Chemical elements of the main subgroup IA of group of the periodic system of elements of D. I. Mendeleev: Li, Na, K, Rb, Cs, Fr. The name comes from hydroxides of alkali metals, called caustic alkalis. Atoms of alkali metals have 1 s-electron on the outer shell, and 2 s- and 6 p-electrons on the previous shell (except for Li). Characterized by low melting temperatures, low densities; soft, cut with a knife. The oxidation state of alkali metals in compounds is always +1. These metals are chemically very active - they are quickly oxidized by atmospheric oxygen, react violently with water, forming alkalis Me. OH (where Me is metal); activity increases from Li to Fr.
Lithium (Latin - lithium), a Li-chemical element of the first group, A-subgroup of the periodic system of D.I. Mendeleev, belongs to the alkali metals, serial number 3, atomic mass is 6.939; under normal conditions, a silvery-white, lightweight metal. Natural lithium consists of two isotopes with mass numbers 6 and 7. An interesting detail: the cost of lithium isotopes is not at all proportional to their abundance. At the beginning of this decade in the United States, relatively pure lithium-7 was almost 10 times more expensive than very high-purity lithium-6. Two more lithium isotopes have been obtained artificially. Their lifetime is extremely short: lithium-8 has a half-life of 0.841 seconds, and lithium-9 has a half-life of 0.168 seconds.
Lithium is a typical element of the earth's crust, a relatively rare element. (content 3.2× 10 -3% by weight), it accumulates in the most recent products of magma differentiation - pegmatites. There is little lithium in the mantle - in ultramafic rocks only 5 × 10 -3% (in basic rocks 1.5 × 10 -3%, in intermediate rocks - 2 × 103%, in acidic rocks 4 × 10 -3%). The proximity of the ionic radii of Li+, Fe 2+ and Mg 2+ allows lithium to enter the lattices of magnesium-iron silicates - pyroxenes and amphiboles. In granitoids it is contained as an isomorphic impurity in micas. Only 28 independent lithium minerals (silicates, phosphates, etc.) are known in pegmatites and in the biosphere. They are all rare. In the biosphere, lithium migrates relatively weakly, its role in living matter is less than that of other alkali metals. It is easily extracted from waters by clays; there is relatively little of it in the World Ocean (1.5 × 10 -5%). In the human body (weighing 70 kg) - 0.67 mg. lithium.
Potassium (Kalium) Potassium is a chemical element of group I of the periodic system of Mendeleev; atomic number 19, atomic mass 39.098; silver-white, very light, soft and fusible metal. The element consists of two stable isotopes - 39 K (93.08%), 41 K (6.91%) and one weakly radioactive 40 K (0.01%) with a half-life of 1.32×109 years.
Occurrence in nature Potassium is a common element: its content in the lithosphere is 2.50% by mass. In magmatic processes, potassium, like sodium, accumulates in acidic magmas, from which granites and other rocks crystallize (average potassium content 3.34%). Potassium is found in feldspars and micas. Basic and ultrabasic rocks rich in iron and magnesium are low in potassium. On the earth's surface, potassium, unlike sodium, migrates weakly. When rocks weather, potassium partially passes into water, but from there it is quickly captured by organisms and absorbed by clays, so river waters are poor in potassium and much less of it enters the ocean than sodium. In the ocean, potassium is absorbed by organisms and bottom silts (for example, it is part of glauconite); Therefore, ocean waters contain only 0.038% potassium - 25 times less than sodium.
In nature, it is the ninth most abundant element (sixth among metals), found only in the form of compounds. It is part of many minerals, rocks, and salt layers. The third most abundant metal in natural waters: 1 liter of sea water contains 0.38 g of K+ ions. Potassium cations are well adsorbed by soil and are difficult to wash out with natural waters. A vital element for all organisms. K+ ions are always found inside cells (unlike Na+ ions). The human body contains about 175 g of potassium, the daily requirement is about 4 g. The lack of potassium in the soil is compensated by applying potassium fertilizers - potassium chloride KCl, potassium sulfate K 2 SO 4 and plant ash.
Interesting facts WHAT IS POTASSIUM CYANIDE NEEDED FOR? For extracting gold and silver from ores. For galvanic gilding and silvering of base metals. For obtaining many organic substances. For nitriding steel - this gives its surface greater strength. Unfortunately, this much-needed substance is extremely poisonous. And KCN looks quite harmless: small white crystals with a brownish or gray tint.
Cesium Cesium was discovered relatively recently, in 1860, in the mineral waters of famous healing springs in the Black Forest (Baden-Baden, etc.). In a short historical period, it has gone through a brilliant path - from a rare, unknown chemical element to a strategic metal. He belongs to the working family of alkali metals, but the blue blood of the last of his kind flows in his veins. . . However, this does not in the least prevent him from communicating with other elements, and even if they are not so famous, he willingly comes into contact with them and establishes strong connections. Currently, he works simultaneously in several industries: in electronics and automation, in radar and cinema, in nuclear reactors and on spaceships. . . ".
Cesium is known to have been the first element discovered by spectral analysis. Scientists, however, had the opportunity to become familiar with this element even before Robert Bunsen and Gustav Kirchhoff created a new research method. In 1846, the German chemist Plattner, analyzing the mineral pollucite, discovered that the sum of its known components was only 93%, but was unable to accurately determine what other element (or elements) was included in this mineral. Only in 1864, after the discovery of Bunsen, the Italian Pisani found cesium in pollucite and established that it was the compounds of this element that Plattner could not identify.
Interesting Facts Cesium and Pressure All alkali metals change greatly when exposed to high pressure. But it is cesium that reacts to it most uniquely and sharply. At a pressure of 100 thousand atm. its volume decreases almost three times - more than that of other alkali metals. In addition, it was under high pressure conditions that two new modifications of elemental cesium were discovered. The electrical resistance of all alkali metals increases with increasing pressure; In cesium this property is especially pronounced.
Francium Among the elements at the end of D.I. Mendeleev’s periodic table, there are those about which non-specialists have heard and know a lot, but there are also those about which even a chemist can tell little. The former include, for example, radon (No. 86) and radium (No. 88). Among the second is their neighbor in the periodic table, element No. 87 - francium. Francium is interesting for two reasons: firstly, it is the heaviest and most active alkali metal; secondly, francium can be considered the most unstable of the first hundred elements of the periodic table. The longest-lived isotope, francium, 223 Fr, has a half-life of only 22 minutes. Such a rare combination in one element of high chemical activity with low nuclear stability determined the difficulties in the discovery and study of this element.
Occurrence in nature In addition to 223 Fr, several isotopes of element No. 87 are now known. But only 223 Fr occurs in nature in any noticeable quantities. Using the law of radioactive decay, it can be calculated that a gram of natural uranium contains 4·10–18 g of 223 Fr. This means that about 500 g of France-223 is in radioactive equilibrium with the entire mass of earthly uranium. There are two more isotopes of element No. 87 in vanishingly small quantities on Earth - 224 Fr (a member of the radioactive thorium family) and 221 Fr. Naturally, it is almost impossible to find an element on Earth whose global reserves do not reach a kilogram. Therefore, all studies of francium and its few compounds were performed on artificial products.
Interesting facts Sodium on a submarine Sodium melts at 98°C, but boils only at 883°C. Consequently, the temperature range of the liquid state of this element is quite large. That is why (and also due to the small neutron capture cross section) sodium began to be used in nuclear energy as a coolant. In particular, American nuclear submarines are equipped with power plants with sodium circuits. The heat generated in the reactor heats the liquid sodium, which circulates between the reactor and the steam generator. In a steam generator, sodium, when cooled, evaporates water, and the resulting high-pressure steam rotates a steam turbine. For the same purposes, an alloy of sodium and potassium is used.
Inorganic photosynthesis Usually, the oxidation of sodium produces an oxide of the composition Na 2 O. However, if sodium is burned in dry air at elevated temperatures, then instead of the oxide, peroxide N 2 O 2 is formed. This substance easily gives up its “extra” oxygen atom and therefore has strong oxidizing properties . At one time, sodium peroxide was widely used to bleach straw hats. Now the proportion of straw hats in the use of sodium peroxide is negligible; The main quantities of it are used for bleaching paper and for air regeneration in submarines. When sodium peroxide interacts with carbon dioxide, the opposite process to respiration occurs: 2 Na 2 O 2 + 2 CO 2 → 2 Na 2 CO 3 + O 2, i.e. carbon dioxide is bound and oxygen is released. Just like a green leaf!
Sodium and gold By the time sodium was discovered, alchemy was no longer in favor, and the idea of turning sodium into gold did not excite the minds of natural scientists. However, now a lot of sodium is consumed to obtain gold. “Gold ore” is treated with a solution of sodium cyanide (and it is obtained from elemental sodium). In this case, gold is converted into a soluble complex compound, from which it is isolated with the help of zinc. Gold miners are among the main consumers of element No. 11. On an industrial scale, sodium cyanide is produced by the reaction of sodium, ammonia and coke at a temperature of about 800°C.
Sodium in water Every schoolchild knows what happens if you throw a piece of sodium into water. More precisely, not into water, but onto water, because sodium is lighter than water. The heat released when sodium reacts with water is enough to melt the sodium. And now a sodium ball runs through the water, driven by the released hydrogen. However, the reaction of sodium with water is not only dangerous fun; on the contrary, it is often useful. Sodium is used to reliably remove traces of water from transformer oils, alcohols, ethers and other organic substances, and with the help of sodium amalgam (i.e., an alloy of sodium with mercury), the moisture content in many compounds can be quickly determined. Amalgam reacts with water much more calmly than sodium itself. To determine moisture content, a certain amount of sodium amalgam is added to a sample of organic matter and the moisture content is determined by the volume of hydrogen released.
Sodium Belt of the Earth It is quite natural that sodium is never found in a free state on Earth - this metal is too active. But in the upper layers of the atmosphere - at an altitude of about 80 km - a layer of atomic sodium was discovered. At this altitude there is virtually no oxygen, water vapor, or anything at all for sodium to react with. Sodium was also discovered in interstellar space using spectral methods.
Rubidium is a metal that can be called chemically touchy. Upon contact with air, it spontaneously ignites and burns with a bright pinkish-violet flame. It explodes with water and also reacts violently on contact with fluorine, chlorine, bromine, iodine, and sulfur. As a true touch-me-not, rubidium must be protected from external influences. For this purpose, it is placed in vessels filled with dry kerosene. . . Rubidium is heavier than kerosene (density of rubidium 1.5) and does not react with it. Rubidium is a radioactive element and slowly releases a stream of electrons to become strontium. The most remarkable property of rubidium is its peculiar sensitivity to light. Under the influence of light rays, rubidium becomes a source of electric current. With the cessation of light irradiation, the current also disappears. R. reacts with water explosively, and hydrogen is released and a solution of R. hydroxide, Rb, is formed. OH.
Rubidium is found in many rocks and minerals, but its concentration is extremely low. Only lepidolites contain slightly more Rb 2 O, sometimes 0.2%, and occasionally up to 1. . . 3%. Rubidium salts are dissolved in the water of seas, oceans and lakes. Their concentration here is very low, on average about 100 µg/l. This means that there is hundreds of times less rubidium in the world's oceans than in the earth's crust.
Interesting facts Rubidium has not ignored many representatives of the plant world: traces of it are found in seaweed and tobacco, in tea leaves and coffee beans, in sugar cane and beets, in grapes and some types of citrus fruits. Why was it called rubidium? Rubidus – Latin for “red”. It would seem that this name is more suitable for copper than for rubidium, which is very ordinary in color. But let's not rush to conclusions. This name was given to element No. 37 by its discoverers Kirchhoff and Bunsen. More than a hundred years ago, while studying various minerals with a spectroscope, they noticed that one of the lepidolite samples sent from Rosen (Saxony) gave special lines in the dark red region of the spectrum. These lines have not been found in the spectra of any known substance. Soon, similar dark red lines were discovered in the spectrum of sediment obtained after the evaporation of healing waters from the mineral springs of the Black Forest. It was natural to assume that these lines belonged to some new, previously unknown element. So in 1861 rubidium was discovered
The copper subgroup includes three elements - copper, silver and gold. Like alkali metal atoms, the atoms of all these elements have one electron in their outer layer; but their penultimate electronic layer contains, unlike alkali metal atoms, eighteen electrons. The structure of the two outer electron shells of the atoms of these elements can be depicted by the formula (where is the number of the period in which the element is located). All elements of the copper subgroup are the penultimate members of the decade elements. However, as can be seen from the above formula, their atoms contain not 9, but 10 electrons at the -sublevel. This is because structure is more stable than structure (see page 93). Comparing the data in Table. 31 with the corresponding values for alkali metals (Table 30), one can see that the radii of the atoms of copper, silver and gold are smaller than the radii of the atoms of the metals of the main subgroup. This causes a significantly higher density, high melting temperatures and large values of the enthalpy of atomization of the metals in question; Smaller atoms are arranged more densely in the lattice, as a result of which the attractive forces between them are strong. The small radius of the atoms also explains the higher ionization energies of the metals of this subgroup than the alkali metals. This leads to large differences in the chemical properties of the metals of both subgroups. Elements of the copper subgroup are low-active metals. They are difficult to oxidize, and, conversely, their ions are easily reduced; They do not decompose water; their hydroxides are relatively weak bases. In the voltage series they come after hydrogen. At the same time, the eighteen-electron layer, which is stable in other elements, is not yet completely stabilized here and is capable of partial loss of electrons. Thus, copper, along with singly charged cations, also forms doubly charged cations, which are even more characteristic of it. Similarly, for gold, the oxidation state is more characteristic than. The oxidation degree of silver in its usual compounds is equal; however, compounds with the oxidation degree of silver and are known.
45. Elements of the 3rd main subgroup
The third group of the periodic system covers a very large number of chemical elements, since its composition, in addition to elements of the main and secondary subgroups, includes elements with serial numbers 58-71 (lanthanides) and with serial numbers 90-103 (actinides). We will consider lanthanides and actinides along with elements of their secondary subgroup. The elements of the main subgroup of the third group - boron, aluminum, gallium, indium and thallium - are characterized by the presence of three electrons in the outer electron layer of the atom. The second outer electron layer of the boron atom contains two electrons, the aluminum atom - in. The metallic properties of the elements under consideration are less pronounced than those of the corresponding elements of the main subgroups of the second and especially the first group, and in boron non-metallic properties predominate. In compounds they exhibit an oxidation state of +3. However, as the atomic mass increases, lower oxidation states also appear. For the last element of the subgroup - thallium - the most stable compounds are those in which its oxidation state is +1. With an increase in the ordinal number, the metallic properties of the elements under consideration, as in other main subgroups, noticeably increase. Thus, boron oxide is acidic in nature, aluminum, gallium and indium oxides are amphoteric, and thallium (III) oxide is basic in nature. In practical terms, the most important of the elements of the third group are boron and aluminum.
46. Elements of the 4th main subgroup
The main subgroup of the fourth group of the periodic table is formed by five elements - carbon, silicon, germanium, tin and lead. When moving from carbon to lead, the sizes of atoms increase. Therefore, it should be expected that the ability to attach electrons, and therefore the non-metallic properties, will weaken, while the ease of electron release will increase. Indeed, germanium already exhibits metallic properties, while in tin and lead they predominate over non-metallic properties. Thus, only the first two members of the described group are non-metals; germanium is classified as both metals and non-metals, tin and lead are metals. The elements of the group under consideration are characterized by oxidation degrees of +2 and +4. Compounds of carbon and silicon, in which the oxidation state of these elements is +2, are few in number and relatively unstable. Table 28. Some properties of carbon and its analogues
47. Elements of the 5th main subgroup
The main subgroup of group V of the periodic table includes nitrogen, phosphorus, arsenic, antimony and bismuth. These elements, having five electrons in the outer layer of the atom, are generally characterized as nonmetals. However, their ability to attach electrons is much less pronounced than that of the corresponding elements of groups VI and VII. Due to the presence of five outer electrons, the highest positive oxidation of the elements of this subgroup is +5, and negative -3. Due to the relatively lower electronegativity, the bond of the elements under consideration with hydrogen is less polar than the bond with hydrogen of elements of groups VI and VII. Therefore, hydrogen compounds of these elements do not abstract hydrogen ions in an aqueous solution and, thus, do not have acidic properties. The physical and chemical properties of the elements of the nitrogen subgroup change with increasing atomic number in the same sequence that was observed in the previously considered groups. But since the non-metallic properties of nitrogen are less pronounced than those of oxygen and, especially, fluorine, the weakening of these properties when moving to the next elements entails the appearance and increase of metallic properties. The latter are already noticeable in arsenic, antimony has both properties approximately equally, and in bismuth the metallic properties predominate over the nonmetallic ones. The most important properties of the elements of the subgroup under consideration are given in Table. 27. Table 27. Some properties of nitrogen and its analogues
48. Organic carbon compounds
Carbon compounds (with the exception of some of the simplest ones) have long been called organic compounds, since in nature they are found almost exclusively in animal and plant organisms, take part in life processes, or are products of vital activity or decay of organisms. Unlike organic compounds, substances such as sand, clay, various minerals, water, carbon oxides, carbonic acid, its salts and others found in “inanimate nature” are called inorganic or mineral substances. The division of substances into organic and inorganic arose due to the uniqueness of organic compounds that have specific properties. For a long time it was believed that carbon-containing substances formed in organisms, in principle, cannot be obtained by synthesis from inorganic compounds. The formation of organic substances was attributed to the influence of a special, inaccessible to knowledge, “vital force”, acting only in living organisms and determining the specificity of organic substances. This doctrine, which was a variety of idealistic ideas about nature, received the name vitalism (from the Latin vis vitalis - life force). Vitalists tried to find in the phenomena of living nature evidence of the existence in the world of some mysterious forces that cannot be studied and do not obey general physical and chemical laws. The vitalist concept was most fully formulated by one of the most authoritative chemists of the first half of the 19th century - the Swedish scientist I. Ya. Berzelius. Vitalistic views hindered progress in the study of the nature of organic substances and were refuted in the course of the development of science. In 1824, the German chemist F. Wöhler, a student of Berzelius, for the first time obtained oxalic acid HOOC-COOH from the inorganic substance cyanogen NC-CN by heating it with water - an organic compound that until then had been extracted only from plants. In 1828, Wöhler carried out the first synthesis of a substance of animal origin: by heating the inorganic compound ammonium cyanate NH4CNO he obtained urea (urea) (NH2)CO; until then this substance had been isolated only from urine. Soon, syntheses of other organic substances were carried out in laboratory conditions: in 1845 in Germany, G. Kolbe synthesized acetic acid, in 1854 in France, M. Berthelot obtained fat synthetically, in 1861 in Russia, A. M. Butlerov carried out the synthesis of a sugary substance, etc. Currently, many organic compounds are obtained through synthesis. Moreover, it turned out that many organic substances are much easier and cheaper to obtain synthetically than to isolate them from natural products. The greatest success of chemistry in the 20th century was the first synthesis of simple proteins - the hormone insulin and the enzyme ribonuclease. Thus, the possibility of synthetic production of even proteins, the most complex organic substances that are indispensable participants in life processes, has been proven; according to the definition of F. Engels: “Life is a way of existence of protein bodies.” With the development of the synthesis of organic compounds, the line separating these compounds from inorganic ones was destroyed, but the name “organic compounds” was preserved. Most currently known carbon compounds are not even found in organisms, but are obtained artificially.
49. Elements of the 8th side group
A side subgroup of the eighth group of the periodic table covers three triads of d-elements. The first triad is formed by the elements Fe, cobalt and nickel, the second triad by ruthenium, rhodium and palladium, and the third triad by osmium, iridium and platinum. Most elements of the subgroup under consideration have two electrons in the outer electron layer of the atom; they are all metals. In addition to the outer electrons, electrons from the previous unfinished layer also take part in the formation of chemical bonds. These elements are characterized by oxidation degrees of 2, 3, 4. Higher oxidation degrees occur less frequently. A comparison of the physical and chemical properties of the elements of the eighth group shows that iron, cobalt and nickel, which are in the first large period, are very similar to each other and at the same time very different from the elements of the other two triads. Therefore, they are usually classified into the iron family. The remaining six elements of the eighth group are combined under the general name of platinum metals.
The main oxidation states of iron are +2 and +3.
When stored in air at temperatures up to 200 °C, iron is gradually covered with a dense film of oxide, which prevents further oxidation of the metal. In humid air, iron becomes covered with a loose layer of rust, which does not prevent the access of oxygen and moisture to the metal and its destruction. Rust does not have a constant chemical composition; approximately its chemical formula can be written as Fe2O3 xH2O.
Iron reacts with oxygen when heated. When iron burns in air, the oxide Fe3O4 is formed, when burned in pure oxygen, the oxide Fe2O3 is formed. If oxygen or air is passed through molten iron, FeO oxide is formed. When sulfur and iron powder are heated, sulfide is formed, the approximate formula of which can be written as FeS.
When heated, iron reacts with halogens. Since FeF3 is non-volatile, iron is resistant to fluorine up to temperatures of 200-300 °C. When iron is chlorinated (at a temperature of about 200 °C), a volatile dimer, Fe3Cl6, is formed. If the interaction of iron and bromine occurs at room temperature or with heating and increased bromine vapor pressure, FeBr3 is formed. When heated, FeCl3 and, especially, FeBr3 split off the halogen and transform into iron(II) halides. When iron and iodine react, iodide Fe3I8 is formed.
When heated, iron reacts with nitrogen, forming iron nitride Fe3N, with phosphorus, forming phosphides FeP, Fe2P and Fe3P, with carbon, forming carbide Fe3C, with silicon, forming several silicides, for example, FeSi.
At elevated pressure, metallic iron reacts with carbon monoxide (II) CO, and liquid, under normal conditions, highly volatile iron pentacarbonyl Fe(CO)5 is formed. Iron carbonyls of the compositions Fe2(CO)9 and Fe3(CO)12 are also known. Iron carbonyls serve as starting materials in the synthesis of organoiron compounds, including ferrocene of the composition (η5-C5H5)2Fe.
Pure metallic iron is stable in water and dilute alkali solutions. Iron does not dissolve in cold concentrated sulfuric and nitric acids due to passivation of the metal surface by a strong oxide film. Hot concentrated sulfuric acid, being a stronger oxidizing agent, interacts with iron.
Iron reacts with hydrochloric and dilute (approximately 20%) sulfuric acids to form iron(II) salts:
Fe + 2HCl → FeCl2 + H2;
Fe + H2SO4 → FeSO4 + H2.
When iron reacts with approximately 70% sulfuric acid, the reaction proceeds to form iron(III) sulfate:
2Fe + 6H2SO4 → Fe2(SO4)3 + 3SO2 + 6H2O.
Iron(II) oxide FeO has basic properties; the base Fe(OH)2 corresponds to it. Iron(III) oxide Fe2O3 is weakly amphoteric; it is matched by an even weaker base than Fe(OH)2, Fe(OH)3, which reacts with acids:
2Fe(OH)3 + 3H2SO4 → Fe2(SO4)3 + 6H2O.
Iron(III) hydroxide Fe(OH)3 exhibits weakly amphoteric properties; it is capable of reacting only with concentrated solutions of alkalis:
Fe(OH)3 + 3KOH → K3.
The resulting hydroxo complexes of iron(III) are stable in strongly alkaline solutions. When solutions are diluted with water, they are destroyed, and Fe(OH)3 precipitates.
Iron(III) compounds in solutions are reduced by metallic iron:
Fe + 2FeCl3 → 3FeCl2.
When storing aqueous solutions of iron(II) salts, oxidation of iron(II) to iron(III) is observed:
4FeCl2 + O2 + 2H2O → 4Fe(OH)Cl2.
Of the iron(II) salts, the most stable in aqueous solutions is Mohr's salt - double ammonium and iron(II) sulfate (NH4)2Fe(SO4)2 6H2O.
Iron(III) is capable of forming double sulfates with singly charged cations such as alum, for example, KFe(SO4)2 - iron-potassium alum, (NH4)Fe(SO4)2 - iron-ammonium alum, etc.
When chlorine gas or ozone acts on alkaline solutions of iron(III) compounds, iron(VI) compounds are formed - ferrates, for example, potassium ferrate(VI) K2FeO4. There are reports of the production of iron(VIII) compounds under the influence of strong oxidizing agents.
To detect iron(III) compounds in solution, a qualitative reaction of Fe3+ ions with SCN− thiocyanate ions is used. When Fe3+ ions interact with SCN− anions, bright red iron thiocyanate Fe(SCN)3 is formed. Another reagent for Fe3+ ions is potassium hexacyanoferrate(II) K4 (yellow blood salt). When Fe3+ and 4− ions interact, a bright blue precipitate of Prussian blue precipitates:
4K4 + 4Fe3+ → 4KFeIII↓ + 12K+.
Potassium hexacyanoferrate(III) K3 (red blood salt) can serve as a reagent for Fe2+ ions in solution. When Fe2+ and 3− ions interact, a Turnboole blue precipitate forms:
3K3 + 3Fe2+→ 3KFe2↓ + 6K+.
It is interesting that Prussian blue and Turnboule blue are two forms of the same substance, since an equilibrium is established in solution:
KFe3 ↔ KFe2.
Nickel is an element of the side subgroup of the eighth group, the fourth period of the periodic system of chemical elements of D.I. Mendeleev, with atomic number 28. It is designated by the symbol Ni (lat. Niccolum). The simple substance nickel is a ductile, malleable transition metal of a silvery-white color; at ordinary temperatures in air it is covered with a thin protective film of oxide. Chemically inactive.
Nickel atoms have an external electron configuration of 3d84s2. The most stable oxidation state for nickel is Ni(II).
Nickel forms compounds with oxidation states +2 and +3. In this case, nickel with an oxidation state of +3 is only available in the form of complex salts. A large number of ordinary and complex compounds are known for nickel +2 compounds. Nickel oxide Ni2O3 is a strong oxidizing agent.
Nickel is characterized by high corrosion resistance - it is stable in air, water, alkalis, and a number of acids. Chemical resistance is due to its tendency to passivation - the formation of a dense oxide film on its surface, which has a protective effect. Nickel actively dissolves in nitric acid.
With carbon monoxide CO, nickel easily forms the volatile and highly toxic carbonyl Ni(CO)4.
Fine nickel powder is pyrophoric (self-ignites in air).
Nickel burns only in powder form. Forms two oxides NiO and Ni2O3 and, accordingly, two hydroxides Ni(OH)2 and Ni(OH)3. The most important soluble nickel salts are acetate, chloride, nitrate and sulfate. Solutions are usually colored green, and anhydrous salts are yellow or brownish-yellow. Insoluble salts include oxalate and phosphate (green), three sulfides NiS (black), Ni2S3 (yellowish-bronze) and Ni3S4 (black). Nickel also forms numerous coordination and complex compounds. For example, nickel dimethylglyoximate Ni(C4H6N2O2)2, which gives a clear red color in an acidic environment, is widely used in qualitative analysis for the detection of nickel
An aqueous solution of nickel sulfate is green in color.
Aqueous solutions of nickel(II) salts contain hexaaquanickel(II) 2+ ion. When an ammonia solution is added to a solution containing these ions, nickel(II) hydroxide, a green, gelatinous substance, precipitates. This precipitate dissolves when excess ammonia is added due to the formation of hexamminnickel(II) 2+ ions.
Nickel forms complexes with tetrahedral and planar square structures. For example, the tetrachloronicickelate(II)2− complex has a tetrahedral structure, while the tetracyanonickelate(II)2− complex has a planar square structure.
Qualitative and quantitative analysis uses an alkaline solution of butanedione dioxime, also known as dimethylglyoxime, to detect nickel(II) ions. When it reacts with nickel(II) ions, the red coordination compound bis(butanedionedioximato)nickel(II) is formed. It is a chelate compound and the butanedione dioximate ligand is bidentate.
The mass fraction of cobalt in the earth's crust is 4×10−3%. Cobalt is part of the minerals: carolite CuCo2S4, linneite Co3S4, cobaltine CoAsS, spherocobaltite CoCO3, smaltite CoAs2, skutterudite (Co, Ni)As3 and others. In total, about 30 cobalt-containing minerals are known. Cobalt is accompanied by iron, nickel, manganese and copper. The content in sea water is approximately (1.7)×10−10%. In air, cobalt oxidizes at temperatures above 300 °C.
Cobalt oxide, stable at room temperature, is a complex oxide Co3O4, having a spinel structure, in the crystal structure of which one part of the nodes is occupied by Co2+ ions, and the other by Co3+ ions; decomposes to form CoO above 900 °C.
At high temperatures, the α-form or β-form of CoO oxide can be obtained.
All cobalt oxides are reduced with hydrogen. Co3O4 + 4H2 → 3Co + 4H2O.
Cobalt (III) oxide can be obtained by calcining cobalt (II) compounds, for example: 2Co(OH)2 + O2 → Co2O3 + H2O.
Platinum (lat. Platinum), Pt, chemical element of group VIII of the periodic system of Mendeleev, atomic number 78, atomic mass 195.09; heavy refractory metal.
Platinum's chemical properties are similar to palladium, but it exhibits greater chemical stability. Reacts only with hot aqua regia: 3Pt + 4HNO3 + 18HCl = 3H2 + 4NO + 8H2O
Platinum dissolves slowly in hot sulfuric acid and liquid bromine. It does not interact with other mineral and organic acids. When heated, it reacts with alkalis and sodium peroxide, halogens (especially in the presence of alkali metal halides): Pt + 2Cl2 + 2NaCl = Na2. When heated, platinum reacts with sulfur, selenium, tellurium, carbon and silicon. Like palladium, platinum can dissolve molecular hydrogen, but the volume of absorbed hydrogen is smaller and the ability of platinum to release it when heated is less.
When heated, platinum reacts with oxygen to form volatile oxides. The following platinum oxides are identified: black PtO, brown PtO2, reddish-brown PtO3, as well as Pt2O3 and Pt3O4.
The hydroxides Pt(OH)2 and Pt(OH)4 are known for platinum. They are obtained by alkaline hydrolysis of the corresponding chloroplatinates, for example: Na2PtCl4 + 2NaOH = 4NaCl + Pt(OH)2, Na2PtCl6 + 4NaOH = 6NaCl + Pt(OH)4. These hydroxides exhibit amphoteric properties: Pt(OH)2 + 2NaOH = Na2, Pt(OH)2 + 4HCl = H2 + 2H2O, Pt(OH)4 + 6HCl = H2 + 4H2O, Pt(OH)4 + 2NaOH = Na2.
Platinum hexafluoride PtF6 is one of the strongest oxidizing agents among all known chemical compounds, capable of oxidizing oxygen, xenon or NO molecules: O2 + PtF6 = O2+−. With its help, in particular, Canadian chemist Neil Bartlett in 1962 obtained the first true chemical compound of xenon, XePtF6.
With the interaction between Xe and PtF6 discovered by N. Bartlett, leading to the formation of XePtF6, the chemistry of inert gases began. PtF6 is obtained by fluorinating platinum at 1000 °C under pressure. Fluoridation of platinum at normal pressure and temperature of 350-400 °C gives Pt(IV) fluoride: Pt + 2F2 = PtF4 Platinum fluorides are hygroscopic and decompose with water. Platinum(IV) tetrachloride with water forms hydrates PtCl4 nH2O, where n = 1, 4, 5 and 7. By dissolving PtCl4 in hydrochloric acid, chloroplatinic acids H and H2 are obtained. Platinum halides such as PtBr4, PtCl2, PtCl2 2PtCl3, PtBr2 and PtI2 have been synthesized. Platinum is characterized by the formation of complex compounds of composition 2- and 2-. While studying platinum complexes, A. Werner formulated the theory of complex compounds and explained the nature of the occurrence of isomers in complex compounds.
Platinum is one of the most inert metals. It is insoluble in acids and alkalis, with the exception of aqua regia. Platinum also reacts directly with bromine, dissolving in it.
When heated, platinum becomes more reactive. It reacts with peroxides, and upon contact with atmospheric oxygen, with alkalis. A thin platinum wire burns in fluorine, releasing a large amount of heat. Reactions with other non-metals (chlorine, sulfur, phosphorus) occur less readily. When heated more strongly, platinum reacts with carbon and silicon, forming solid solutions, similar to the iron group metals.
In its compounds, platinum exhibits almost all oxidation states from 0 to +6, of which +2 and +4 are the most stable. Platinum is characterized by the formation of numerous complex compounds, of which many hundreds are known.
The main subgroup of Group I of the Periodic Table consists of lithium Li, sodium Na, potassium K, rubidium Rb, cesium Cs and francium Fr.
The atoms of these elements have one s-electron at the outer energy level: ns1. When entering into chemical interactions, atoms easily give up an electron from the outer energy level, exhibiting a constant oxidation state of +1 in compounds.
Elements of this subgroup belong to metals. Their common name is alkali metals.
In nature, sodium and potassium are the most common. The mass fraction of sodium in the earth's crust is 2.64%, potassium - 2.60%. Alkali metals do not occur in nature in a free state. The main natural Na compounds are the minerals halite, or rock salt, NaCl, and mirabilite, or Glauber's salt (Na2SO4 10H2O). The most important potassium compounds include sylvin (KCl), carnallite (KCl MgCl2 6H2O), sylvinite
Francium is a radioactive element. Traces of this element have been found in the decay products of natural uranium. Due to the short lifetime of Fr isotopes, it is difficult to obtain in large quantities, therefore the properties of metallic France and its compounds have not yet been sufficiently studied.
Properties: Alkali metals are silvery-white substances with low density. Lithium is the lightest of them. These are soft metals; the softness of Na, K, Rb, Cs is similar to wax. Alkali metals are fusible. The melting point of cesium is 28.5°C, the highest melting point of lithium (180.5°C). They have good electrical conductivity.
Alkali metals have high chemical activity, their activity increases in the series Li-Na-K-Rb-Cs-Fr. They are strong reducing agents in reactions.
1. Interaction with simple substances.
Alkali metals react with oxygen. All of them are easily oxidized by atmospheric oxygen, and rubidium and cesium even spontaneously ignite.
4Li + O2® 2Li2O(lithium oxide)
2Na + O2® Na2O2 (sodium peroxide)
K+O2® KO2 (potassium superoxide)
Alkali metals spontaneously ignite in fluorine, chlorine, bromine vapor, forming halides:
2Na+Br2®2NaBr (halide)
When heated, they interact with many non-metals:
2Na + S ® Na2S (sulfides)
6Li + N2® 2Li3N (nitrides)
2Li + 2C ® 2Li2C2 (carbides)
2. Interaction with water. All alkali metals react with water, reducing it to hydrogen. The activity of interaction of metals with water increases from lithium to cesium.
2Na + 2H2O ® 2NaOH + H2
2Li + 2H2O ® 2LiOH + H2
3. Interact with acids. Alkali metals react with hydrochloric and dilute sulfuric acids to release hydrogen:
2Na + 2HCl ® 2NaCl +H2
Concentrated sulfuric acid is reduced mainly to hydrogen sulfide:
8Na + 5H2SO4® 4Na2SO4+ H2S + 4H2O
In this case, a parallel reaction of the reduction of sulfuric acid to sulfur oxide (IV) and elemental sulfur is possible.
When an alkali metal reacts with dilute nitric acid, ammonia or ammonium nitrate is predominantly produced, and with concentrated acid, nitrogen or nitric oxide (I) is produced:
8Na +10HNO3(dil.)® 8NaNO3+ NH4NO3+ 3 H2O
8K +10HNO3(conc.)® 8KNO3+ NO2+ 5H2O
However, as a rule, several products are formed simultaneously.
4. Interaction with metal oxides and salts. Due to their high chemical activity, alkali metals can reduce many metals from their oxides and salts:
BeO +2Na ®Be + Na2O
CaCl2+ 2Na® Ca + 2NaCl
Receipt:
Metallic sodium is produced industrially by electrolysis of molten sodium chloride with inert electrodes. In the melt, sodium chloride dissociates into ions:
NaCl↔ Na+ + Cl-
During electrolysis, the Na+ cation is reduced at the cathode, and the Cl- anion is oxidized at the anode:
cathode: 2 Na++2е ® 2Na
anode: 2 Cl--2e ® Cl2
2Na++ 2Cl-® 2Na + Cl2 or 2NaCl®2Na + Cl
Thus, sodium and chlorine are formed during electrolysis. Sometimes sodium is obtained by electrolysis of molten sodium hydroxide.
Another way to obtain sodium is to reduce soda with coal at high temperatures:
Na2CO3+ 2C®2Na + 3CO
Potassium is obtained by replacing it with sodium from a melt of potassium chloride or potassium hydroxide:
KCl + Na ® K + NaCl
Potassium can also be obtained by electrolysis of melts of its compounds (KCl; KOH).
Lithium metal is produced by electrolysis of molten lithium chloride or reduction of lithium oxide with aluminum.
Rubidium and cesium are obtained by reducing their halides with metals in a vacuum:
2RbCl + Ca = 2Rb + CaCl2 ; 2CsCl + Mg = 2Cs + CaCl2
Alkali metal oxides (R2O):
Lithium and sodium oxides are white substances, potassium oxide is light yellow, rubidium is yellow, and cesium is orange. All oxides are reactive compounds and have pronounced basic properties, and in the series from lithium oxide to cesium oxide, the basic properties increase.
Oxidation of the metal produces only lithium oxide:
4Li + O2® 2Li2O
The remaining oxides are obtained indirectly. Thus, sodium oxide is obtained by reducing the sodium compound with sodium metal:
Na2O2+ 2Na ® 2Na2O
2NaOH + 2Na ® 2Na2O + H2
Alkali metal oxides easily react with water to form hydroxides, for example:
Li2O + H2O ® 2LiOH
They react with acidic oxides and acids to form salts:
Na2O + SO3 ® Na2SO4
K2O + 2HNO3 ® 2KNO3+ H2O
Alkali metal hydroxides (ROH):
They are white crystalline substances. All alkali metal hydroxides are strong bases, soluble in water. The common name is alkali.
Hydroxides are formed by the interaction of alkali metals or their oxides with water:
2Li + 2H2O ® 2LiOH + H2
Li2O + H2O ® 2LiOH
Sodium and potassium hydroxides, which are of great practical importance, are produced industrially by electrolysis of chlorides:
2NaCl + 2H2O ® 2NaOH + H2 + Cl2
cathode: 2H++ 2ē ® H02
anode: 2Cl-– 2ē ® Cl02
Alkali metal hydroxides exhibit all the characteristic properties of bases: they interact with acids and amphoteric oxides, amphoteric hydroxides, acids, and salts. Some metals that form amphoteric hydroxides dissolve in aqueous solutions of alkalis, for example:
Zn + 2NaOH + 2H2O = Na2 + H2
C Si Ge Sn Pb
general characteristics
Electronic configuration ns 2 np 2
Characteristic oxidation states: -4; 0; +2; +4.
The maximum valency of these elements, both in terms of donation and gain of electrons, is four. Due to the increase in the volume of atoms in the transition from carbon to lead, the process of electron acceptance weakens, and the ease of their loss increases, so the metallic properties of atoms increase from top to bottom.
Carbon
Occurrence in nature, application, physical properties. The forms of carbon in nature are diverse. In addition to the tissues of living organisms and the products of their destruction (coal, oil, etc.), it is part of many minerals, having mostly the general formula MSO 3, where M is a divalent metal. The most common of these minerals is calcite (CaCO 3), which sometimes forms huge accumulations in certain areas of the earth's surface. Carbon is found in the atmosphere as carbon dioxide, which is also found in a dissolved state in all natural waters.
In the form of charcoal, carbon has been known to mankind since time immemorial. It received its modern name in 1787. Natural carbon is composed of two isotopes - 12 C (98.892%) and 13 C (1.108%). The mass of the carbon-12 isotope is taken as a unit of atomic and molecular masses. In various natural objects, the ratio of both isotopes may vary slightly.
Free carbon occurs in nature in the form of two simple substances - diamond and graphite (the most stable form of carbon under normal conditions). These include the so-called “amorphous” carbon, the simplest representative of which is charcoal. Diamond has a density of 3.5 g/cm 3 and is the hardest of all minerals. The purest diamonds are colorless and transparent. Graphite is a gray mass that has a metallic sheen and is greasy to the touch. with a density of 2.2 g/cm 3. It is very soft - it is easily scratched with a fingernail and, when rubbed, leaves gray stripes on the paper. “Amorphous” carbon is quite close in properties to graphite.
Formation of natural diamonds occurred through the crystallization of carbon in the deep layers of the Earth (200-300 km from the surface) at temperatures of the order of 3000 ° C and pressures of the order of 200 thousand atm. Their primary deposits are associated with a very rare outcrop of a special rock - kimberlite - to the surface, and alluvial deposits are occasionally found in alluvial layers. Industrial mines contain on average only 0.5 g of diamond per ton of rock. Rich deposits were discovered in Yakutia (1955). The structure of diamond can be represented in the form of tetrahedra with a carbon atom in the center, which are repeated at infinity in three dimensions (Fig. 1). Diamond has an atomic crystal lattice.
Rice. 1. Diagram of the arrangement of C atoms in diamond. Rice. 2. Regular diamond cut.
Despite its hardness, diamond is fragile and easily breaks when struck. It conducts heat well, but practically does not conduct electric current. Not all diamonds are colorless; some have color ranging from subtle to intense. In relation to X-rays, a diamond is transparent (unlike fakes), but to ultraviolet rays some crystals are transparent, others are not.
Diamond is highly inert: it is not affected by either acids or alkalis. In air, diamond burns at a temperature of about 900 °C, and in oxygen - about 700 °C. After combustion, some ash remains (0.02 wt.% or more), which indicates the presence of impurities in natural diamonds (mainly aluminum, silicon, calcium and magnesium). When heated above 1200 °C in the absence of air, diamond graphitization begins.
The most beautiful diamonds are polished and, called diamonds (Fig. 2), used as jewelry. For their pricing, the unit of mass used for precious stones is the carat (0.2 g). The largest diamond mined (“Cullinan”) weighed 3026 carats, i.e. more than 600 g
The exceptional hardness of diamond determines its value for technology. The industry uses all those stones (the vast majority) that have some kind of flaw (ugly coloring, cracks, etc.) that makes them unsuitable as jewelry.
There is an assumption that the starting material for the natural synthesis of diamonds was carbon, which resulted from the reduction (at high temperatures and under high pressure) of carbonate rocks with divalent iron according to the approximate summary scheme:
CaCO 3 + 5 FeO → Ca(FeO 2) 2 + Fe 3 O 4 + C.
Fe +2 Fe +3 (Fe +3 O 4)
The very high pressure necessary for the crystallization of carbon in the form of diamond was created due to its random local increases.
Attempts to artificially produce diamonds were made many times, but for the first time they were successful only in 1953. Converting graphite to diamond can only be carried out at very high pressures, at high temperatures and in the presence of catalysts, of which some elements of the triads turned out to be the most suitable. Seed diamond crystals arise at the interface between graphite and the molten catalyst metal. They remain coated with a film of liquid carbon-containing metal, through which the carbon then diffuses from the graphite to the diamond as it grows. Modern technology makes it possible to obtain 20 g of diamonds in one chamber in a few minutes.
Another method of synthesis is also interesting - by acting on graphite (mixed with a catalyst) shock wave created by explosion. The instantaneous nature of this action is compensated by the occurrence of extremely high pressure and temperature at the moment of explosion. So in one of the experiments with a shock wave under a pressure of 300 thousand atm. almost all of the graphite taken turned into very small diamond crystals (up to 40 microns in size).
Artificial diamonds are small crystals, the predominant shape of which usually changes from cubic (at relatively low synthesis temperatures) to octahedral (at high temperatures). Their color is also different: from black at low temperatures to green, yellow and white at high temperatures. For example, in one of the experiments under a pressure of 200 thousand atm. By instantaneously (within thousandths of a second) heating of graphite with an electric discharge to 5000°C, colorless pure water diamonds were obtained. The color of artificial diamonds significantly depends on the nature of the impurities included in the crystals (and thereby on the composition of the original graphite mixture). For example, impurity Nickel gives greenish tones, and at the same time nickel and boron give blue tones.
A major consumer of graphite is the ceramics industry, producing crucibles for melting metals from a mixture of graphite and clay (graphite crucibles). Made from pressed graphite rocket gas rudders. In metallurgy, it is used to coat molds during casting. Due to the good electrical conductivity of graphite, it is used to make electrodes for electrical and electrometallurgical processes. A significant amount of graphite is used for manufacturing mineral paints and (mixed with clay) pencils. An interesting application of graphite is the use of its powder (alone or together with machine oil) as a lubricant for rubbing parts of mechanisms.
Less well known are two other allotropes of carbon - carbyne and fullerene .
There may be a linear form of elemental carbon different from both graphite and diamond carbine .
In addition, fullerenes C 70, C 74, C 84, etc., having the shape of a spheroid, were obtained (Fig. 2).
Rice. 2. Molecules C 60 and C 70.
Chemical properties. Carbon in the free state is typical reducing agent. When oxidized by oxygen in excess air, it turns into carbon monoxide (IV):
if there is a deficiency - into carbon monoxide (II):
Both reactions are highly exothermic.
When carbon is heated in an atmosphere, carbon monoxide (IV) is formed carbon monoxide:
Carbon reduces many metals from their oxides:
This is how reactions occur with oxides of cadmium, copper, and lead. When carbon interacts with oxides of alkaline earth metals, aluminum and some other metals, carbides:
This is explained by the fact that active metals are stronger reducing agents than carbon, therefore, when heated, the metals formed oxidize excess carbon, giving carbides:
Hydrogen compounds. U carbon and group IV elements form hydrides with the general formula E n H 2 n +2. For carbon, n can take large values; silicon – n = 1÷6; germanium – n = 1÷3; tin and lead – n = 1.
CH 4 – hydrogen carbide (methane). The gas is colorless and odorless, chemically inert, does not interact with acids and alkalis, easily ignites, and when mixed with air it is an explosive “explosive mixture”.
Methane derivatives - methanides: beryllium and aluminum carbide Be 2 C and Al 4 C 3. Refractory substances that decompose with water:
Al 4 C 3 + H 2 O → Al(OH) 3 + CH 4
Carbon forms a large number of percarbides:
C 2 H 6 ethane; C 2 H 4 ethene; C 2 H 2 ethyn.
Percarbides of metals s and d-elements, groups I and II (A) and aluminum are called acetylenides.
AgNO 3 + C 2 H 2 → Ag 2 C 2 + HNO 3
silver acetylide
Al + C 2 H 2 → Al 2 (C 2) 3 + H 2
aluminum acetylide
Calcium acetylide (percarbide) is prepared by heating calcium oxide with carbon:
CaO + C t → CaC 2 + CO
(percarbide), simply called calcium carbide, it decomposes with water:
CaC 2 + H 2 O → Ca(OH) 2 + C 2 H 2 - used to produce acetylene
1) typical oxidizing agents
2) transition elements
3) * – elements
1) gaseous substances, colorless and odorless
2) liquids at room temperature
3) metals
2) two electrons in the outer level
3) completely filled outer level
1) the ability to lose two outer electrons, forming a cation
with oxidation state +2
with oxidation state +1
3) the ability to acquire one electron to the outer level, forming
anion with oxidation state -1
1) very hard
2) the most widespread in the earth’s crust
3) radioactive
1) self-ignite in air
2) stored in water
3) stored in kerosene
all these metals are...
1) typical insulators
2) strong reducing agents
3) oxidizing agents
1) peroxides of the composition М2О2
2) oxides of the composition Me2O
t hv
1) K2O2 + 2K === 2K2O 3) 2KO2 + O3 === K2O + 3O2
t
2) KO2 + 3K === 2K2O
forming...
1) hydroxides
2) hydrates
3) hydrides
1) hydrides and oxygen
2) alkalis and hydrogen
3) peroxides and hydrogen
1) KSO4
2) KHSO4
3) K2SO4
1) which in the series of metal stresses come after H
2) which in the voltage series of metals are up to N
3) everyone will react
1) Al and Na
2) K and Na
3) K and Mn
1) silver and iron
2) ferric chloride and silver
3) copper (I I) nitrate and silver
a) Na + H2 →
b) NaO2 + H2O →
1. All elements of the main subgroup of group I of the periodic table belong to ...
3) * – elements
2. All s - elements, except hydrogen and helium, are...
3) metals
3. Atoms of alkaline elements have...
1) one electron per outer level
4. Atoms of alkaline elements have...
2) the ability to lose a single external electron, forming a cation
with oxidation state +1
5. Francium, completing group I, is...
3) radioactive
6. All s – metals are very active and therefore...
3) stored in kerosene
7. Since the outer electrons of s - metals easily pass to other elements, all these metals are ...
2) strong reducing agents
8. All alkali metals burn in an oxygen atmosphere, forming...
3) Me2O2 peroxides and Me2O oxides
9. Potassium oxide can be obtained as a result of the reaction...
T
2) KO2 + 3K === 2K2O
10. All s – metals combine with hydrogen even with slight heating,
forming...
3) hydrides
11. When alkali metals interact with water, they form...
2) alkalis and hydrogen
12. Potassium sulfate is a substance with the formula...
3) K2SO4
13. Metals of group I of the main subgroup will interact with acids...
3) everyone will react
14. A pair of metals with the most similar properties:
2) K and Na
15. A substitution reaction is possible between...
4) silver nitrate and iron
I I task: Carry out the transformations, write down the corresponding reaction equations.
Na → Na2O2 → NaO2 → Na2CO3 → Na Cl
Na2O2 + Na = Na2O
Na2O + CO2 = Na2CO3
Na2CO3 +2HCl = 2NaCl +H2O +CO2
I I I task: complete the reaction equations
a) 2Na + H2 → 2NaH
b) Na2O + H2O → 2NaOH
1. All elements of the main subgroup of group I of the periodic table belong to
3) * – elements
2. All s - elements, except hydrogen and helium, are...
3) metals
3. Atoms of alkaline elements have...
1) one electron per outer level
4. Atoms of alkaline elements have...
2) the ability to lose a single external electron, forming a cation
with oxidation state +1
5. Francium, completing group I, is...
3) radioactive
6. All s – metals are very active and therefore...
There is no correct answer here: lithium does not ignite in air and can not be stored in kerosene, but most likely leave it, for other metals it is true
3) stored in kerosene
7. Since the outer electrons of s – metals easily transfer to other elements,
all these metals are...
2) strong reducing agents
8. All alkali metals burn in an oxygen atmosphere, forming...
3) Me2O2 peroxides and Me2O oxides, although not for everyone either
9. Potassium oxide can be obtained as a result of the reaction
2) KO2 + 3K === 2K2O
10. All s-metals combine with hydrogen even with slight heating, forming...
3) hydrides
11. When alkali metals interact with water, they form...
2) alkalis and hydrogen
12. Potassium sulfate is a substance with the formula...
3) K2SO4
13. Metals of group I of the main subgroup will interact with acids...
3) everyone will react
14. A pair of metals with the most similar properties:
2) K and Na
15. A substitution reaction is possible between...
4) silver nitrate and iron
I I task: Carry out the transformations, write down the corresponding reaction equations.
Na → Na2O2 → NaO2 → Na2CO3 → Na Cl
2Nа+О2=Nа2О2
Na2O2+2Na=2Na2O
Na2O+CO2=Na2CO3
Na2СО3+2НCl=2NаСl+СО2+Н2О
I I I task: complete the reaction equations
a) 2Na + H2 →2NaH
b) 2NaO2 + H2O →NaOH + NaHO2 + O2