at ambient conditions. – 0.09 g/l, 14.5 times lighter than air, t boil = -252.8C, t pl = - 259.2C. Hydrogen is poorly soluble in water and organic solvents; it is highly soluble in some metals: nickel, palladium, platinum.

According to modern cosmochemistry, hydrogen is the most common element in the Universe. The main form of existence of hydrogen in outer space is individual atoms. Hydrogen is the 9th most abundant element on Earth among all elements. The main amount of hydrogen on Earth is in a bound state - in the composition of water, oil, natural gas, coal, etc. Hydrogen is rarely found in the form of a simple substance - in the composition of volcanic gases.

Hydrogen production

There are laboratory and industrial methods for producing hydrogen. Laboratory methods include the interaction of metals with acids (1), as well as the interaction of aluminum with aqueous solutions of alkalis (2). Among industrial methods for producing hydrogen, electrolysis of aqueous solutions of alkalis and salts (3) and methane conversion (4) play an important role:

Zn + 2HCl = ZnCl 2 + H 2 (1);

2Al + 2NaOH + 6H 2 O = 2Na +3 H 2 (2);

2NaCl + 2H 2 O = H 2 + Cl 2 + 2NaOH (3);

CH 4 + H 2 O ↔ CO + H 2 (4).

Examples of problem solving

EXAMPLE 1

EXAMPLE 2

HYDROGEN (Latin Hydrogenium), H, chemical element of group VII of the short form (group 1 of the long form) of the periodic system; atomic number 1, atomic mass 1.00794; non-metal. There are two stable isotopes in nature: protium 1H (99.985% by mass) and deuterium D, or 2H (0.015%). Artificially produced radioactive tritium 3 H, or T (ß-decay, T 1/2 12.26 years), is formed in nature in negligible quantities in the upper layers of the atmosphere as a result of the interaction of cosmic radiation mainly with N and O nuclei. Artificially obtained extremely unstable radioactive isotopes 4 H, 5 H, 6 H.

Historical reference. Hydrogen was first studied in 1766 by G. Cavendish and he called it “flammable air.” In 1787, A. Lavoisier showed that this gas forms water when burned, included it in the list of chemical elements and proposed the name hydrogène (from the Greek?δωρ - water and γενν?ω - to give birth).

Prevalence in nature. Hydrogen content in atmospheric air 3.5-10% by mass, 1% in the earth’s crust. The main reservoir of hydrogen on Earth is water (11.19% hydrogen by mass). Hydrogen is a biogenic element and is part of the compounds that form coal, oil, natural combustible gases, many minerals, etc. In near-Earth space, hydrogen in the form of a flow of protons forms the Earth's internal radiation belt. Hydrogen is the most abundant element in space; in the form of plasma it makes up about 70% of the mass of the Sun and stars, the bulk of the interstellar medium and gaseous nebulae, is present in the atmosphere of a number of planets in the form of H 2, CH 4, NH 3, H 2 O, etc.

Properties. The configuration of the electron shell of the hydrogen atom is 1s 1; in compounds exhibits oxidation states +1 and -1. Electronegativity according to Pauling 2.1; radii (pm): atomic 46, covalent 30, van der Waals 120; ionization energy Н°→ Н + 1312.0 kJ/mol. In the free state, hydrogen forms a diatomic H 2 molecule, the internuclear distance is 76 pm, the dissociation energy is 432.1 kJ/mol (0 K). Depending on the relative orientation of the nuclear spins, there are ortho-hydrogen (parallel spins) and para-hydrogen (antiparallel spins), differing in magnetic, optical and thermal properties and usually contained in a 3:1 ratio; the conversion of para-hydrogen to ortho-hydrogen requires 1418 J/mol of energy.

Hydrogen is a colorless, tasteless and odorless gas; t PL -259.19 °C, t KIP -252.77 °C. Hydrogen is the lightest and most thermally conductive of all gases: at 273 K, density is 0.0899 kg/m 3, thermal conductivity is 0.1815 W/(m K). Insoluble in water; dissolves well in many metals (best in Pd - up to 850% by volume); diffuses through many materials (eg steel). Burns in air and forms explosive mixtures. Solid hydrogen crystallizes in a hexagonal lattice; at pressures above 10 4 MPa, a phase transition is possible with the formation of a structure built from atoms and possessing metallic properties - the so-called metallic hydrogen.

Hydrogen forms compounds with many elements. With oxygen it forms water (at temperatures above 550 °C the reaction is accompanied by an explosion), with nitrogen - ammonia, with halogens - hydrogen halides, with metals, intermetallic compounds, as well as with many non-metals (for example, chalcogens) - hydrides, with carbon - hydrocarbons. Reactions with CO are of practical importance (see Synthesis gas). Hydrogen reduces the oxides and halides of many metals to metals, and unsaturated hydrocarbons to saturated ones (see Hydrogenation). The nucleus of the hydrogen atom - the H + proton - determines the acidic properties of compounds. In aqueous solutions, H + forms hydronium ion H 3 O + with a water molecule. Composed of molecules various connections Hydrogen tends to form hydrogen bonds with many electronegative elements.

Application. Hydrogen gas is used in the industrial synthesis of ammonia, hydrochloric acid, methanol and higher alcohols, synthetic liquid fuels, etc., for the hydrogenation of fats and other organic compounds; in oil refining - for hydrotreating and hydrocracking of oil fractions; in metallurgy - to obtain metals (for example, W, Mo, Re from their oxides and fluorides), to create a protective environment when processing metals and alloys; in the production of quartz glass products using a hydrogen-oxygen flame, for atomic-hydrogen welding of refractory steels and alloys, etc., as lifting gas for balloons. Liquid hydrogen is a fuel in rocket and space technology; also used as a refrigerant.

For information on the main methods of production, as well as storage, transportation and use of hydrogen as an energy carrier, see Hydrogen Energy.

Lit. look at Art. Hydrogen energy.

The hydrogen atom has electronic formula outer (and only) electronic level 1 s 1 . On the one hand, due to the presence of one electron on the outer electronic level The hydrogen atom is similar to alkali metal atoms. However, just like halogens, it only needs one electron to fill the outer electronic level, since the first electronic level can contain no more than 2 electrons. It turns out that hydrogen can be placed simultaneously in both the first and the penultimate (seventh) group of the periodic table, which is sometimes done in various options periodic system:

From the point of view 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 diatomic molecules (H 2) like them.

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

Interaction of hydrogen with simple substances

with metals

Of the metals, hydrogen reacts only with alkali and alkaline earth metals! Alkali metals include metals of the main subgroup Group I(Li, Na, K, Rb, Cs, Fr), and alkaline earth metals - metals of the main subgroup of group II, except 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. The reaction occurs when heated:

It should be noted that interaction with active metals is the only case when molecular hydrogen H2 is an oxidizing agent.

with non-metals

Of the 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, it can 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 activity series of metals up to aluminum (inclusive), however, it is capable of reducing many metal oxides to the right of aluminum when heated:

with non-metal oxides

Of the non-metal oxides, hydrogen reacts when heated with the oxides of nitrogen, halogens and carbon. Of all the interactions of hydrogen with non-metal oxides, especially noteworthy is its reaction with carbon monoxide CO.

The mixture of CO and H2 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 acids, as well as with acids containing functional groups capable of reduction with 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

Halogens are called chemical elements Group VIIA (F, Cl, Br, I, At), as well as the simple substances they form. Here and further in the text, unless otherwise stated, halogens will be understood as simple substances.

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

It should be noted that this specific physical property Yoda, how his ability to sublimation or, in other words, sublimation. Sublimation, is a phenomenon in which a substance in a solid state does not melt when heated, but, bypassing the liquid phase, immediately passes into the gaseous state.

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 periodic table period in which the halogen is located. As you can see, the halogen atoms only need one electron to reach the eight-electron outer shell. From this it is logical to assume the predominantly oxidizing properties of free halogens, which is confirmed in practice. As is known, the electronegativity of nonmetals decreases when moving down a subgroup, and therefore the activity of halogens decreases in the series:

F 2 > Cl 2 > Br 2 > I 2

Interaction of halogens with simple substances

All halogens are highly reactive substances 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. Such 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 remaining 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 interact with hydrogen, they form hydrogen halides with the general formula HHal. In this case, 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 heat. Also proceeds with 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, phosphorus pentafluoride is formed:

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

Moreover, in case white phosphorus in an atmosphere of fluorine, chlorine or liquid bromine, the reaction begins spontaneously.

The interaction of phosphorus with iodine can lead to the formation of only phosphorus triodide due to its 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 oxidation states +1 and +2, which are extremely unusual for it. These interactions are very specific, and for passing the Unified State Exam in chemistry, the ability to write equations for these interactions is not necessary. Therefore, the following three equations are given rather for reference:

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 remaining 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 capable of displacing less active halogens from the hydrohalic acids and metal halides they form:

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

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

Reaction of halogens with water

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

Bromine and chlorine react differently with water than fluorine. If fluorine acted as an oxidizing agent, then chlorine and bromine are 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 degree that it can be neglected and it can be assumed that the reaction does not occur at all.

Interaction of halogens with alkali solutions

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

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

Unlike fluorine, other halogens in alkali solutions are 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, it is possible to flow through two different directions. In particular, in the cold the 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 hypoiodite is not stable not only when heated, but also at ordinary temperatures and even in the cold.