The most common element in the universe is hydrogen. In the matter of stars, it has the form of nuclei - protons - and is a material for thermonuclear processes. Almost half of the Sun's mass also consists of H 2 molecules. Its content in the earth's crust reaches 0.15%, and atoms are present in oil, natural gas, and water. Together with oxygen, nitrogen and carbon, it is an organogenic element that is part of all living organisms on Earth. In our article we will study the physical and chemical properties of hydrogen, determine the main areas of its application in industry and its significance in nature.

Position in Mendeleev's periodic table of chemical elements

The first element to discover the periodic table is hydrogen. Its atomic mass is 1.0079. It has two stable isotopes (protium and deuterium) and one radioactive isotope (tritium). Physical properties are determined by the place of the nonmetal in the table of chemical elements. Under normal conditions, hydrogen (its formula is H2) is a gas that is almost 15 times lighter than air. The structure of the element's atom is unique: it consists of only a nucleus and one electron. The molecule of the substance is diatomic; the particles in it are connected using a covalent nonpolar bond. Its energy intensity is quite high - 431 kJ. This explains the low chemical activity of the compound under normal conditions. The electronic formula of hydrogen is: H:H.

The substance also has a number of properties that have no analogues among other non-metals. Let's look at some of them.

Solubility and thermal conductivity

Metals conduct heat best, but hydrogen is close to them in thermal conductivity. The explanation for the phenomenon lies in the very high speed of thermal movement of light molecules of a substance, therefore in a hydrogen atmosphere a heated object cools down 6 times faster than in air. The compound can be highly soluble in metals; for example, almost 900 volumes of hydrogen can be absorbed by one volume of palladium. Metals can enter into chemical reactions with H2, in which the oxidizing properties of hydrogen are manifested. In this case, hydrides are formed:

2Na + H 2 =2 NaH.

In this reaction, atoms of the element accept electrons from metal particles, becoming anions with a single negative charge. The simple substance H2 in this case is an oxidizing agent, which is usually not typical for it.

Hydrogen as a reducing agent

What unites metals and hydrogen is not only high thermal conductivity, but also the ability of their atoms in chemical processes to give up their own electrons, that is, to oxidize. For example, basic oxides react with hydrogen. The redox reaction ends with the release of pure metal and the formation of water molecules:

CuO + H 2 = Cu + H 2 O.

The interaction of a substance with oxygen when heated also leads to the production of water molecules. The process is exothermic and is accompanied by the release of a large amount of thermal energy. If a gas mixture of H 2 and O 2 reacts in a 2:1 ratio, then it is called because it explodes when ignited:

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

Water is and plays a vital role in the formation of the Earth’s hydrosphere, climate, and weather. It ensures the circulation of elements in nature, supports all the life processes of organisms - the inhabitants of our planet.

Interaction with non-metals

The most important chemical properties of hydrogen are its reactions with non-metallic elements. Under normal conditions, they are quite chemically inert, so the substance can only react with halogens, for example with fluorine or chlorine, which are the most active among all non-metals. Thus, a mixture of fluorine and hydrogen explodes in the dark or in the cold, and with chlorine - when heated or in the light. The reaction products will be hydrogen halides, aqueous solutions of which are known as fluoride and chloride acids. C interacts at a temperature of 450-500 degrees, a pressure of 30-100 mPa and in the presence of a catalyst:

N₂ + 3H₂ ⇔ p, t, kat ⇔ 2NH₃.

The considered chemical properties of hydrogen are of great importance for industry. For example, you can obtain a valuable chemical product - ammonia. It is the main raw material for the production of nitrate acid and nitrogen fertilizers: urea, ammonium nitrate.

Organic matter

Between carbon and hydrogen leads to the production of the simplest hydrocarbon - methane:

C + 2H 2 = CH 4.

The substance is the most important component of natural and They are used as a valuable type of fuel and raw material for the organic synthesis industry.

In the chemistry of carbon compounds, the element is part of a huge number of substances: alkanes, alkenes, carbohydrates, alcohols, etc. Many reactions of organic compounds with H 2 molecules are known. They have a common name - hydrogenation or hydrogenation. Thus, aldehydes can be reduced with hydrogen to alcohols, unsaturated hydrocarbons - to alkanes. For example, ethylene is converted to ethane:

C 2 H 4 + H 2 = C 2 H 6.

The chemical properties of hydrogen, such as, for example, the hydrogenation of liquid oils: sunflower, corn, rapeseed, are of important practical importance. It leads to the production of solid fat - lard, which is used in the production of glycerin, soap, stearin, and hard margarine. To improve the appearance and taste of a food product, milk, animal fats, sugar, and vitamins are added to it.

In our article, we studied the properties of hydrogen and found out its role in nature and human life.

HYDROGEN
N (lat. hydrogenium),
the lightest gaseous chemical element is a member of subgroup IA of the periodic table of elements, sometimes it is classified as subgroup VIIA. In the earth's atmosphere, hydrogen exists in an unbound state for only a fraction of a minute; its amount is 1-2 parts per 1,500,000 parts of air. It is usually released with other gases during volcanic eruptions, from oil wells and in places where large quantities of organic matter decompose. Hydrogen combines with carbon and/or oxygen in organic matter such as carbohydrates, hydrocarbons, fats and animal proteins. In the hydrosphere, hydrogen is part of water, the most common compound on Earth. In rocks, soils, and other parts of the earth's crust, hydrogen combines with oxygen to form water and the hydroxide ion OH-. Hydrogen makes up 16% of all atoms in the earth's crust, but only about 1% by mass, since it is 16 times lighter than oxygen. The mass of the Sun and stars is 70% hydrogen plasma: this is the most common element in space. The concentration of hydrogen in the Earth's atmosphere increases with altitude due to its low density and ability to rise to high altitudes. Meteorites found on the surface of the Earth contain 6-10 hydrogen atoms per 100 silicon atoms.
Historical reference. Another German doctor and naturalist Paracelsus in the 16th century. established the flammability of hydrogen. In 1700 N. Lemery discovered that the gas released by the action of sulfuric acid on iron explodes in air. Hydrogen as an element was identified by G. Cavendish in 1766 and called it “combustible air,” and in 1781 he proved that water is a product of its interaction with oxygen. The Latin hydrogenium, which comes from the Greek combination “giving birth to water,” was assigned to this element by A. Lavoisier.
General characteristics of hydrogen. Hydrogen is the first element in the periodic table of elements; its atom consists of one proton and one electron rotating around it
(see also PERIODIC SYSTEM OF ELEMENTS).
One of 5000 hydrogen atoms is distinguished by the presence of one neutron in the nucleus, increasing the mass of the nucleus from 1 to 2. This isotope of hydrogen is called deuterium 21H or 21D. Another, rarer isotope of hydrogen contains two neutrons in the nucleus and is called tritium 31H or 31T. Tritium is radioactive and decays to release helium and electrons. The nuclei of different isotopes of hydrogen differ in the spins of their protons. Hydrogen can be obtained a) by the action of an active metal on water, b) by the action of acids on certain metals, c) by the action of bases on silicon and some amphoteric metals, d) by the action of superheated steam on coal and methane, as well as on iron, e) by electrolytic decomposition water and thermal decomposition of hydrocarbons. The chemical activity of hydrogen is determined by its ability to donate an electron to another atom or share it almost equally with other elements when forming a chemical bond, or to attach an electron of another element in a chemical compound called a hydride. Hydrogen produced by industry is used in huge quantities for the synthesis of ammonia, nitric acid, and metal hydrides. The food industry uses hydrogen to hydrogenate (hydrogenate) liquid vegetable oils into solid fats (such as margarine). During hydrogenation, saturated organic oils containing double bonds between carbon atoms are converted into saturated ones having single carbon-carbon bonds. High-purity (99.9998%) liquid hydrogen is used in space rockets as a highly efficient fuel.
Physical properties. Hydrogen requires very low temperatures and high pressure to liquefy and solidify (see properties table). Under normal conditions, hydrogen is a colorless gas, odorless and tasteless, very light: 1 liter of hydrogen at 0° C and atmospheric pressure has a mass of 0.08987 g (cf. the density of air and helium 1.2929 and 0.1785 g/l, respectively ; therefore, a balloon filled with helium and having the same lift as a balloon filled with hydrogen should have 8% more volume). The table shows some physical and thermodynamic properties of hydrogen. PROPERTIES OF ORDINARY HYDROGEN
(at 273.16 K, or 0 ° C)
Atomic number 1 Atomic mass 11H 1.00797 Density, g/l

at normal pressure 0.08987 at 2.5*10 5 atm 0.66 at 2.7*10 18 atm 1.12*10 7

in water, volume/100 volumes of H2O (under standard conditions) 2.148 in benzene, ml/g (35.2° C, 150.2 atm) 11.77 in ammonia, ml/g (25° C) at 50 atm 4 .47 at 1000 atm 79.25

Nonmetals form volatile hydrides of the general formula MHx (x is an integer) with a relatively low boiling point and high vapor pressure. These hydrides differ significantly from salt hydrides, in which the hydrogen has a more negative charge. In volatile hydrides (e.g. hydrocarbons), covalent bonding between nonmetals and hydrogen predominates. As the nonmetallic character increases, compounds with partially ionic bonds are formed, for example H+Cl-, (H2)2+O2-, N3-(H3)3+. Some examples of the formation of various hydrides are given below (the heat of hydride formation is indicated in parentheses):

Isomerism and isotopes of hydrogen. The atoms of hydrogen isotopes are not alike. Ordinary hydrogen, protium, is always a proton around which one electron rotates, located at a huge distance from the proton (relative to the size of the proton). Both particles have spin, so hydrogen atoms can differ in either the electron spin, the proton spin, or both. Hydrogen atoms that differ in the spin of the proton or electron are called isomers. The combination of two atoms with parallel spins results in the formation of an “orthohydrogen” molecule, and those with opposite spins of protons result in a “parahydrogen” molecule. Chemically, both molecules are identical. Orthohydrogen has a very weak magnetic moment. At room or elevated temperatures, both isomers, orthohydrogen and parahydrogen, are usually in equilibrium in a ratio of 3:1. When cooled to 20 K (-253° C), the parahydrogen content increases to 99%, as it is more stable. When liquefied by industrial purification methods, the orthoform transforms into paraform with the release of heat, which causes hydrogen loss from evaporation. The rate of conversion of orthoform to paraform increases in the presence of a catalyst, such as charcoal, nickel oxide, chromium oxide supported on alumina. Protium is an unusual element because it has no neutrons in its nucleus. If a neutron appears in the nucleus, then such hydrogen is called deuterium 21D. Elements with the same number of protons and electrons and different numbers of neutrons are called isotopes. Natural hydrogen contains a small proportion of HD and D2. Similarly, natural water contains low concentrations (less than 0.1%) of DOH and D2O. Heavy water D2O, which has a mass greater than that of H2O, differs in physical and chemical properties, for example, the density of ordinary water is 0.9982 g/ml (20° C), and that of heavy water is 1.105 g/ml, the melting point of ordinary water is 0. 0 ° C, and heavy - 3.82 ° C, boiling point - 100 ° C and 101.42 ° C, respectively. Reactions involving D2O proceed at a lower speed (for example, electrolysis of natural water containing an admixture of D2O with the addition of alkali NaOH ). The rate of electrolytic decomposition of protium oxide H2O is greater than that of D2O (taking into account the constant increase in the proportion of D2O undergoing electrolysis). Due to the similar properties of protium and deuterium, it is possible to replace protium with deuterium. Such connections are referred to as so-called tags. By mixing deuterium compounds with ordinary hydrogen-containing substances, it is possible to study the paths, nature and mechanism of many reactions. This method is used to study biological and biochemical reactions, such as digestion processes. A third isotope of hydrogen, tritium (31T), occurs naturally in trace amounts. Unlike stable deuterium, tritium is radioactive and has a half-life of 12.26 years. Tritium decays to helium (32He) releasing a b particle (electron). Tritium and metal tritides are used to produce nuclear energy; for example, in a hydrogen bomb the following thermonuclear fusion reaction occurs: 21H + 31H -> 42He + 10n + 17.6 MeV
Hydrogen production. Often, the further use of hydrogen is determined by the nature of the production itself. In some cases, for example in the synthesis of ammonia, small amounts of nitrogen in the starting hydrogen, of course, are not a harmful impurity. An admixture of carbon(II) monoxide will also not be a problem if hydrogen is used as a reducing agent. 1. The largest production of hydrogen is based on the catalytic conversion of hydrocarbons with steam according to the scheme CnH2n + 2 + nH2O (r) nCO + (2n + 1)H2 and CnH2n + 2 + 2nH2O (r) nCO2 + (3n + 1)H2. The process temperature depends on the composition of the catalyst. It is known that the reaction temperature with propane can be reduced to 370° C using bauxite as a catalyst. Up to 95% of the CO produced in this case is consumed in a further reaction with water vapor: H2O + CO -> CO2 + H2
2. The water gas method accounts for a significant portion of the total hydrogen production. The essence of the method is the reaction of water vapor with coke to form a mixture of CO and H2. The reaction is endothermic (DH° = 121.8 kJ/mol) and is carried out at 1000° C. The heated coke is treated with steam; The purified gas mixture released contains some hydrogen, a large percentage of CO and a small admixture of CO2. To increase the H2 yield, CO monoxide is removed by further steam treatment at 370°C, which produces more CO2. Carbon dioxide is fairly easy to remove by passing the gas mixture through a scrubber sprayed with countercurrent water. 3. Electrolysis. In the electrolytic process, hydrogen is actually a by-product of the production of the main products, chlor alkali (NaOH). Electrolysis is carried out in a slightly alkaline aqueous environment at 80° C and a voltage of about 2V, using an iron cathode and a nickel anode:

4. Iron-steam method, in which steam at 500-1000 ° C is passed over iron: 3Fe + 4H2O Fe3O4 + 4H2 + 160.67 kJ. The hydrogen produced by this method is usually used to hydrogenate fats and oils. The composition of iron oxide depends on the process temperature; at nC + (n + 1)H2
6. The next largest production volume is the methanol-steam method: CH3OH + H2O -> 3H2 + CO2. The reaction is endothermic and is carried out at HYDROGEN 260° C in conventional steel reactors at pressures up to 20 atm. 7. Catalytic decomposition of ammonia: 2NH3 -> The reaction is reversible. When hydrogen requirements are small, this process is uneconomical. There are also various methods for producing hydrogen, which, although not of great industrial importance, in some cases may prove to be the most economically advantageous. Very pure hydrogen is obtained by hydrolysis of purified alkali metal hydrides; in this case, a lot of hydrogen is formed from a small amount of hydride: LiH + H2O -> LiOH + H2
(This method is convenient when directly using the resulting hydrogen.) When acids interact with active metals, hydrogen is also released, but it is usually contaminated with acid vapor or another gaseous product, for example, phosphine PH3, hydrogen sulfide H2S, arsine AsH3. The most active metals, reacting with water, displace hydrogen and form an alkaline solution: 2H2O + 2Na -> H2 + 2NaOH A common laboratory method for obtaining H2 in the Kipp apparatus is by reacting zinc with hydrochloric or sulfuric acid:
Zn + 2HCl -> ZnCl2 + H2. Alkaline earth metal hydrides (for example, CaH2), complex salt hydrides (for example, LiAlH4 or NaBH4) and some borohydrides (for example, B2H6) release hydrogen when reacting with water or during thermal dissociation. Brown coal and steam at high temperatures also react to release hydrogen.
Hydrogen purification. The degree of required purity of hydrogen is determined by its field of application. Carbon dioxide impurities are removed by freezing or liquefaction (for example, by passing the gaseous mixture through liquid nitrogen). The same impurity can be completely removed by bubbling through water. CO can be removed by catalytic conversion to CH4 or CO2 or by liquefaction by treatment with liquid nitrogen. The oxygen impurity formed during the electrolysis process is removed in the form of water after a spark discharge.
Application of hydrogen. Hydrogen is used mainly in the chemical industry for the production of hydrogen chloride, ammonia, methanol and other organic compounds. It is used in the hydrogenation of oils, as well as coal and petroleum (to convert low-grade fuels into high-quality ones). In metallurgy, some non-ferrous metals are reduced from their oxides using hydrogen. Hydrogen is used to cool powerful electric generators. Hydrogen isotopes are used in nuclear energy. Hydrogen-oxygen flame is used for cutting and welding metals.
LITERATURE
Nekrasov B.V. Fundamentals of general chemistry. M., 1973 Liquid hydrogen. M., 1980 Hydrogen in metals. M., 1981

Collier's Encyclopedia. - Open Society. 2000 .

See what "HYDROGEN" is in other dictionaries:

Table of nuclides General information Name, symbol Hydrogen 4, 4H Neutrons 3 Protons 1 Nuclide properties Atomic mass 4.027810(110) ... Wikipedia

Table of nuclides General information Name, symbol Hydrogen 5, 5H Neutrons 4 Protons 1 Nuclide properties Atomic mass 5.035310(110) ... Wikipedia

Nuclide table General information Name, symbol Hydrogen 6, 6H Neutrons 5 Protons 1 Nuclide properties Atomic mass 6.044940(280) ... Wikipedia

Table of nuclides General information Name, symbol Hydrogen 7, 7H Neutrons 6 Protons 1 Properties of the nuclide Atomic mass 7.052750 (1080) ... Wikipedia

Hydrogen

Hydrogen is the first element and one of two representatives of the first period of the Periodic Table. The hydrogen atom consists of two particles - a proton and an electron, between which there are only attractive forces. Hydrogen and group IA metals exhibit an oxidation state of +1, are reducing agents, and have similar optical spectra. However, in the state of a singly charged H+ cation (proton), hydrogen has no analogues. In addition, the ionization energy of a hydrogen atom is much greater than the ionization energy of alkali metal atoms.

On the other hand, both hydrogen and halogens are one electron short of completing the outer electron layer. Like halogens, hydrogen exhibits an oxidation state of –1 and oxidizing properties. Hydrogen is similar to halogens both in its state of aggregation and in the composition of E 2 molecules. But the molecular orbital (MO) of H 2 has nothing in common with those of halogen molecules, while at the same time the MO of H 2 has a certain similarity with the MO of diatomic molecules of alkali metals that exist in the vapor state.

Hydrogen is the most common element in the Universe and makes up the bulk of the Sun, stars and other cosmic bodies. On Earth it ranks 9th in prevalence; It is rare in a free state, and the main part of it is found in water, clays, coal and brown coal, oil, etc., as well as complex substances of living organisms.

Natural hydrogen is a mixture of stable isotopes of protium 1 H (99.985%) and deuterium 2 H (2 D), radioactive tritium 3 H (3 T).

Simple substances. Possible molecules of light hydrogen are H 2 (diprotium), heavy hydrogen are D 2 (dideuterium), T 2 (ditritium), HD (protodeuterium), NT (prototritium), DT (deuterotritium).

H 2 (dihydrogen, diprotium)– a colorless, difficult to liquefy gas, very slightly soluble in water, better in organic solvents, chemisorbed by metals (Fe, Ni, Pt, Pd). Under normal conditions, it is relatively little active and directly interacts only with fluorine; at elevated temperatures reacts with metals, non-metals, metal oxides. The reducing ability of atomic hydrogen H0, formed during the thermal decomposition of molecular hydrogen or as a result of reactions directly in the zone of the reduction process, is especially high.

Hydrogen exhibits reducing properties when interacting with nonmetals, metal oxides, and halides:

H 2 0 + Cl 2 = 2H +1 Cl; 2H 2 + O 2 = 2H 2 O; CuO + H 2 = Cu + H 2 O

As an oxidizing agent, hydrogen interacts with active metals:

2Na + H 2 0 = 2NaH –1

Production and use of hydrogen. In industry, hydrogen is produced mainly from natural and associated gases, fuel gasification products and coke oven gas. Hydrogen production is based on catalytic reactions of interaction with water vapor (conversion) of hydrocarbons (mainly methane) and carbon monoxide (II), respectively:

CH 4 + H 2 O = CO + 3H 2 (cat. Ni, 800°C)

CO + H 2 O = CO 2 + H 2 (cat. Fe, 550°C)

An important way to obtain hydrogen is to separate it from coke oven gas and oil refining gases by deep cooling. Electrolysis of water (the electrolyte is usually an aqueous solution of alkali) provides the purest hydrogen.

In laboratory conditions, hydrogen is usually obtained by the action of zinc on solutions of sulfuric or hydrochloric acid:

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

Hydrogen is used in the chemical industry for the synthesis of ammonia, methanol, hydrogen chloride, for the hydrogenation of solid and liquid fuels, fats, etc. In the form of water gas (mixed with CO) it is used as fuel. When hydrogen burns in oxygen, a high temperature arises (up to 2600°C), which makes it possible to weld and cut refractory metals, quartz, etc. Liquid hydrogen is used as one of the most efficient jet fuels.

Hydrogen compounds (–I). Hydrogen compounds with less electronegative elements, in which it is negatively polarized, are classified as hydrides, i.e. mainly its compounds with metals.

In simple salt-like hydrides there is an H – anion. The most polar bond is observed in hydrides of active metals - alkali and alkaline earth (for example, KH, CaH 2). Chemically, ionic hydrides behave like basic compounds.

LiH + H 2 O = LiOH + H 2

Covalent hydrides include hydrides of non-metallic elements that are less electronegative than hydrogen itself (for example, hydrides of the composition SiH 4 and BH 3). By chemical nature, non-metal hydrides are acidic compounds.

SiH 4 + 3H 2 O = H 2 SiO 3 + 4H 2

Upon hydrolysis, basic hydrides form an alkali, and acidic hydrides form an acid.

Many transition metals form hydrides with a predominantly metallic bond character and non-stoichiometric composition. The idealized composition of metal hydrides most often corresponds to the formulas: M +1 H (VH, NbH, TaH), M +2 H 2 (TiH 2, ZrH 2) and M +3 H 3 (UN 3, RaH 3).

Hydrogen compounds (I). Positive polarization of hydrogen atoms is observed in its numerous compounds with covalent bonds. Under normal conditions, these are gases (HCl, H 2 S, H 3 N), liquids (H 2 O, HF, HNO 3), solids (H 3 PO 4, H 2 SiO 3). The properties of these compounds strongly depend on the nature of the electronegative element.

Lithium

Lithium is quite widespread in the earth's crust. It is part of many minerals, found in coal, soils, sea water, and also in living organisms. The most valuable minerals are spodumene LiAl(SiO 3) 2, amblygonitis LiAl(PO 4)F and lepidolite Li 2 Al 2 (SiO 3) 3 (F,OH) 2.

Simple substance. Li (lithium) – A silvery-white, soft, low-melting alkali metal, the lightest of the metals. Reactive; in air it is covered with an oxide-nitride film (Li 2 O, Li 3 N). Will ignite when heated moderately (above 200°C); turns the flame of a gas burner dark red. Strong reducing agent. Compared to sodium and the alkali metals themselves (potassium subgroup), lithium is a chemically less active metal. Under normal conditions, it reacts violently with all halogens. When heated, it directly combines with sulfur, coal, hydrogen and other non-metals. When heated, it burns in CO 2. Lithium forms intermetallic compounds with metals. In addition, it forms solid solutions with Na, Al, Zn and some other metals. Lithium energetically decomposes water, releasing hydrogen from it, and interacts even more easily with acids.

2Li + H 2 O = 2LiOH + H 2

2Li + 2НCl = 2LiСl + Н 2

3Li + 4HNO 3 (diluted) = 2LiNO 3 + NO + 2H 2 O

Lithium is stored under a layer of petroleum jelly or paraffin in sealed containers.

Receipt and application. Lithium is obtained by vacuum-thermal reduction of spodumene or lithium oxide; silicon or aluminum is used as a reducing agent.

2Li 2 O + Si = 4Li + SiO 2

3Li 2 O + 2Al = 6Li + A1 2 O 3

In electrolytic reduction, a melt of the eutectic mixture LiCl-KCl is used.

Lithium imparts a number of valuable physical and chemical properties to alloys. Thus, aluminum alloys containing up to 1% Li increase mechanical strength and corrosion resistance, the introduction of 2% Li into commercial copper significantly increases its electrical conductivity, etc. The most important area of ​​application of lithium is nuclear energy (as a coolant in nuclear reactors) . It is used as a source of tritium (3 H).

Lithium(I) compounds. Binary lithium compounds are colorless crystalline substances; are salts or salt-like compounds. In terms of their chemical nature, solubility and nature of hydrolysis, they resemble derivatives of calcium and magnesium. LiF, Li 2 CO 3, Li 3 PO 4, etc. are poorly soluble.

Peroxide compounds are of little character for lithium. However, Li 2 O 2 peroxide, Li 2 S 2 persulfide and Li 2 C 2 percarbide are known for it.

Lithium oxide Li 2 O is a basic oxide, obtained by the interaction of simple substances. Reacts actively with water, acids, acidic and amphoteric oxides.

Li 2 O + H 2 O = 2LiOH

Li 2 O + 2HCl (diluted) = 2LiCl + H 2 O

Li 2 O + CO 2 = Li 2 CO 3

Lithium hydroxide LiOH is a strong base, but in solubility and strength it is inferior to the hydroxides of other alkali metals, and unlike them, LiOH decomposes when heated:

2LiOH ↔ Li 2 O + H 2 O (800-1000°C, in the atmosphere H 2)

LiOH is produced by electrolysis of aqueous solutions of LiCl. Used as an electrolyte in batteries.

When lithium salts co-crystallize or fuse with similar compounds of other alkali metals, eutectic mixtures are formed (LiNO 3 –KNO 3, etc.); less commonly, double compounds are formed, for example M +1 LiSO 4, Na 3 Li(SO 4) 2 ∙6H 2 O and solid solutions.

Melts of lithium salts and their mixtures are non-aqueous solvents; Most metals dissolve in them. These solutions are intensely colored and are very strong reducing agents. The dissolution of metals in molten salts is important for many electrometallurgical and metallothermic processes, for refining metals, and carrying out various syntheses.

Sodium

Sodium is one of the most abundant elements on Earth. Essential sodium minerals: rock salt or halite NaCl, mirabilite or Glauber's salt Na 2 SO 4 ∙10H 2 O, cryolite Na 3 AlF 6, borax Na 2 B 4 O 7 ∙10H 2 O, etc.; is part of many natural silicates and aluminosilicates. Sodium compounds are found in the hydrosphere (about 1.5∙10 t), in living organisms (for example, in human blood Na + ions account for 0.32%, in muscle tissue - up to 1.5%).

Simple substance. Na (sodium) – silvery-white, light, very soft, low-melting alkali metal. Very reactive; When exposed to air, it becomes covered with an oxide film (tarnishes), and ignites when heated moderately. Stable in an atmosphere of argon and nitrogen (reacts with nitrogen only when heated). Strong reducing agent; Reacts vigorously with water, acids, and non-metals. It forms an amalgam with mercury (unlike pure sodium, the reaction with water proceeds calmly). Colors the flame of a gas burner yellow.

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

2Na + 2НCl(dil.) = 2NaCl + H 2

2Na + 2NaOH(l) = 2Na 2 O + H 2

2Na + H 2 = 2NaH

2Na + Hal 2 = 2NaHal (room, Hal = F, Cl; 150-200° C, Hal = Br, I)

2Na + NH 3 (g) = 2NaNH 2 + H 2

Sodium forms intermetallic compounds with many metals. Thus, with tin it gives a number of compounds: NaSn 6, NaSn 4, NaSn 3, NaSn 2, NaSn, Na 2 Sn, Na 3 Sn, etc.; gives solid solutions with some metals.

Sodium is stored in sealed containers or under a layer of kerosene.

Preparation and use of sodium. Sodium is obtained by electrolysis of molten NaCl and, less commonly, NaOH. In the electrolytic reduction of NaCl, a eutectic mixture is used, for example, NaCl-KCl (melting point is almost 300°C lower than the melting point of NaCl).

2NaCl(l) = 2Na + Cl 2 (electric current)

Sodium is used in metallothermy, organic synthesis, nuclear power plants (as a coolant), aircraft engine valves, chemical industries, where uniform heating is required within the range of 450-650 ° C.

Sodium compounds (I). The most characteristic are ionic compounds with a crystalline structure, characterized by refractory properties and are highly soluble in water. Some derivatives with complex anions are sparingly soluble, such as Na hexahydroxostibate (V); NaHCO 3 is slightly soluble (unlike carbonate).

When interacting with oxygen, sodium (unlike lithium) forms not an oxide, but a peroxide: 2Na + O 2 = Na 2 O 2

Sodium oxide Na 2 O is obtained by reducing Na 2 O 2 with sodium metal. Low-resistant ozonide NaO 3 and sodium superoxide NaO 2 are also known.

Of the sodium compounds, its chloride, hydroxide, carbonates and numerous other derivatives are important.

Sodium chloride NaCl is the basis for a number of important industries, such as the production of sodium, caustic soda, soda, chlorine, etc.

Sodium hydroxide ( caustic soda, caustic soda) NaOH is a very strong base. It is used in a variety of industries, the main of which are the production of soaps, paints, cellulose, etc. NaOH is obtained by electrolysis of aqueous solutions of NaCl and chemical methods. Thus, the lime method is common - the interaction of a solution of sodium carbonate (soda) with calcium hydroxide (slaked lime):

Na 2 CO 3 + Ca(OH) 2 = 2NaOH + CaCO 3

Sodium carbonates Na 2 CO 3 ( soda ash), Na 2 CO 3 ∙10H 2 O ( crystal soda), NaHCO 3 ( baking soda) are used in the chemical, soap, paper, textile, and food industries.

Potassium subgroup(potassium, rubidium, cesium, francium)

Elements of the potassium subgroup are the most typical metals. They are most characterized by compounds with a predominantly ionic type of bond. Complexation with inorganic ligands for K + , Rb + , Cs + is uncharacteristic.

The most important potassium minerals are: sylvin KCl, sylvinite NaCl∙KCl, carnallite KCl∙MgCl 2 ∙6H 2 O, Cainite KCl∙MgSO 4 ∙3H 2 O. Potassium (together with sodium) is part of living organisms and all silicate rocks. Rubidium and cesium are found in potassium minerals. Francium is radioactive and has no stable isotopes (the longest-lived isotope is Fr with a half-life of 22 minutes).

Simple substances. K (potassium) – silvery-white, soft, low-melting alkali metal. Extremely reactive, powerful reducing agent; reacts with air O 2, water (ignition of the released H 2 occurs), dilute acids, non-metals, ammonia, hydrogen sulfide, molten potassium hydroxide. Practically does not react with nitrogen (unlike lithium and sodium). Forms intermetallic compounds with Na, Tl, Sn, Pb and Bi. Colors the flame of a gas burner purple.

Rb (rubidium) – white, soft, very low-melting alkali metal. Extremely reactive; the strongest reducing agent; reacts vigorously with air O 2, water (ignition of the metal and released H 2 occurs), dilute acids, non-metals, ammonia, hydrogen sulfide. Does not react with nitrogen. Colors the flame of a gas burner purple.

Cs (cesium) – white (light yellow when cut), soft, very low-melting alkali metal. Extremely reactive, powerful reducing agent; reacts with air O 2, water (ignition of the metal and the released H 2 occurs), dilute acids, non-metals, ammonia, hydrogen sulfide. Does not react with nitrogen. Colors the flame of a gas burner blue.

Fr (French) – white, very low-melting alkali metal. Radioactive. The most reactive of all metals, its chemical behavior is similar to cesium. In air it becomes covered with an oxide film. Strong reducing agent; Reacts vigorously with water and acids, releasing H2. The francium compounds FrClO 4 and Fr 2 were isolated by precipitation with the corresponding poorly soluble salts Rb and Cs.

Potassium and its analogues are stored in sealed containers, as well as under a layer of paraffin or petroleum jelly. Potassium, in addition, is well preserved under a layer of kerosene or gasoline.

Receipt and application. Potassium is obtained by electrolysis of molten KCl and by the sodium thermal method from molten potassium hydroxide or chloride. Rubidium and cesium are most often obtained by vacuum-thermal reduction of their chlorides with calcium metal. All alkali metals can be easily purified by sublimation in a vacuum.

Metals of the potassium subgroup lose electrons relatively easily when heated and illuminated, and this ability makes them a valuable material for the manufacture of solar cells.

Compounds of potassium (I), rubidium (I), cesium (I). Derivatives of potassium and its analogues are mainly salts and salt-like compounds. In composition, crystal structure, solubility and the nature of solvolysis, their compounds show great similarity with similar sodium compounds.

In accordance with the increase in chemical activity in the K–Rb–Cs series, the tendency to form peroxide compounds increases. Thus, upon combustion they form superoxides EO 2. Peroxides E 2 O 2 and ozonides EO 3 can also be obtained indirectly. Peroxides, superoxides and ozonides are strong oxidizing agents, easily decomposed by water and dilute acids:

2KO 2 + 2H 2 O = 2KON + H 2 O 2 + O 2

2KO 2 + 2НCl = 2КCl + Н 2 О 2 + О 2

4KO 3 + 2H 2 O = 4KON + 5O 2

EON hydroxides are the strongest bases (alkalis); when heated, like NaOH, they sublimate without decomposition. When dissolved in water, a significant amount of heat is released. The most important in technology is KOH (caustic potash), obtained by electrolysis of an aqueous solution of KCl.

In contrast to similar compounds Li + and Na +, their oxochlorates (VII) EOCl 4, chloroplatinates (IV) E 2 PlCl 6, nitrite cobaltates (III) E 3 [Co(NO 2) 6] and some others are sparingly soluble.

Of the derivatives of the subgroup, potassium compounds are of greatest importance. About 90% of potassium salts are consumed as fertilizer. Its compounds are also used in the production of glass and soap.

Copper subgroup(copper, silver, gold)

For copper, the most typical compounds are with oxidation states +1 and +2, for gold +1 and +3, and for silver +1. All of them have a pronounced tendency to form complexes.

All elements of group IB are relatively rare. The most important of the natural copper compounds are the following minerals: copper pyrite (chalcopyrite) CuFeS 2 , copper shine Cu 2 S, as well as cuprite Cu 2 O, malachite CuCO 3 ∙Cu(OH) 2, etc. Silver is part of sulfide minerals of other metals (Pd, Zn, Cd, etc.). For Cu, Ag and Au, arsenide, stibide and sulfidarsenide minerals are also quite common. Copper, silver and especially gold are found in nature in a native state.

All soluble compounds of copper, silver and gold are poisonous.

Simple substances. Si (copper) – red, soft, malleable metal. Does not change in air in the absence of moisture and CO 2; when heated, it becomes dull (formation of an oxide film). Weak reducing agent (noble metal); does not react with water. It is transferred into solution with non-oxidizing acids or ammonia hydrate in the presence of O 2, potassium cyanide. It is oxidized by concentrated sulfuric and nitric acids, aqua regia, oxygen, halogens, chalcogens, and metal oxides. Reacts when heated with hydrogen halides.

Cu + H 2 SO 4 (conc., horizontal) = CuSO 4 + SO 2 + H 2 O

Cu + 4HNO 3 (conc.) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O

3Cu + 8HNO 3 (diluted) = 3Cu(NO 3) 2 + 2NO + 4H 2 O

2Cu + 4НCl(diluted) + O 2 = 2CuCl 2 + 2Н 2 O

Cu + Cl 2 (humidity, room) = CuCl 2

2Cu + O 2 (load) = 2CuO

Cu + 4KCN(conc.) + H 2 O = 2K + 2KOH + H 2

4Cu + 2O2 + 8NH3 + 2H2O = 4OH

2Cu + CO 2 + O 2 + H 2 O = Cu 2 CO 3 (OH) 2 ↓

Ag (silver) – white, heavy, ductile metal. Low-active (noble metal); does not react with oxygen, water, dilute hydrochloric and sulfuric acids. Weak reducing agent; reacts with oxidizing acids. Turns black in the presence of moist H2S.

Ag + 2H 2 SO 4 (conc., horizontal) = Ag 2 SO 4 ↓ + SO 2 + H 2 O

3Ag + 4HNO 3 (diluted) = 3AgNO 3 + NO + 2H 2 O

4Ag + H 2 S + O 2 (air) = 2Ag 2 S + 2H 2 O

2Ag + Hal 2 (load) = 2AgHal

4Ag + 8KCN + 2H2O + O2 = 4K + 4KOH

Ai (gold) – yellow, malleable, heavy, high-melting metal. Stable in dry and humid air. Noble metal; does not react with water, non-oxidizing acids, concentrated sulfuric and nitric acids, alkalis, ammonia hydrate, oxygen, nitrogen, carbon, sulfur. Does not form simple cations in solution. Transferred into solution "royal vodka", mixtures of halogens and hydrohalic acids, oxygen in the presence of alkali metal cyanides. Oxidized by sodium nitrate during fusion, krypton difluoride.

Au + HNO 3 (conc.) + 4HCl (conc.) = H + NO + 2H 2 O

2Au + 6H 2 SeO 4 (conc., horizontal) = Au 2 (SeO 4) 3 + 3SeO 2 + 6H 2 O

2Au + 3Cl 2 (up to 150°C) = 2AuCl 3

2Au + Cl 2 (150-250°C) = 2AuCl

Au + 3Hal + 2HAl(conc.) = H + NO + 2H 2 O (Hal = Cl, Br, I)

4Au + 8NaCN + 2H 2 O + O 2 = 4Na + 4KOH

Au + NaN0 3 = NaAuО 2 + NO

Receipt and application. Copper is obtained by pyrometallurgical reduction of oxidized sulfide concentrates. The sulfur dioxide SO2 released during the roasting of sulfides is used for the production of sulfuric acid, and the slag is used for the production of slag concrete, stone casting, slag wool, etc. Recovered blister copper is purified by electrochemical refining. Precious metals, selenium, tellurium, etc. are extracted from the anode slurry. Silver is obtained by processing polymetallic (silver-lead-zinc) sulfide ores. After oxidative roasting, zinc is distilled off, copper is oxidized, and rough silver is subjected to electrochemical refining. In the cyanide method of gold mining, the gold-bearing rock is first washed with water, then treated with a NaCN solution in air; in this case, gold forms a Na complex, from which it is precipitated with zinc:

Na + Zn = Na 2 + 2Au↓

This method can also be used to isolate silver from low-grade ores. In the mercury method, gold-bearing rock is treated with mercury to obtain amalgams gold, then the mercury is distilled off.

Cu, Ag and Au form alloys with each other and with many other metals. Of the copper alloys, the most important are bronze(90% Cu, 10% Sn), red brass(90% Cu, 10% Zn), cupronickel(68% Cu, 30% Ni, 1% Mn, 1% Fe), nickel silver(65% Cu, 20% Zn, 15% Ni), brass(60% Cu, 40% Zn), as well as coin alloys.

Due to its high thermal and electrical conductivity, malleability, good casting qualities, high tensile strength and chemical resistance, copper is widely used in industry, electrical engineering, and mechanical engineering. Copper is used to make electrical wires and cables, various industrial equipment (boilers, distillation cubes, etc.)

Due to their softness, silver and gold are usually alloyed with other metals, most often with copper. Silver alloys are used for the manufacture of jewelry and household items, coins, radio components, silver-zinc batteries, and in medicine. Gold alloys are used for electrical contacts, dental prosthetics, and jewelry.

Compounds of copper (I), silver (I) and gold (I). The +1 oxidation state is most characteristic of silver; In copper and, especially, in gold, this oxidation state appears less frequently.

Binary compounds Cu (I), Ag (I) and Au (I) are solid crystalline salt-like substances, mostly insoluble in water. Derivatives of Ag (I) are formed by the direct interaction of simple substances, and Cu (I) and Au (I) are formed by the reduction of the corresponding compounds Cu (II) and Au (III).

Ammino complexes of the [E(NH 3) 2 ] + type are stable for Cu (I) and Ag (I), and therefore most Cu (I) and Ag (I) compounds dissolve quite easily in the presence of ammonia, as follows:

CuCl + 2NH 3 = Cl

Ag 2 O + 4NH 3 + H 2 O = 2(OH)

Hydroxides of the [E(NH 3) 2 ](OH) type are much more stable than EON and are close in strength to alkalis. EON hydroxides are unstable, and when trying to obtain them through exchange reactions, the oxides CuO (red), Ag 2 O (dark brown) are released, as follows:

2AgNO 3 + 2NaOH = Ag 2 O + 2NaNO 3 + H 2 O

Oxides of E 2 O exhibit acidic properties when interacting with the corresponding basic compounds, cuprates (I), argentates (I) and aurates (I) are formed.

Cu 2 O + 2NaOH (conc.) + H 2 O = 2Na

EHal halides, insoluble in water and acids, dissolve quite significantly in solutions of hydrohalic acids or basic halides:

CuCl + HC1 = H AgI + KI = K

Water-insoluble cyanides ECN, sulfides E2S, etc. behave similarly.

Most Cu (I) and Au (I) compounds are easily oxidized (even by atmospheric oxygen), turning into stable derivatives Cu (II) and Au (III).

4CuCl + O 2 + 4HCl = 4CuCl 2 + 2H 2 O

For connections. Cu(I) and Au(I) are characterized by disproportionation:

2CuC1 = CuCl 2 + Cu

3AuCl + KCl = K + 2Au

Most E(I) compounds easily decompose with slight heating and exposure to light, so they are usually stored in dark glass jars. The photosensitivity of silver halides is used to prepare photosensitive emulsions. Copper (I) oxide is used for coloring glass, enamels, and also in semiconductor technology.

Copper(II) compounds . The oxidation state +2 is characteristic only of copper. When Cu(II) salts are dissolved in water or when CuO (black) and Cu(OH)2 (blue) interact with acids, blue aqua complexes 2+ are formed. Most crystalline hydrates have the same color, for example, Cu(NO 3) 2 ∙6H 2 O; Crystalline hydrates of Cu(II) are also found, having a green and dark brown color.

When ammonia acts on solutions of copper (II) salts, ammonia compounds are formed:

Cu(OH) 2 ↓ + 4NH 3 + 2H 2 = (OH) 2

Copper(II) is also characterized by anionic complexes – cuprates(II). Thus, Cu(OH) 2, when heated in concentrated alkali solutions, partially dissolves, forming blue hydroxocuprates (II) type M 2 +1. In aqueous solutions, hydroxocuprates (II) easily decompose.

In an excess of basic halides, CuHal 2 forms halogenocuprates (II) of type M +1 and M 2 +1 [CuHal 4 ]. Anionic complexes of Cu (II) with cyanide, carbonate, sulfate and other anions are also known.

Of the copper (II) compounds, the most technically important is the crystalline hydrate CuSO 4 ∙5H 2 O ( copper sulfate) is used to produce paints, to control pests and plant diseases, serves as a starting product for the production of copper and its compounds, etc.

Compounds of copper (III), silver (III), gold (III). The oxidation state +3 is most characteristic of gold. Compounds of copper (III) and silver (III) are unstable and are strong oxidizing agents.

The starting product for the preparation of many gold compounds is AuCl 3, which is obtained by reacting Au powder with excess Cl 2 at 200°C.

Au(III) halides, oxide and hydroxide are amphoteric compounds with predominant acidic properties.

NaOH + Au(OH) 3 = Na

Au(OH) 3 + 4HN0 3 = H + 3H 2 O

AuHal 3 + M +1 Hal = M

Hydrogen nitrate and cyanoaurates (III) are isolated in a free state. In the presence of alkali metal salts, aurates are formed, for example: M +1, M +1, etc.

Gold (V) and (VII) compounds. The interaction of gold and krypton(II) fluoride produced gold pentafluoride AuF 5:

2Au + 5KrF 2 = 2AuF 5 + 5Kr

Pentafluoride AuF 5 exhibits acidic properties and forms fluoroaurates (V) with basic fluorides.

NaF + AuF5 = Na

Au(V) compounds are very strong oxidizing agents. Thus, AuF 5 oxidizes even XeF 2:

AuF 5 + XeF 2 = XeF 4 + AuF 3

Compounds such as XeFAuF 6, XeF 5 AuF 6 and some others are also known.

An extremely unstable fluoride AuF 7 is known.

Hydrogen H is a chemical element, one of the most common in our Universe. The mass of hydrogen as an element in the composition of substances is 75% of the total content of atoms of other types. It is part of the most important and vital compound on the planet - water. A distinctive feature of hydrogen is also that it is the first element in D.I. Mendeleev’s periodic system of chemical elements.

Discovery and exploration

The first mention of hydrogen in the writings of Paracelsus dates back to the sixteenth century. But its isolation from the gas mixture of air and the study of flammable properties were carried out already in the seventeenth century by the scientist Lemery. Hydrogen was thoroughly studied by an English chemist, physicist and natural scientist who experimentally proved that the mass of hydrogen is the smallest in comparison with other gases. In subsequent stages of the development of science, many scientists worked with him, in particular Lavoisier, who called him “the birther of water.”

Characteristics by position in PSHE

The element that opens the periodic table of D.I. Mendeleev is hydrogen. The physical and chemical properties of the atom show a certain duality, since hydrogen is simultaneously classified as belonging to the first group, the main subgroup, if it behaves like a metal and gives up a single electron in the process of a chemical reaction, and to the seventh - in the case of complete filling of the valence shell, that is, acceptance negative particle, which characterizes it as similar to halogens.

Features of the electronic structure of the element

The properties of the complex substances in which it is included, and of the simplest substance H2, are primarily determined by the electronic configuration of hydrogen. The particle has one electron with Z= (-1), which rotates in its orbit around a nucleus containing one proton with unit mass and a positive charge (+1). Its electronic configuration is written as 1s 1, which means the presence of one negative particle in the very first and only s-orbital for hydrogen.

When an electron is removed or given up, and an atom of this element has such a property that it is related to metals, a cation is obtained. In essence, the hydrogen ion is a positive elementary particle. Therefore, hydrogen deprived of an electron is simply called a proton.

Physical properties

To describe hydrogen briefly, it is a colorless, slightly soluble gas with a relative atomic mass of 2, 14.5 times lighter than air, with a liquefaction temperature of -252.8 degrees Celsius.

From experience you can easily verify that H 2 is the lightest. To do this, it is enough to fill three balls with various substances - hydrogen, carbon dioxide, ordinary air - and simultaneously release them from your hand. The one filled with CO 2 will reach the ground the fastest, after it the one inflated with the air mixture will descend, and the one containing H 2 will rise to the ceiling.

The small mass and size of hydrogen particles justify its ability to penetrate various substances. Using the example of the same ball, it is easy to verify this; after a couple of days it will deflate on its own, since the gas will simply pass through the rubber. Hydrogen can also accumulate in the structure of some metals (palladium or platinum), and evaporate from it when the temperature rises.

The property of low solubility of hydrogen is used in laboratory practice to isolate it by displacing hydrogen (the table shown below contains the main parameters) to determine the scope of its application and methods of production.

Isotopic composition

Like many other representatives of the periodic system of chemical elements, hydrogen has several natural isotopes, that is, atoms with the same number of protons in the nucleus, but a different number of neutrons - particles with zero charge and unit mass. Examples of atoms with a similar property are oxygen, carbon, chlorine, bromine and others, including radioactive ones.

The physical properties of hydrogen 1H, the most common of the representatives of this group, differ significantly from the same characteristics of its counterparts. In particular, the characteristics of the substances they contain differ. Thus, there is ordinary and deuterated water, which contains, instead of a hydrogen atom with a single proton, deuterium 2 H - its isotope with two elementary particles: positive and uncharged. This isotope is twice as heavy as ordinary hydrogen, which explains the dramatic difference in the properties of the compounds they make up. In nature, deuterium is found 3200 times less frequently than hydrogen. The third representative is tritium 3H; it has two neutrons and one proton in its nucleus.

Methods of production and isolation

Laboratory and industrial methods are quite different. Thus, gas is produced in small quantities mainly through reactions involving mineral substances, while large-scale production uses organic synthesis to a greater extent.

The following chemical interactions are used in the laboratory:

For industrial purposes, gas is produced by the following methods:

1. Thermal decomposition of methane in the presence of a catalyst to its constituent simple substances (the value of such an indicator as temperature reaches 350 degrees) - hydrogen H2 and carbon C.
2. Passing steamy water through coke at 1000 degrees Celsius to form carbon dioxide CO 2 and H 2 (the most common method).
3. Conversion of methane gas on a nickel catalyst at temperatures reaching 800 degrees.
4. Hydrogen is a by-product from the electrolysis of aqueous solutions of potassium or sodium chlorides.

Chemical interactions: general provisions

The physical properties of hydrogen largely explain its behavior in reaction processes with a particular compound. The valency of hydrogen is 1, since it is located in the first group in the periodic table, and the degree of oxidation varies. In all compounds, except hydrides, hydrogen in d.o. = (1+), in molecules of the type CN, CN 2, CN 3 - (1-).

The hydrogen gas molecule, formed by creating a generalized electron pair, consists of two atoms and is quite stable energetically, which is why under normal conditions it is somewhat inert and reacts when normal conditions change. Depending on the degree of oxidation of hydrogen in the composition of other substances, it can act as both an oxidizing agent and a reducing agent.

Substances with which hydrogen reacts and forms

Elemental interactions to form complex substances (often at elevated temperatures):

1. Alkali and alkaline earth metal + hydrogen = hydride.
2. Halogen + H 2 = hydrogen halide.
3. Sulfur + hydrogen = hydrogen sulfide.
4. Oxygen + H 2 = water.
5. Carbon + hydrogen = methane.
6. Nitrogen + H 2 = ammonia.

Interaction with complex substances:

1. Production of synthesis gas from carbon monoxide and hydrogen.
2. Reduction of metals from their oxides using H 2.
3. Saturation of unsaturated aliphatic hydrocarbons with hydrogen.

Hydrogen bond

The physical properties of hydrogen are such that they allow it, when in combination with an electronegative element, to form a special type of bond with the same atom from neighboring molecules that have lone electron pairs (for example, oxygen, nitrogen and fluorine). The clearest example in which it is better to consider this phenomenon is water. It can be said to be stitched with hydrogen bonds, which are weaker than covalent or ionic ones, but due to the fact that there are many of them, they have a significant impact on the properties of the substance. Essentially, hydrogen bonding is an electrostatic interaction that binds water molecules into dimers and polymers, giving rise to its high boiling point.

Hydrogen in mineral compounds

All contain a proton, a cation of an atom such as hydrogen. A substance whose acidic residue has an oxidation state greater than (-1) is called a polybasic compound. It contains several hydrogen atoms, which makes dissociation in aqueous solutions multi-stage. Each subsequent proton becomes more and more difficult to remove from the acid residue. The acidity of the medium is determined by the quantitative content of hydrogen in the medium.

Application in human activities

Cylinders with the substance, as well as containers with other liquefied gases, such as oxygen, have a specific appearance. They are painted dark green with the word “Hydrogen” written in bright red. Gas is pumped into a cylinder under a pressure of about 150 atmospheres. The physical properties of hydrogen, in particular the lightness of the gaseous state of aggregation, are used to fill balloons, balloons, etc. with it mixed with helium.

Hydrogen, the physical and chemical properties of which people learned to use many years ago, is currently used in many industries. The bulk of it goes to the production of ammonia. Hydrogen also participates in (hafnium, germanium, gallium, silicon, molybdenum, tungsten, zirconium and others) oxides, acting in the reaction as a reducing agent, hydrocyanic and hydrochloric acids, as well as artificial liquid fuel. The food industry uses it to convert vegetable oils into solid fats.

The chemical properties and use of hydrogen in various processes of hydrogenation and hydrogenation of fats, coals, hydrocarbons, oils and fuel oil were determined. It is used to produce precious stones, incandescent lamps, and forge and weld metal products under the influence of an oxygen-hydrogen flame.

HYDROGEN, H (lat. hydrogenium; a. hydrogen; n. Wasserstoff; f. hydrogene; i. hidrogeno), is a chemical element of the periodic system of Mendeleev’s elements, which is simultaneously classified as groups I and VII, atomic number 1, atomic mass 1, 0079. Natural hydrogen has stable isotopes - protium (1 H), deuterium (2 H, or D) and radioactive - tritium (3 H, or T). For natural compounds, the average ratio D/H = (158±2).10 -6 The equilibrium content of 3 H on Earth is ~5.10 27 atoms.

Physical properties of hydrogen

Hydrogen was first described in 1766 by the English scientist G. Cavendish. Under normal conditions, hydrogen is a colorless, odorless, and tasteless gas. In nature, it is found in the free state in the form of H2 molecules. The dissociation energy of the H 2 molecule is 4.776 eV; ionization potential of the hydrogen atom is 13.595 eV. Hydrogen is the lightest substance known, at 0°C and 0.1 MPa 0.0899 kg/m 3 ; boiling t - 252.6°C, melting t - 259.1°C; critical parameters: t - 240°C, pressure 1.28 MPa, density 31.2 kg/m 3. The most thermally conductive of all gases is 0.174 W/(m.K) at 0°C and 1 MPa, specific heat capacity 14.208.10 3 J(kg.K).

Chemical properties of hydrogen

Liquid hydrogen is very light (density at -253°C is 70.8 kg/m 3) and fluid (at -253°C it is 13.8 cP). In most compounds, hydrogen exhibits an oxidation state of +1 (similar to alkali metals), less often -1 (similar to metal hydrides). Under normal conditions, molecular hydrogen is inactive; solubility in water at 20°C and 1 MPa 0.0182 ml/g; highly soluble in metals - Ni, Pt, Pd, etc. With oxygen it forms water with the release of heat 143.3 MJ/kg (at 25°C and 0.1 MPa); at 550°C and above the reaction is accompanied by an explosion. When interacting with fluorine and chlorine, reactions also occur explosively. The main hydrogen compounds: H 2 O, ammonia NH 3, hydrogen sulfide H 2 S, CH 4, metal and halogen hydrides CaH 2, HBr, Hl, as well as organic compounds C 2 H 4, HCHO, CH 3 OH, etc.

Hydrogen in nature

Hydrogen is a widespread element in nature, its content is 1% (by weight). The main reservoir of hydrogen on Earth is water (11.19%, by mass). Hydrogen is one of the main components of all natural organic compounds. In a free state, it is present in volcanic and other natural gases, in (0.0001%, by number of atoms). It makes up the bulk of the mass of the Sun, stars, interstellar gas, and gas nebulae. In the atmospheres of planets it is present in the form of H 2, CH 4, NH 3, H 2 O, CH, NHOH, etc. It is part of the corpuscular radiation of the Sun (proton flows) and cosmic rays (electron flows).

Production and use of hydrogen

Raw materials for the industrial production of hydrogen are oil refinery gases, gasification products, etc. The main methods for producing hydrogen are: the reaction of hydrocarbons with water vapor, partial oxidation of hydrocarbons, oxide conversion, electrolysis of water. Hydrogen is used for the production of ammonia, alcohols, synthetic gasoline, hydrochloric acid, hydrotreating of petroleum products, and cutting metals with a hydrogen-oxygen flame.

Hydrogen is a promising gaseous fuel. Deuterium and tritium have found application in nuclear energy.