The concept of alkaline earth metals includes some of the elements of group II of the Mendeleev system: beryllium, magnesium, calcium, strontium, barium, radium. The last four metals have the most pronounced signs of an alkaline earth classification, therefore, in some sources, beryllium and magnesium are not included in the list, being limited to four elements.
The metal got its name due to the fact that when their oxides interact with water, an alkaline medium is formed. Physical properties of alkaline earth metals: all elements have a gray metallic color, under normal conditions they have a solid structure, with an increase in the serial number, their density increases, and have a very high melting point. Unlike alkali metals, elements of this group are not cut with a knife (with the exception of strontium). Chemical properties of alkaline earth metals: they have two valence electrons, the activity increases with increasing serial number, in reactions they act as a reducing agent.
The characteristics of alkaline earth metals indicate their high activity. This is especially true for elements with a large serial number. For example, beryllium under normal conditions does not interact with oxygen and halogens. To start the reaction mechanism, it must be heated to a temperature of over 600 degrees Celsius. Magnesium under normal conditions has an oxide film on its surface and also does not react with oxygen. Calcium oxidizes, but rather slowly. But strontium, barium and radium are oxidized almost instantly, so they are stored in an oxygen-free environment under a kerosene layer.
All oxides enhance the basic properties with an increase in the serial number of the metal. Beryllium hydroxide is an amphoteric compound that does not react with water, but dissolves well in acids. Magnesium hydroxide is a weak alkali, insoluble in water, but reactive with strong acids. Calcium hydroxide is a strong base that is poorly soluble in water and reacts with acids. Barium and strontium hydroxides are strong bases that are readily soluble in water. And radium hydroxide is one of the strongest alkalis that reacts well with water and almost all types of acids.
Methods of obtaining
Hydroxides of alkaline earth metals are obtained by the action of water on a pure element. The reaction takes place at room conditions (except for beryllium, which requires an increase in temperature) with the evolution of hydrogen. When heated, all alkaline earth metals react with halogens. The obtained compounds are used in the manufacture of a wide range of products from chemical fertilizers to ultra-precise microprocessor parts. Compounds of alkaline earth metals exhibit the same high activity as pure elements, therefore they are used in many chemical reactions.
Most often this occurs during exchange reactions, when it is necessary to displace a less active metal from a substance. They take part in redox reactions as a strong reducing agent. Divalent cations of calcium and magnesium impart so-called hardness to water. Overcoming this phenomenon occurs by the deposition of ions using physical impact or the addition of special emollients to the water. Salts of alkaline earth metals are formed by dissolving elements in acid or as a result of exchange reactions. The resulting compounds have a strong covalent bond, and therefore have low electrical conductivity.
In nature, alkaline earth metals cannot be in their pure form, since they quickly interact with the environment, forming chemical compounds. They are part of the minerals and rocks contained in the thickness of the earth's crust. Calcium is the most common, magnesium is slightly inferior to it, barium and strontium are quite common. Beryllium is a rare metal, while radium is a very rare one. For all the time that has passed since the discovery of radium, only one and a half kilograms of pure metal have been mined all over the world. Like most radioactive elements, radium has four isotopes.
Alkaline earth metals are obtained by decomposing complex substances and isolating a pure substance from them. Beryllium is obtained by reducing it from fluoride when exposed to high temperatures. Barium is reduced from its oxide. Calcium, magnesium and strontium are obtained by electrolysis of their chloride melt. Pure radium is the most difficult to synthesize. It is mined by impacting on uranium ore. Scientists estimate that on average there are 3 grams of pure radium per ton of ore, although there are also rich deposits, which contain as much as 25 grams per ton. Methods of precipitation, fractional crystallization and ion exchange are used to isolate the metal.
Application of alkaline earth metals
The range of applications for alkaline earth metals is very extensive and covers many industries. Beryllium is in most cases used as an alloying addition to various alloys. It increases the hardness and strength of materials, protects the surface well from corrosion. Also, due to the weak absorption of radioactive radiation, beryllium is used in the manufacture of X-ray machines and in nuclear power.
Magnesium is used as one of the reducing agents in titanium production. Its alloys are characterized by high strength and lightness, therefore they are used in the manufacture of aircraft, cars, and rockets. Magnesium oxide burns with a bright, blinding flame, which is reflected in the military, where it is used to make incendiary and tracer shells, flares and stun grenades. It is one of the most important elements for the regulation of the normal life process of the body, therefore it is part of some drugs.
Pure calcium is practically not used. It is needed to restore other metals from their compounds, as well as in the production of preparations for strengthening bone tissue. Strontium is used for the reduction of other metals and as the main component for the production of superconducting materials. Barium is added to many alloys that are designed to work in corrosive environments, as it has excellent protective properties. Radium is used in medicine for short-term skin irradiation in the treatment of malignant tumors.
These are the elements of group I of the periodic system: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr); very soft, plastic, low-melting and light, as a rule, silvery white; chemically very active; react violently with water, forming alkalis(where the name comes from).
All alkali metals are extremely active, in all chemical reactions they exhibit reducing properties, give up their only valence electron, turning into a positively charged cation, and exhibit a single oxidation state of +1.
The regenerative capacity increases in the order –– Li – Na – K – Rb – Cs.
All alkali metal compounds are ionic.
Almost all salts are water soluble.
Low melting points,
Low density values,
Soft, cut with a knife
Due to their activity, alkali metals are stored under a layer of kerosene to block the access of air and moisture. Lithium is very light and floats to the surface in kerosene, so it is stored under a layer of petroleum jelly.
Chemical properties of alkali metals
1. Alkali metals actively interact with water:
2Na + 2H 2 O → 2NaOH + H 2
2Li + 2H 2 O → 2LiOH + H 2
2. Reaction of alkali metals with oxygen:
4Li + O 2 → 2Li 2 O (lithium oxide)
2Na + O 2 → Na 2 O 2 (sodium peroxide)
K + O 2 → KO 2 (potassium superoxide)
In air, alkali metals are instantly oxidized. Therefore, they are stored under a layer of organic solvents (kerosene, etc.).
3. In the reactions of alkali metals with other non-metals, binary compounds are formed:
2Li + Cl 2 → 2LiCl (halides)
2Na + S → Na 2 S (sulfides)
2Na + H 2 → 2NaH (hydrides)
6Li + N 2 → 2Li 3 N (nitrides)
2Li + 2C → Li 2 C 2 (carbides)
4. Reaction of alkali metals with acids
(rarely carried out, there is a competing reaction with water):
2Na + 2HCl → 2NaCl + H 2
5. Interaction of alkali metals with ammonia
(sodium amide is formed):
2Li + 2NH 3 = 2LiNH 2 + H 2
6. Interaction of alkali metals with alcohols and phenols, which in this case exhibit acidic properties:
2Na + 2C 2 H 5 OH = 2C 2 H 5 ONa + H 2;
2K + 2C 6 H 5 OH = 2C 6 H 5 OK + H 2;
7. Qualitative reaction to alkali metal cations - coloring of the flame in the following colors:
Li + - carmine red 
Na + - yellow
K +, Rb + and Cs + - purple
Obtaining alkali metals
Metallic lithium, sodium and potassium get by electrolysis of molten salts (chlorides), and rubidium and cesium - by reduction in vacuum when their chlorides are heated with calcium: 2CsCl + Ca = 2Cs + CaCl 2
The vacuum-thermal production of sodium and potassium is also used on a small scale:
2NaCl + CaC 2 = 2Na + CaCl 2 + 2C;
4KCl + 4CaO + Si = 4K + 2CaCl 2 + Ca 2 SiO 4.
Active alkali metals are released in vacuum-thermal processes due to their high volatility (their vapors are removed from the reaction zone).
Features of the chemical properties of group I s-elements and their physiological action
The electronic configuration of the lithium atom is 1s 2 2s 1. It has the largest atomic radius in the 2nd period, which facilitates the detachment of a valence electron and the appearance of a Li + ion with a stable configuration of an inert gas (helium). Consequently, its compounds are formed with the transfer of an electron from lithium to another atom and the appearance of an ionic bond with a small fraction of covalence. Lithium is a typical metallic element. As a substance, it is an alkali metal. It differs from other members of group I by its small size and the smallest, in comparison with them, activity. In this respect, it resembles the Group II element, magnesium, located diagonally from Li. In solutions, the Li + ion is highly solvated; it is surrounded by several tens of water molecules. Lithium in terms of the energy of solvation - the addition of solvent molecules, is closer to the proton than to the cations of alkali metals.
The small size of the Li + ion, the high charge of the nucleus and only two electrons create conditions for the appearance of a rather significant field of a positive charge around this particle, therefore, in solutions, a significant number of molecules of polar solvents are attracted to it and its coordination number is large, the metal is able to form a significant number of organolithium compounds ...
The third period begins with sodium, therefore, at the external level, it has only 1e - , occupying a 3s orbital. The radius of the Na atom is the largest in the 3rd period. These two features determine the nature of the item. Its electronic configuration is 1s 2 2s 2 2p 6 3s 1 . The only oxidation state of sodium is +1. Its electronegativity is very small, therefore, sodium is present in compounds only in the form of a positively charged ion and gives the chemical bond an ionic character. The Na + ion is much larger in size than Li +, and its solvation is not so great. However, it does not exist in free solution in solution.
The physiological significance of K + and Na + ions is associated with their different adsorbability on the surface of the components that make up the earth's crust. Sodium compounds are only slightly subject to adsorption, while potassium compounds are firmly held by clay and other substances. Cell membranes, being the cell-medium interface, are permeable to K + ions, as a result of which the intracellular concentration of K + is much higher than that of Na + ions. At the same time, the concentration of Na + in blood plasma exceeds the content of potassium in it. This circumstance is associated with the emergence of the membrane potential of cells. Ions K + and Na + are one of the main components of the liquid phase of the body. Their ratio with Ca 2+ ions is strictly defined, and its violation leads to pathology. The introduction of Na + ions into the body does not have a noticeable harmful effect. An increase in the content of K + ions is harmful, but under normal conditions, an increase in its concentration never reaches dangerous values. The influence of Rb +, Cs +, Li + ions has not yet been sufficiently studied.
Of the various injuries associated with the use of alkali metal compounds, burns with hydroxide solutions are most common. The action of alkalis is associated with the dissolution of skin proteins in them and the formation of alkaline albuminates. Alkali is released again as a result of their hydrolysis and acts on the deeper layers of the body, causing ulcers. Under the influence of alkalis, nails become dull and brittle. Damage to the eyes, even with very dilute alkali solutions, is accompanied not only by superficial damage, but by disturbances in the deeper parts of the eye (iris) and leads to blindness. During the hydrolysis of alkali metal amides, alkali and ammonia are simultaneously formed, which cause tracheobronchitis of the fibrinous type and inflammation of the lungs.
Potassium was obtained by G. Davy almost simultaneously with sodium in 1807 by electrolysis of wet potassium hydroxide. From the name of this compound - "caustic potassium" and the element got its name. The properties of potassium differ markedly from those of sodium, which is due to the difference in the radii of their atoms and ions. In potassium compounds, the bond is more ionic, and in the form of the K + ion, it has a lower polarizing effect than sodium, due to its large size. The natural mixture consists of three isotopes 39 K, 40 K, 41 K. One of them is 40 K ‑ radioactive and a certain proportion of the radioactivity of minerals and soil is associated with the presence of this isotope. Its half-life is long - 1.32 billion years. It is quite easy to determine the presence of potassium in a sample: vapors of the metal and its compounds color the flame in a violet-red color. The spectrum of the element is quite simple and proves the presence of 1e - on the 4s-orbital. Its study served as one of the grounds for finding general patterns in the structure of the spectra.
In 1861, while studying the salt of mineral springs by spectral analysis, Robert Bunsen discovered a new element. Its presence was proved by dark red lines in the spectrum, which were not given by other elements. According to the color of these lines, the element was named rubidium (rubidus-dark red). In 1863 R. Bunsen obtained this metal in its pure form by reduction of rubidium tartrate (tartrate salt) with soot. A feature of the element is the easy excitability of its atoms. Electronic emission from it appears under the influence of red rays of the visible spectrum. This is due to the small difference in the energies of the atomic 4d and 5s orbitals. Of all alkaline elements with stable isotopes, rubidium (like cesium) has one of the largest atomic radii and a small ionization potential. These parameters determine the nature of the element: high electropositiveness, extreme chemical activity, low melting point (39 0 C) and low resistance to external influences.
The discovery of cesium, like rubidium, is associated with spectral analysis. In 1860, R. Bunsen discovered two bright blue lines in the spectrum that did not belong to any element known by that time. Hence the name "cesius" (caesius), which means sky blue. It is the last element in the alkali metal subgroup that is still found in measurable amounts. The largest atomic radius and the smallest first ionization potentials determine the character and behavior of this element. It has a pronounced electropositiveness and pronounced metallic qualities. The desire to donate the outer 6s electron leads to the fact that all its reactions are extremely violent. The small difference in the energies of the atomic 5d and 6s orbitals is responsible for the slight excitability of atoms. Electronic emission from cesium is observed under the influence of invisible infrared rays (heat). The specified feature of the atomic structure determines the good electrical conductivity of the current. All this makes cesium indispensable in electronic devices. V Lately more and more attention is paid to cesium plasma as the fuel of the future and in connection with the solution of the problem of thermonuclear fusion.
In air, lithium actively reacts not only with oxygen, but also with nitrogen and is covered with a film consisting of Li 3 N (up to 75%) and Li 2 O. The rest of the alkali metals form peroxides (Na 2 O 2) and superoxides (K 2 O 4 or KO 2).
The listed substances react with water:
Li 3 N + 3 H 2 O = 3 LiOH + NH 3;
Na 2 O 2 + 2 H 2 O = 2 NaOH + H 2 O 2;
K 2 O 4 + 2 H 2 O = 2 KOH + H 2 O 2 + O 2.
For the regeneration of air on submarines and spaceships, in insulating gas masks and breathing apparatus of combat swimmers (underwater saboteurs), a mixture of "oxon" was used:
Na 2 O 2 + CO 2 = Na 2 CO 3 + 0.5O 2;
K 2 O 4 + CO 2 = K 2 CO 3 + 1.5 O 2.
It is currently the standard filling of regenerative insulating gas mask cartridges for firefighters.
Alkali metals react with hydrogen when heated to form hydrides:
Lithium hydride is used as a strong reducing agent.
Hydroxides alkali metals corrode glass and porcelain dishes, they cannot be heated in quartz dishes:
SiO 2 + 2NaOH = Na 2 SiO 3 + H 2 O.
Sodium and potassium hydroxides do not split off water when heated up to their boiling points (more than 1300 0 С). Some sodium compounds are called sodas:
a) soda ash, anhydrous soda, linen soda or just soda - sodium carbonate Na 2 CO 3;
b) crystalline soda - sodium carbonate crystalline hydrate Na 2 CO 3. 10H 2 O;
c) bicarbonate or drinking - sodium bicarbonate NaHCO 3;
d) sodium hydroxide NaOH is called caustic soda or caustic.
The second group of the periodic table of D.I.Mendeleev contains a group of elements that are very similar in their properties to alkali metals, but inferior to them in activity. It includes beryllium and magnesium, as well as calcium, strontium, barium and radium. They are collectively known as alkaline earth elements. In our article, we will get acquainted with their distribution in nature and their application in industry, as well as study the most important chemical properties of alkaline earth metals.
general characteristics
All atoms of the above listed elements contain two electrons on the outer energy layer. Interacting with other substances, they always give up their negative particles, passing into the state of cations with a charge of 2+. In redox reactions, elements behave like strong reducing agents. As the nuclear charge increases, the chemical properties of alkaline earth metals and their activity increase. In air, they quickly oxidize, forming an oxide film on their surface. The general formula for all oxides is RO. They correspond to hydroxides with the formula R (OH) 2. Their basic properties and solubility in water also increase with an increase in the ordinal number of the element.
Special properties of beryllium and magnesium
In some of their properties, the first two representatives of the main subgroup of the second group are somewhat different from other alkaline earth elements. This manifests itself, in particular, during their interaction with water. For example, the chemical properties of beryllium are such that it does not react with H 2 O at all. Magnesium, on the other hand, interacts with water only when heated. But all alkaline earth elements easily react with it at ordinary temperatures. What substances are formed in this case?
Alkaline earth metal bases
Being active elements, calcium, barium and other members of the group quickly displace hydrogen from water, resulting in their hydroxides. The interaction of alkaline earth metals with water proceeds violently, with the release of heat. Solutions of bases of calcium, barium, strontium are soapy to the touch, and in contact with the skin and mucous membranes of the eyes cause severe burns. The first aid in such cases will be the treatment of the wound surface with a weak solution of acetic acid. It neutralizes alkali and reduces the risk of necrosis of damaged tissues.
Chemical properties of alkaline earth metals
Interaction with oxygen, water and non-metals is the main list of properties of metals included in the second group of the periodic table of chemical elements. For example, calcium, even under normal conditions, reacts with halogens: fluorine, chlorine, bromine and iodine. When heated, it combines with sulfur, carbon and nitrogen. Hard oxidation - combustion, ends with the formation of calcium oxide: 2Ca + O 2 = 2 CaO. The interaction of metals with hydrogen leads to the formation of hydrides. They are white refractory substances with ionic crystal lattices. The important chemical properties of alkaline earth metals include their interaction with water. As mentioned earlier, the product of this substitution reaction will be a metal hydroxide. Note also that in the main subgroup of the second group, calcium occupies the most significant place. Therefore, let us dwell on its characteristics in more detail.
Calcium and its compounds
The content of the element in the earth's crust is up to 3.5%, which indicates its wide distribution in minerals such as limestone, chalk, marble and calcite. Natural calcium contains six types of isotopes. It is also found in natural water sources. Compounds of alkali metals are studied in detail in the course of inorganic chemistry. For example, in 9th grade, students learn that calcium is a lightweight, but strong, silvery-white metal. Its melting and boiling point is higher than that of alkaline elements. The main production method is electrolysis of a mixture of molten salts of calcium chloride and fluoride. Its main chemical properties include its reactions with oxygen, water and non-metals. Of the compounds of alkali metals, calcium oxide and base are of the greatest importance for industry. The first compound is obtained from chalk or limestone by burning them out.
Further, calcium hydroxide is formed from calcium oxide and water. A mixture of it with sand and water is called mortar. It continues to be used as a plaster and for joining bricks when laying walls. A solution of calcium hydroxide, called lime water, is used as a reagent to detect carbon dioxide. When carbon dioxide is passed through a transparent aqueous solution of Ca (OH) 2, its turbidity is observed due to the formation of an insoluble precipitate of calcium carbonate.
Magnesium and its characteristics
The chemistry of alkaline earth metals studies the properties of magnesium, focusing on some of its features. It is a very light, silvery-white metal. Magnesium, melted in an atmosphere with high humidity, actively absorbs hydrogen molecules from water vapor. Cooling down, the metal almost completely releases them back into the air. It reacts very slowly with water due to the formation of a poorly soluble compound - magnesium hydroxide. Alkalis have no effect on magnesium at all. The metal does not react with some acids: concentrated sulfate and hydrofluoric, due to its passivation and the formation of a protective film on the surface. Most of the mineral acids dissolve the metal, which is accompanied by a violent evolution of hydrogen. Magnesium is a strong reducing agent, it replaces many metals from their oxides or salts:
BeO + Mg = MgO + Be.
The metal, together with beryllium, manganese, and aluminum, is used as an alloying addition to steel. Magnesium-containing alloys - electrons - have especially valuable properties. They are used in aircraft and automobile manufacturing, as well as in optical technology parts.
The role of elements in the life of organisms
Let us give examples of alkaline earth metals, the compounds of which are common in nature. Magnesium is the central atom in chlorophyll molecules in plants. It participates in the process of photosynthesis and is part of the active centers of the green pigment. Magnesium atoms fix light energy, then converting it into energy chemical bonds organic compounds: glucose, amino acids, glycerin and fatty acids. An important role is played by the element as a necessary component of enzymes that regulate metabolism in the human body. Calcium is a macronutrient that ensures the efficient passage of electrical impulses through the nervous tissue. The presence of its phosphate salts in the composition of bones and tooth enamel gives them hardness and strength.
Beryllium and its properties
Alkaline earth metals also include beryllium, barium and strontium. Consider beryllium. The element is not widespread in nature, it is mainly found in the composition of minerals, for example, beryl. Its varieties, containing multi-colored impurities, form gems: emeralds and aquamarines. A feature of the physical properties is fragility and high hardness. A distinctive feature of the atom of an element is the presence at the second outside energy level, not eight, as in all other alkaline earth metals, but only two electrons.
Therefore, the radius of the atom and ion is disproportionately small, the ionization energy is large. This determines the high strength of the crystal lattice of the metal. The chemical properties of beryllium also distinguish it from other elements of the second group. It reacts not only with acids, but also with alkali solutions, displacing hydrogen and forming hydroxyberyllates:
Be + 2NaOH + 2H 2 O = Na 2 + H 2.
Metal has a number of unique characteristics. Due to its ability to transmit X-rays, it is used to make X-ray tube windows. In the nuclear industry, the element is considered the best moderator and reflector of neutrons. In metallurgy, it is used as a valuable alloying additive that increases the anti-corrosion properties of alloys.
Strontium and barium
The elements are quite common in nature and, like the alkaline earth metal magnesium, are part of the minerals. Let's call them: barite, celestine, strontianite. Barium has the appearance of a silvery-white plastic metal. Like calcium, it is represented by several isotopes. In air, it actively interacts with its components - oxygen and nitrogen, forming barium oxide and nitride. For this reason, the metal is stored under a layer of paraffin or mineral oil, avoiding contact with air. Both metals form peroxides when heated to 500 ° C.
Of these, barium peroxide, used as a fabric bleach, has a practical application. The chemical properties of the alkaline earth metals, barium and strontium, are similar to those of calcium. However, their interaction with water is much more active, and the resulting bases are stronger than calcium hydroxide. Barium is used as an additive to liquid metal coolants, which reduces corrosion, in optics, and in the manufacture of vacuum electronic devices. Strontium is in demand in the production of solar cells and phosphors.
Qualitative reactions using alkaline earth metal ions
Barium and strontium compounds are examples of alkaline earth metals widely used in pyrotechnics due to their bright ion coloration of flames. So, strontium sulfate or carbonate gives a carmine-red glow of the flame, and the corresponding barium compounds - yellow-green. To detect calcium ions in the laboratory, several grains of calcium chloride are poured onto the burner flame, the flame turns brick-red.
A solution of barium chloride is used in analytical chemistry to detect ions of an acidic residue of sulfate acid in a solution. If, when draining the solutions, a white precipitate of barium sulfate forms, it means that there were particles of SO 4 2- in it.
In our article, we studied the properties of alkaline earth metals and gave examples of their use in various industries.
General characteristics of IIA group of the Periodic System of Elements.
The following elements are located in this group: Be, Mg, Ca, Sr, Ba, Ra. They have a common electronic configuration: (n-1) p6ns2, except for Be 1s22s2. Due to the latter, the properties of Be slightly differ from the properties of the subgroup as a whole. The properties of magnesium also differ from those of the subgroup, but to a lesser extent. In the Ca - Sr - Ba - Ra series, the properties change sequentially. The relative electronegativity in the Be - Ra series decreases because with an increase in atomic size, valence electrons are donated more readily. The properties of the elements of the IIA subgroup are determined by the ease of recoil of two ns electrons. In this case, E2 + ions are formed. In the study of X-ray diffraction, it turned out that in some compounds the elements of the IIA subgroup exhibit univalence. An example of such compounds is EG, which is obtained by adding E to the EG2 melt. All elements of this series are not found in nature in a free state due to their high activity.
Alkaline earth metals.
Calcium, strontium, barium and radium are called alkaline earth metals. They are named so because their oxides give water an alkaline environment.
History of alkaline earth metals
Limestone, marble and gypsum were already used by the Egyptians in the construction business in ancient times (5000 years ago). Until the end of the 18th century, chemists considered lime to be a simple substance. In 1746 I. Pott obtained and described a fairly pure calcium oxide. In 1789 Lavoisier suggested that lime, magnesia, barite are complex substances. Long before the discovery of strontium and barium, their "undeciphered" compounds were used in pyrotechnics to produce red and green lights, respectively. Until the mid-40s of the last century, strontium was primarily the metal of "amusing fires". In 1787, a new mineral was found in a lead mine near the Scottish village of Strontian, which was named strontianite SrCO3. A. Crawford suggested the existence of a still unknown "land". In 1792 T. Hop proved that the found mineral contains a new element - strontium. At the same time, insoluble strontium disaccharate (С12Н22О4.2SrO) was isolated with the help of Sr (OH) 2, when sugar was obtained from molasses. Sr production increased. However, it was soon noticed that a similar calcium saccharate was also insoluble, and calcium oxide was undoubtedly cheaper. Interest in strontium immediately disappeared and again increased to it only in the 40s of the last century. Heavy spar was the first known barium compound. It was discovered at the beginning of the 17th century. Italian alchemist Casciarolo. He also found that this mineral, after strong heating with coal, glows in the dark with a red light and gave it the name "lapis solaris" (solar stone). In 1808, Davy, subjecting a mixture of wet slaked lime with mercury oxide to electrolysis with a mercury cathode, prepared an amalgam of calcium, and after removing mercury from it, he obtained a metal called "calcium" (from the Latin Calx, genus calcis - lime). In the same way, Davy was obtained by Wa and Sr. An industrial method for producing calcium was developed by Suter and Redlich in 1896 at the Rathenau plant (Germany). In 1904, the first calcium production plant began to operate.
Radium was predicted by Mendeleev in 1871 and discovered in 1898 by the spouses Maria and Pierre Curie. They found that uranium ores are more radioactive than uranium itself. Radium compounds were the cause. They processed the remains of uranium ore with alkali, and what did not dissolve - with hydrochloric acid. The residue after the second procedure was more radioactive than the ore. Radium was found in this fraction. The Curies reported their discovery in a report for 1898.
Alkaline earth metals prevalence
The content of calcium in the lithosphere is 2.96% of the total mass of the earth's crust, strontium - 0.034%, barium - 0.065%, radium - 1.10-10%. In nature, calcium consists of isotopes with mass numbers 40 (96.97%), 42 (0.64%), 43 (0.14%), 44 (2.06%), 46 (0.003%), 48 (0 ,nineteen%); strontium - 84 (0.56%), 86 (9.86%), 87 (7.02%), 88 (82.56%); barium - 130 (0.1%), 132 (0.1%), 134 (2.42%), 135 (6.59%), 136 (7.81), 137 (11, 32%), 138 (71.66). Radium is radioactive. The most stable natural isotope is 226Ra. The main minerals of alkaline earth elements are carbon and sulfate salts: CaCO3 - calcite, CaSO4 - andhydrite, SrCO3 - strontianite, SrSO4 - celestine, BaCO3 - witherite. BaSO4 is a heavy spar. Fluorite CaF2 is also a useful mineral.
Ca plays an important role in life processes. The human body contains 0.7-1.4 wt.% Calcium, 99% of which is bone and dental tissue. Plants also contain large amounts of calcium. Calcium compounds are found in natural waters and soil. Barium, strontium and radium are contained in the human body in insignificant quantities.
Obtaining alkaline earth metals
First, oxides or chlorides of E.O are obtained. EO is obtained by calcining ECO3, and ECl2 by the action of hydrochloric acid on ECO3. All alkaline earth metals can be obtained by alumothermal reduction of their oxides at a temperature of 1200 ° C according to the approximate scheme: 3EO + 2Al = Al2O3 + 3E. The process is carried out in a vacuum to avoid oxidation of E. Calcium (like all other E) can be obtained by electrolysis of CaCl2 melt, followed by distillation in vacuum or thermal dissociation of CaC2. Ba and Sr can be obtained by pyrolysis of E2N3, E (NH3) 6, EH2. Radium is mined along the way from uranium ores.
Physical properties of alkaline earth metals
Ca and its analogs are silvery white metals. Calcium is the hardest of them all. Strontium and especially barium are much softer than calcium. All alkaline earth metals are ductile and lend themselves well to forging, cutting and rolling. Calcium under normal conditions crystallizes in the fcc structure with a period of a = 0.556 nm (CN = 12), and at temperatures above 464 ° C in the bcc structure. Ca forms alloys with Li, Mg, Pb, Cu, Cd, Al, Ag, Hg. Strontium has an HCC - structure; at a temperature of 488 ° C, strontium undergoes a polymorphic transformation and crystallizes in a hexagonal structure. It is paramagnetic. Barium crystallizes in the bcc structure. Ca and Sr are capable of forming a continuous series of solid solutions among themselves, and delamination areas appear in the Ca-Ba and Sr-Ba systems. In the liquid state, strontium is mixed with Be, Hg, Ga, In, Sb, Bi, Tl, Al, Mg, Zn, Sn, Pb. With the latter four, Sr forms intermetallic compounds. The electrical conductivity of alkaline earth metals decreases with increasing pressure, contrary to the reverse process for other typical metals. Below are some of the constants for alkaline earth metals:
Ca Sr Ba Ra
Atomic radius, nm 0.197 0.215 0.221 0.235
Radius of the E2 + ion, nm 0.104 0.127 0.138 0.144
Energy cr. lattice, μJcmol 194.1 164.3 175.8 130
, gcm3 1.54 2.63 3.5 5.5-6
Tm., ОС 852 770 710 800
B.p., оС 1484 1380 1640 1500
Electrical conductivity (Hg = 1) 22 4 2
Heat of fusion kcalg-atom 2.1 2.2 1.8
Heat of vaporization kcalg-atom 36 33 36
Specific heat, J (kg.K) 624 737 191.93 136
Liquefiability Pa-1.10-11 5.92 8.36
Chemical properties of alkaline earth metals and their compounds
The fresh surface of E quickly darkens due to the formation of an oxide film. This film is relatively dense - over time, all of the metal is slowly oxidized. The film consists of EO, as well as EO2 and E3N2. Normal electrode potentials of the reactions E-2e = E2 + are = -2.84V (Ca), = -2.89 (Sr). E are very active elements: they dissolve in water and acids, displace most metals from their oxides, halides, sulfides. Primarily (200-300оС) calcium interacts with water vapor according to the scheme: 2Са + Н2О = СаО + СаН2. Secondary reactions are of the form: CaH2 + 2H2O = Ca (OH) 2 + 2H2 and CaO + H2O = Ca (OH) 2. In strong sulfuric acid, E is almost insoluble due to the formation of a film of poorly soluble ESO4. With dilute mineral acids, E reacts violently with the evolution of hydrogen. When heated above 800 ° C, calcium reacts with methane according to the scheme: 3Ca + CH4 = CaH2 + CaC2. When heated, E reacts with hydrogen, sulfur, and gaseous ammonia. In terms of chemical properties, radium is closest to Ba, but it is more active. At room temperature, it noticeably combines with oxygen and nitrogen in the air. In general, its chemical properties are slightly more pronounced than that of its counterparts. All radium compounds slowly decompose under the influence of their own radiation, acquiring a yellowish or brown color. Radium compounds have the property of autoluminescence. As a result of radioactive decay, 1 g of Ra releases 553.7 J of heat every hour. Therefore, the temperature of radium and its compounds is always 1.5 degrees higher than the ambient temperature. It is also known that 1 g of radium per day emits 1 mm3 of radon (226Ra = 222Rn + 4He), which is the basis for its use as a source of radon for radon baths.
E hydrides are white, crystalline salt-like substances. They are obtained directly from the elements by heating. The temperatures of the onset of the reaction E + H2 = EH2 are equal to 250 ° C (Ca), 200 ° C (Sr), 150 ° C (Ba). Thermal dissociation of EN2 begins at 600 ° C. In the hydrogen atmosphere, CaH2 does not decompose at the melting point (816оС). In the absence of moisture, alkaline earth metal hydrides are stable in air at ambient temperatures. They do not react with halogens. However, when heated, the chemical activity of EN2 increases. They are able to reduce oxides to metals (W, Nb, Ti, Ce, Zr, Ta), for example 2CaH2 + TiO2 = 2CaO + 2H2 + Ti. The reaction of CaH2 with Al2O3 proceeds at 750 ° C: 3CaH2 + Al2O3 = 3CaO + 3H2 + 2Al, and then: CaH2 + 2Al = CaAl2 + H2. CaH2 reacts with nitrogen at 600 ° C according to the following scheme: 3CaH2 + N2 = Ca3N2 + 3H2. When EN2 is ignited, they burn out slowly: EN2 + O2 = H2O + CaO. Explosive when mixed with solid oxidizing agents. Under the action of water on EN2, hydroxide and hydrogen are released. This reaction is highly exothermic: EN2 moistened with water in air ignites spontaneously. EN2 reacts with acids, for example, according to the scheme: 2HCl + CaH2 = CaCl2 + 2H2. EN2 is used to obtain pure hydrogen, as well as to determine traces of water in organic solvents. E nitrides are colorless refractory substances. They are obtained directly from elements at elevated temperatures. They decompose with water according to the scheme: E3N2 + 6H2O = 3E (OH) 2 + 2NH3. E3N2 react when heated with CO according to the scheme: E3N2 + 3CO = 3EO + N2 + 3C. The processes that occur when E3N2 is heated with coal look like this:
E3N2 + 5C = ECN2 + 2ES2; (E = Ca, Sr); Ba3N2 + 6C = Ba (CN) 2 + 2BaC2;
Strontium nitride reacts with HCl to give Sr and ammonium chlorides. E3P2 phosphides are formed directly from elements or by calcining tri-substituted phosphates with coal:
Ca3 (PO4) 2 + 4C = Ca3P2 + 4CO
They are hydrolyzed by water according to the scheme: E3R2 + 6H2O = 2PH3 + 3E (OH) 2. With acids, phosphides of alkaline earth metals give the corresponding salt and phosphine. This is the basis of their application for the production of phosphine in the laboratory.
Complex ammonia compounds E (NH3) 6 are solids with a metallic luster and high electrical conductivity. They are obtained by the action of liquid ammonia on E. They spontaneously ignite in air. Without access to air, they decompose into the corresponding amides: E (NH3) 6 = E (NH2) 2 + 4NH3 + H2. When heated, they vigorously decompose in the same way.
Carbides of alkaline earth metals that are obtained by calcining E with coal are decomposed by water with the release of acetylene: ES2 + 2H2O = E (OH) 2 + C2H2. The reaction with BaC2 proceeds so violently that it ignites in contact with water. The heats of formation of ES2 from the elements for Ca and Ba are 14 and 12 kcal mol. When heated with nitrogen, ES2 gives CaCN2, Ba (CN) 2, SrCN2. There are known silicides (ESi and ESi2). They can be obtained by heating directly from the elements. They are hydrolyzed by water and react with acids to give H2Si2O5, SiH4, the corresponding compound E, and hydrogen. Known borides EV6 obtained from the elements when heated.
Calcium oxides and its analogs are white refractory (TbCaO = 2850оС) substances that energetically absorb water. This is the basis for the use of BaO to obtain absolute alcohol. They react violently with water, releasing a lot of heat (except for SrO, the dissolution of which is endothermic). EOs dissolve in acids and ammonium chloride: EO + 2NH4Cl = SrCl2 + 2NH3 + H2O. EO is obtained by calcining carbonates, nitrates, peroxides or hydroxides of the corresponding metals. The effective charges of barium and oxygen in BaO are 0.86. SrO at 700 ° C reacts with potassium cyanide:
KCN + SrO = Sr + KCNO.
Strontium oxide dissolves in methanol to form Sr (OCH3) 2. With the thermal magnesium reduction of BaO, an intermediate oxide Ba2O can be obtained, which is unstable and disproportionate.
Alkaline earth metal hydroxides are white, water-soluble substances. They are strong foundations. In the Ca-Sr-Ba series, the basic character and solubility of hydroxides increase. pPR (Ca (OH) 2) = 5.26, pPR (Sr (OH) 2) = 3.5, pPR (Ba (OH) 2) = 2.3. Ba (OH) 2.8H2O, Sr (OH) 2.8H2O, Ca (OH) 2.H2O are usually isolated from hydroxide solutions. EO add water to form hydroxides. This is the basis for the use of CaO in construction. A close mixture of Ca (OH) 2 and NaOH in a weight ratio of 2: 1 is called soda lime, and is widely used as a CO2 absorber. Ca (OH) 2, when standing in air, absorbs CO2 according to the scheme: Ca (OH) 2 + CO2 = CaCO3 + H2O. Ca (OH) 2 reacts with carbon monoxide at about 400 ° C: CO + Ca (OH) 2 = CaCO3 + H2. Barite water reacts with СS2 at 100 оС: СS2 + 2Ва (ОН) 2 = ВаСО3 + Ва (НS) 2 + Н2О. Aluminum reacts with barite water: 2Al + Ba (OH) 2 + 10H2O = Ba2 + 3H2. E (OH) 2 are used to open carbonic anhydride.
E forms white peroxides. They are much less stable than oxides and are strong oxidizing agents. Of practical importance is the most stable BaO2, which is a white, paramagnetic powder with a density of 4.96 g1 cm3, etc. 450 °. BaО2 is stable at ordinary temperatures (it can be stored for years), poorly soluble in water, alcohol and ether, soluble in dilute acids with the release of salt and hydrogen peroxide. The thermal decomposition of barium peroxide is accelerated by the oxides, Cr2O3, Fe2O3 and CuO. Barium peroxide reacts when heated with hydrogen, sulfur, carbon, ammonia, ammonium salts, potassium ferricyanide, etc. Barium peroxide reacts with concentrated hydrochloric acid, releasing chlorine: BaO2 + 4HCl = BaCl2 + Cl2 + 2H2O. It oxidizes water to hydrogen peroxide: H2O + BaO2 = Ba (OH) 2 + H2O2. This reaction is reversible and in the presence of even carbonic acid the equilibrium is shifted to the right. ВаО2 is used as a starting product for obtaining Н2О2, as well as an oxidizing agent in pyrotechnic compositions. However, BaO2 can also act as a reducing agent: HgCl2 + BaO2 = Hg + BaCl2 + O2. BaO2 is obtained by heating BaO in a stream of air up to 500 ° C according to the following scheme: 2BaO + O2 = 2VaO2. When the temperature rises, the opposite process takes place. Therefore, when Ba burns, only oxide is released. SrO2 and CaO2 are less stable. The general method for obtaining EO2 is the interaction of E (OH) 2 with H2O2, with the release of EO2.8H2O. Thermal decomposition of EO2 begins at 380 oC (Ca), 480 oC (Sr), 790 oC (Ba). When EO2 is heated with concentrated hydrogen peroxide, yellow unstable substances - EO4 superoxides - can be obtained.
E salts are usually colorless. Chlorides, bromides, iodides and nitrates are readily soluble in water. Fluorides, sulfates, carbonates and phosphates are poorly soluble. Ion Ba2 + is toxic. Halides E are divided into two groups: fluorides and all the others. Fluorides are almost insoluble in water and acids, and do not form crystalline hydrates. On the other hand, chlorides, bromides, and iodides are readily soluble in water and are released from solutions in the form of crystalline hydrates. Some properties of EG2 are presented below:
CaF2 CaCl2 CaBr2 CaI2 SrF2 SrCl2 SrBr2 SrI2 BaF2 BaCl2 BaBr2 BaI2
Warm. sample, kcalmol. 290 191 164 128 189 198 171 134 286 205 181 145
Ecr. lattice, kcalmol. 617 525 508 487 580 504 489 467 547 468 463 440
Tm., ОС 1423 782 760 575 1473 872 643 515 1353 962 853 740
B.p., оС 2500 2000 1800 718 2460 2030 2260 1830
D (EG) in pairs, nm. 2.1 2.51 2.67 2.88 2.20 2.67 2.82 3.03 2.32 2.82 2.99 3.20
When obtained by exchange decomposition in solution, fluorides are released in the form of voluminous mucous sediments, which quite easily form colloidal solutions. EG2 can be obtained by acting on the corresponding E. The degrees of dissociation according to the scheme ESl2 = E2 + + 2Cl– are equal: BeCl2 - 0.009%, MgCl2 - 14.6%, CaCl2 - 43.3%, SrCl2 - 60.6%, BaCl2 - 80.2%. Halides (except for fluorides) E contain water of crystallization: CaCl2.6H2O, SrCl2.6H2O and BaCl2.2H2O. X-ray structural analysis established the structure of E [(OH2) 6] G2 for Ca and Sr crystal hydrates. With slow heating of EG2 crystalline hydrates, anhydrous salts can be obtained. CaCl2 easily forms supersaturated solutions. Natural CaF2 (fluorite) is used in the ceramic industry, and it is also used for the production of HF and is a fluorine mineral. Anhydrous CaCl2 is used as a desiccant due to its hydroscopic nature. Crystalline hydrate of calcium chloride is used for the preparation of refrigeration mixtures. BaCl2 - used in cx and for opening SO42- (Ba2 + + SO42- = BaSO4). By fusion of EG2 and EN2, hydrohalides can be obtained: EG2 + EN2 = 2ENH. These substances melt without decomposition but are hydrolyzed by water: 2ENH + 2H2O = EG2 + 2H2 + E (OH) 2. The water solubility of chlorates, bromates and iodates in water decreases along the series of Ca - Sr - Ba and Cl - Br - I. Ba (ClO3) 2 - is used in pyrotechnics. Perchlorates E are readily soluble not only in water but also in organic solvents. The most important of the E (ClO4) 2 is Ba (ClO4) 2.3H2O. Anhydrous barium perchlorate is a good desiccant. Its thermal decomposition begins only at 400 ° C. Calcium hypochlorite Ca (ClO) 2.nH2O (n = 2,3,4) is obtained by the action of chlorine on milk of lime. It is an oxidizing agent and is highly soluble in water. Chlorine lime can be obtained by acting on solid slaked lime with chlorine. It decomposes with water and smells like chlorine in the presence of moisture. Reacts with CO2 air:
CO2 + 2CaOCl2 = CaCO3 + CaCl2 + Cl2O.
Bleach is used as an oxidizing agent, bleach and disinfectant.
For alkaline earth metals, azides E (N3) 2 and thiocyanates E (CNS) 2.3H2O are known. Azidas are much less explosive than lead azide. Rodanides easily lose water when heated. They are highly soluble in water and organic solvents. Ba (N3) 2 and Ba (CNS) 2 can be used to obtain azides and thiocyanates of other metals from sulfates by exchange reaction.
Calcium and strontium nitrates usually exist in the form of crystalline hydrates Ca (NO3) 2.4H2O and Sr (NO3) 2.4H2O. The formation of crystalline hydrate is not characteristic of barium nitrate. When Ca (NO3) 2.4H2O and Sr (NO3) 2.4H2O are heated, I easily lose water. In an inert atmosphere, E nitrates are thermally stable up to 455 oC (Ca), 480 oC (Sr), 495 oC (Ba). The melt of calcium nitrate crystalline hydrate has an acidic environment at 75 ° C. A feature of barium nitrate is the low rate of dissolution of its crystals in water. Only barium nitrate, for which the unstable K2 complex is known, is prone to complexation. Calcium nitrate is soluble in alcohols, methyl acetate, acetone. Strontium and barium nitrates are almost insoluble there. The melting points of E nitrates are estimated at 600 ° C, however, at the same temperature, decomposition begins: E (NO3) 2 = E (NO2) 2 + O2. Further decomposition occurs at a higher temperature: E (NO2) 2 = EO + NO2 + NO. E nitrates have long been used in pyrotechnics. Highly volatile salts E color the flame in the corresponding colors: Ca - in orange-yellow, Sr - in red-carmine, Ba - in yellow-green. Let's understand the essence of this using the example of Sr: Sr2 + has two HAOs: 5s and 5p or 5s and 4d. Let us give energy to this system - we will heat it up. Electrons from orbitals closer to the nucleus will transfer to these HLWs. But such a system is not stable and will release energy in the form of a quantum of light. It is Sr2 + that emits quanta with a frequency corresponding to the lengths of red waves. When receiving pyrotechnic compositions, it is convenient to use saltpeter, because it not only colors the flame, but is also an oxidizing agent, releasing oxygen when heated. Pyrotechnic compositions consist of a solid oxidizing agent, a solid reducing agent, and some organic substances that discolor the reducing agent flame and act as a binding agent. Calcium nitrate is used as a fertilizer.
All phosphates and hydrogen phosphates of E are poorly soluble in water. They can be obtained by dissolving an appropriate amount of CaO or CaCO3 in phosphoric acid. They also precipitate during exchange reactions such as: (3-x) Ca2 + + 2HxPO4- (3-x) = Ca (3-x) (HxPO4) 2. Monosubstituted calcium orthophosphate, which, along with Ca (SO4), is part of superphosphate, is of practical importance (as a fertilizer). It is received according to the scheme:
Ca3 (PO4) 2 + 2H2SO4 = Ca (H2PO4) 2 + 2CаSO4
Oxalates are also slightly soluble in water. Of practical importance is calcium oxalate, which dehydrates at 200 ° C, and decomposes at 430 ° C according to the scheme: CaC2O4 = CaCO3 + CO. Acetates E are released in the form of crystalline hydrates and are readily soluble in water.
Sulfates E are white, poorly soluble in water substances. The solubility of CaSO4.2H2O per 1000 g of water at ordinary temperature is 8.10-3 mol, SrSO4 - 5.10-4 mol, BaSO4 - 1.10-5 mol, RaSO4 - 6.10-6 mol. In the Ca - Ra series, the solubility of sulfates rapidly decreases. Ba2 + is a sulfate ion reagent. Calcium sulphate contains crystallization water. Above 66 ° C, anhydrous calcium sulfate is released from the solution, below - gypsum CaSO4.2H2O. Heating gypsum above 170 ° C is accompanied by the release of hydrated water. When gypsum is mixed with water, this mass quickly hardens due to the formation of crystalline hydrate. This property of gypsum is used in construction. The Egyptians used this knowledge 2000 years ago. The solubility of ESO4 in strong sulfuric acid is much higher than in water (BaSO4 up to 10%), which indicates complexation. The corresponding complexes ESO4.H2SO4 can be obtained in a free state. Double salts with alkali metal and ammonium sulfates are known only for Ca and Sr. (NH4) 2 is soluble in water and is used in analytical chemistry to separate Ca from Sr, because (NH4) 2 is slightly soluble. Gypsum is used for the combined production of sulfuric acid and cement, because when heated with a reducing agent (coal), gypsum decomposes: CaSO4 + C = CaO + SO2 + CO. At a higher temperature (900 oC) sulfur is reduced even more according to the scheme: CaSO4 + 3C = CaS + CO2 + 2CO. A similar decomposition of Sr and Ba sulfates begins at higher temperatures. BaSO4 is non-toxic and is used in medicine and in the production of mineral paints.
Sulfides E are white solids that crystallize like NaCl. The heats of their formation and the energies of the crystal lattices are equal (kcalmol): 110 and 722 (Ca), 108 and 687 (Sr), 106 and 656 (Ba). They can be obtained by synthesis from elements by heating or by calcining sulfates with coal: ESO4 + 3C = ES + CO2 + 2CO. CaS is the least soluble (0.2 hl). ES enters into the following reactions when heated:
ES + H2O = EO + H2S; ES + G2 = S + EG2; ES + 2O2 = ESO4; ES + xS = ESx + 1 (x = 2,3).
Sulfides of alkaline earth metals in a neutral solution are completely hydrolyzed according to the scheme: 2ES + 2H2O = E (HS) 2 + E (OH) 2. Acid sulfides can also be obtained in a free state by evaporation of a solution of sulfides. They react with sulfur:
E (HS) 2 + xS = ESx + 1 + H2S (x = 2,3,4).
Of crystalline hydrates, BaS.6H2O and Ca (HS) 2.6H2O, Ba (HS) 2.4H2O are known. Ca (HS) 2 is used for hair removal. ES are subject to the phenomenon of phosphorescence. There are known polysulfides E: ES2, ES3, ES4, ES5. They are obtained by boiling a suspension of ES in water with sulfur. ES are oxidized in air: 2ES + 3O2 = 2ESO3. By passing air through a CaS suspension, CaS thiosulfate can be obtained according to the following scheme: 2CaS + 2O2 + H2O = Ca (OH) 2 + CaS2O3. It is highly soluble in water. In the Ca - Sr - Ba series, the solubility of thiosulfates decreases. Tellurides E are slightly soluble in water and are also susceptible to hydrolysis, but to a lesser extent than sulfides.
The solubility of E chromates in the Ca - Ba series decreases as sharply as in the case of sulfates. These yellow substances are obtained by the interaction of soluble salts of E with chromates (or dichromates) of alkali metals: E2 + + CrO42- = ECrO4. Calcium chromate is released in the form of crystalline hydrate - CaCrO4.2H2O (pPR CaCrO4 = 3.15). It loses water even before its melting point. SrCrO4 and ВаCrO4 do not form crystalline hydrates. pSP SrCrO4 = 4.44, pSP BaCrO4 = 9.93.
E carbonates are white, poorly soluble in water substances. When heated, ECO3 transforms into EO, removing CO2. In the Ca - Ba series, the thermal stability of carbonates increases. The most practically important of these is calcium carbonate (limestone). It is directly used in construction, and also serves as a raw material for the production of lime and cement. The annual world extraction of lime from limestone is estimated at tens of millions of tons. Thermal dissociation of CaCO3 is endothermic: CaCO3 = CaO + CO2 and requires 43 kcal per mole of limestone. Calcination of CaCO3 is carried out in shaft furnaces. A by-product of roasting is valuable carbon dioxide. CaO is an important building material. When mixed with water, crystallization occurs due to the formation of hydroxide, and then carbonate according to the following schemes:
CaO + H2O = Ca (OH) 2 and Ca (OH) 2 + CO2 = CaCO3 + H2O.
A colossally important practical role is played by cement - a greenish-gray powder consisting of a mixture of various silicates and calcium aluminates. Mixed with water, it hardens through hydration. During its production, a mixture of CaCO3 with clay is fired to the start of sintering (1400-1500 ° C). Then the mixture is ground. The composition of the cement can be expressed by the percentage ratio of the components CaO, SiO2, Al2O3, Fe2O3, and CaO is the base, and everything else is acid anhydrides. The composition of silicate (Portlad) cement is composed mainly of Ca3SiO5, Ca2SiO4, Ca3 (AlO3) 2 and Ca (FeO2) 2. Its seizure takes place according to the schemes:
Ca3SiO5 + 3H2O = Ca2SiO4.2H2O + Ca (OH) 2
Ca2SiO4 + 2Н2О = Ca2SiO4.2Н2О
Ca3 (AlO3) 2 + 6Н2О = Ca3 (AlO3) 2.6Н2О
Ca (FeO2) 2 + nH2O = Ca (FeO2) 2.nH2O.
Natural chalk is added to the composition of various putties. Fine-crystalline CaCO3 precipitated from a solution is included in the composition of tooth powders. BaO is obtained from BaCO3 by calcining with coal according to the scheme: BaCO3 + C = BaO + 2CO. If the process is carried out at a higher temperature in a stream of nitrogen, barium cyanide is formed: ВаСО3 + 4С + N2 = 3CO + Ba (CN) 2. Ba (CN) 2 is readily soluble in water. Ba (CN) 2 can be used for the production of cyanides of other metals by exchange decomposition with sulfates. Bicarbonates E are soluble in water and can be obtained only in solution, for example, by passing carbon dioxide into a suspension of CaCO3 in water: CO2 + CaCO3 + H2O = Ca (HCO3) 2. This reaction is reversible and shifts to the left when heated. The presence of calcium and magnesium bicarbonates in natural waters determines the hardness of the water.
Water hardness and how to eliminate it
Soluble calcium and magnesium salts determine the overall hardness of the water. If they are present in water in small quantities, then the water is called soft. With a high content of these salts (100 - 200 mg of calcium salts - in 1 liter. In terms of ions), the water is considered hard. In such water, soap does not foam well, since calcium and magnesium salts form insoluble compounds with it. In hard water, food products are poorly boiled, and when boiled, it gives scale on the walls of household utensils and steam boilers. Scale has low thermal conductivity, causes an increase in fuel consumption or power consumption of an electrical appliance and accelerates the wear of the walls of the vessel for boiling water. When heated, acidic calcium and magnesium carbonates decompose and transform into insoluble basic carbonates: Ca (HCO3) = H2O + CO2 + CaCO3 ↓ The solubility of calcium sulfate CaSO4 also decreases when heated, therefore it is part of the scale. The hardness caused by the presence of acidic calcium and magnesium carbonates in the water is called carbonate or temporary, because it can be removed. In addition to carbonate hardness, non-carbonate hardness is also distinguished, which depends on the content of ESl2 and ESO4 in water, where E is Ca, Mg. These salts are not removed by boiling, and therefore non-carbonate hardness is also called permanent hardness. Carbonate and non-carbonate hardness add up to total hardness. To eliminate it completely, water is sometimes distilled. But it's expensive. To eliminate carbonate hardness, water can be boiled, but this is also expensive and scale forms. The hardness is removed by adding the appropriate amount of Ca (OH) 2: Ca (OH) 2 + Ca (HCO3) 2 = CaCO3 ↓ + 2H2O. The general hardness is eliminated either by adding Na2CO3, or with the help of so-called cation exchangers. When using sodium carbonate, soluble calcium and magnesium salts are also converted into insoluble carbonates: Ca2 + + Na2CO3 = 2Na + + CaCO3 ↓. Elimination of stiffness with cation exchangers is a more perfect process. Cation exchangers - high molecular weight sodium organic compounds, the composition of which can be expressed by the formula Na2R, where R is a complex acid residue. When water is filtered through a layer of cation exchanger, Na + cations of the crystal lattice are exchanged for Ca2 + and Mg2 + cations from the solution according to the scheme: Ca2 + + Na2R = 2Na + + CaR. Consequently, Ca ions from the solution pass into the cation exchanger, and the Na + ions pass from the cation exchanger into the solution. To restore the used cation exchanger, it is washed with a concentrated solution of sodium chloride. In this case, the opposite process occurs: Ca2 + ions in the crystal lattice in the cation exchanger are replaced by Na + ions from the solution. The regenerated cation exchanger is again used for water purification. Filters based on permutite work in a similar way:
Na2 + Ca2 + = 2Na + + Ca
Being in nature
Due to its high chemical activity, free calcium is not found in nature.
Most of the calcium is contained in silicates and aluminosilicates of various rocks (granites, gneisses, etc.), especially in feldspar - anorthite Ca.
In the form of sedimentary rocks, calcium compounds are represented by chalk and limestone, consisting mainly of the mineral calcite (CaCO3). The crystalline form of calcite - mra-mor - is much less common in nature.
Calcium minerals such as calcite CaCO3, anhydrite CaSO4, alabaster CaSO4 0.5H2O and gypsum CaSO4 2H2O, fluorite CaF2, apatite Ca5 (PO4) 3 (F, Cl, OH), dolomite MgCO3 CaCO3 are quite widespread. The presence of calcium and magnesium salts in natural water determines its hardness.
Calcium, vigorously migrating in the earth's crust and accumulating in various geochemical systems, forms 385 minerals (fourth in the number of minerals).
Calcium accounts for 3.38% of the mass of the earth's crust (5th most abundant after oxygen, silicon, aluminum and iron). The content of the element in seawater is 400 mg / l.
Isotopes
Calcium occurs naturally in the form of a mixture of six isotopes: 40Ca, 42Ca, 43Ca, 44Ca, 46Ca and 48Ca, among which the most common - 40Ca - is 96.97%.
Of the six natural isotopes of calcium, five are stable. The sixth isotope 48Ca, the heaviest of the six and very rare (its isotopic abundance is only 0.187%), was recently discovered to undergo double beta decay with a half-life of 5.3 × 1019 years.
Receiving
Free metallic calcium is obtained by electrolysis of a melt consisting of CaCl2 (75-80%) and KCl or from CaCl2 and CaF2, as well as by aluminothermic reduction of CaO at 1170-1200 ° C:
4CaO + 2Al → CaAl2O4 + 3Ca.
Chemical properties
Calcium is a typical alkaline earth metal. The reactivity of calcium is high, but lower than that of all other alkaline earth metals. It easily interacts with oxygen, carbon dioxide and moisture in the air, which is why the surface of metallic calcium is usually dull gray, therefore, in the laboratory, calcium is usually stored, like other alkaline earth metals, in a tightly closed jar under a layer of kerosene or liquid paraffin. ...
In the series of standard potentials, calcium is located to the left of hydrogen. The standard electrode potential of the Ca2 + / Ca0 pair is -2.84 V, so that calcium actively reacts with water, but without ignition:
Ca + 2H2O → Ca (OH) 2 + H2 + Q.
Calcium reacts with active non-metals (oxygen, chlorine, bromine) under normal conditions:
2Са + О2 → 2СаО
Ca + Br2 → CaBr2.
When heated in air or oxygen, calcium ignites. Calcium reacts with less active non-metals (hydrogen, boron, carbon, silicon, nitrogen, phosphorus and others) when heated, for example:
Ca + H2 → CaH2, Ca + 6B = CaB6,
3Ca + N2 → Ca3N2, Ca + 2C → CaC2,
3Ca + 2P → Ca3P2 (calcium phosphide), calcium phosphides of the compositions CaP and CaP5 are also known;
2Ca + Si → Ca2Si (calcium silicide), calcium silicides of the compositions CaSi, Ca3Si4 and CaSi2 are also known.
The course of the above reactions, as a rule, is accompanied by the release of a large amount of heat (that is, these reactions are exothermic). In all compounds with non-metals, the oxidation state of calcium is +2. Most of the calcium compounds with non-metals are readily decomposed by water, for example:
CaH2 + 2H2O → Ca (OH) 2 + 2H2,
Ca3N2 + 6Н2О → 3Са (ОН) 2 + 2NH3.
The Ca2 + ion is colorless. When soluble calcium salts are introduced into the flame, the flame turns brick-red.
Calcium salts such as chloride CaCl2, bromide CaBr2, iodide CaI2 and nitrate Ca (NO3) 2 are readily soluble in water. Fluoride CaF2, carbonate CaCO3, sulfate CaSO4, orthophosphate Ca3 (PO4) 2, oxalate CaC2O4 and some others are insoluble in water.
Of great importance is the fact that, unlike calcium carbonate CaCO3, acidic calcium carbonate (bicarbonate) Ca (HCO3) 2 is soluble in water. In nature, this leads to the following processes. When cold rain or river water saturated with carbon dioxide penetrates underground and falls on limestones, their dissolution is observed:
CaCO3 + CO2 + H2O → Ca (HCO3) 2.
In the same places where water saturated with calcium bicarbonate comes out to the surface of the earth and is heated by the sun's rays, the opposite reaction takes place:
Ca (HCO3) 2 → CaCO3 + CO2 + H2O.
This is how large masses of substances are transferred in nature. As a result, huge gaps can form underground, and beautiful stone "icicles" - stalactites and stalagmites - form in the caves.
The presence of dissolved calcium bicarbonate in water largely determines the temporary hardness of the water. It is called temporary because when boiling water, bicarbonate decomposes, and CaCO3 precipitates. This phenomenon leads, for example, to the fact that scale forms in the teapot over time.
Application of metallic calcium
The main use of metallic calcium is as a reducing agent in the production of metals, especially nickel, copper and stainless steel. Calcium and its hydride are also used to obtain hard-to-reduce metals such as chromium, thorium and uranium. Calcium lead alloys are used in batteries and bearing alloys. Calcium granules are also used to remove traces of air from vacuum equipment.
Biological role
Calcium is a common macronutrient in plants, animals and humans. In humans and other vertebrates, most of it is contained in the skeleton and teeth in the form of phosphates. The skeletons of most invertebrate groups (sponges, coral polyps, molluscs, etc.) are composed of various forms of calcium carbonate (lime). Calcium ions are involved in blood clotting processes, as well as in ensuring a constant osmotic pressure of the blood. Calcium ions also serve as one of the universal secondary mediators and regulate a variety of intracellular processes - muscle contraction, exocytosis, including the secretion of hormones and neurotransmitters, etc. The calcium concentration in the cytoplasm of human cells is about 10-7 mol, in intercellular fluids about 10 3 mol.
STRONTIUM
Being in nature
Strontium is not found in free form. It is found in about 40 minerals. Of these, the most important is celestine SrSO4 (51.2% Sr). Strontianite SrCO3 (64.4% Sr) is also mined. These two minerals are of industrial importance. Most often, strontium is present as an impurity in various calcium minerals.
Other strontium minerals include:
SrAl3 (AsO4) SO4 (OH) 6 - kemmlicite;
Sr2Al (CO3) F5 - stenonite;
SrAl2 (CO3) 2 (OH) 4 H2O - strontium dresserite;
SrAl3 (PO4) 2 (OH) 5 H2O - goyasite;
Sr2Al (PO4) 2OH - buzzenite;
SrAl3 (PO4) SO4 (OH) 6 - svanbergite;
Sr (AlSiO4) 2 - slosonite;
Sr (AlSi3O8) 2 5H2O - brewsterite;
Sr5 (AsO4) 3F - fermorite;
Sr2 (B14O23) 8H2O - stronziojinorite;
Sr2 (B5O9) Cl H2O - stronziohilgardite;
SrFe3 (PO4) 2 (OH) 5 H2O - lusunite;
SrMn2 (VO4) 2 4H2O - santafeite;
Sr5 (PO4) 3OH - white;
SrV (Si2O7) - haradite.
According to the level of physical prevalence in the earth's crust, strontium takes the 23rd place - its mass fraction is 0.014% (in the lithosphere - 0.045%). The mole fraction of metal in the earth's crust is 0.0029%. Strontium is found in seawater (8 mg / l).
In nature, strontium occurs as a mixture of 4 stable isotopes 84Sr (0.56%), 86Sr (9.86%), 87Sr (7.02%), 88Sr (82.56%).
Receiving
There are 3 ways to obtain metallic strontium:
thermal decomposition of some compounds
electrolysis
reduction of oxide or chloride.
Chemical properties
Strontium in its compounds always exhibits a valency of +2. In terms of properties, strontium is close to calcium and barium, occupying an intermediate position between them.
In the electrochemical series of voltages, strontium is among the most active metals (its normal electrode potential is −2.89 V. It reacts vigorously with water, forming a hydroxide:
Sr + 2H2O = Sr (OH) 2 + H2
Interacts with acids, displaces heavy metals from their salts. It reacts weakly with concentrated acids (H2SO4, HNO3).
Metallic strontium is rapidly oxidized in air, forming a yellowish film, in which, in addition to SrO oxide, SrO2 peroxide and Sr3N2 nitride are always present. When heated in air, it ignites; powdered strontium in air is prone to spontaneous combustion.
Reacts vigorously with non-metals - sulfur, phosphorus, halogens. Reacts with hydrogen (above 200 ° C), nitrogen (above 400 ° C). Virtually does not react with alkalis.
At high temperatures, it reacts with CO2 to form carbide:
5Sr + 2CO2 = SrC2 + 4SrO
Easily soluble strontium salts with Cl–, I–, NO3– anions. Salts with anions F−, SO42−, CO32−, PO43− are slightly soluble.
Application
The main fields of application of strontium and its chemical compounds are the radio-electronic industry, pyrotechnics, metallurgy, and the food industry.
Metallurgy
Strontium is used for alloying copper and some of its alloys, for introduction into battery lead alloys, for desulfurization of cast iron, copper and steels.
Metallothermy
Strontium with a purity of 99.99-99.999% is used for the reduction of uranium.
Magnetic materials
Hard magnetic strontium ferrites are widely used as materials for the production of permanent magnets.
Pyrotechnics
In pyrotechnics, carbonate, nitrate, strontium perchlorate are used to color the flame in a carmine-red color. The magnesium-strontium alloy has the strongest pyrophoric properties and is used in pyrotechnics for incendiary and signal compositions.
Nuclear power engineering
Strontium uranate plays an important role in the production of hydrogen (strontium-uranate cycle, Los Alamos, USA) by the thermochemical method (atomic-hydrogen energy), and, in particular, methods are being developed for the direct fission of uranium nuclei in strontium uranate to obtain heat during the decomposition of water for hydrogen and oxygen.
Strontium oxide is used as a component of superconducting ceramics.
Chemical power sources
Strontium fluoride is used as a component of solid-state fluorionic storage batteries with enormous energy and energy density.
Alloys of strontium with tin and lead are used for casting down conductors of storage batteries. Strontium-cadmium alloys for electrochemical cell anodes.
Biological role
Effect on the human body
One should not confuse the effect on the human body of natural (non-radioactive, low-toxic and, moreover, widely used for the treatment of osteoporosis) and radioactive strontium isotopes.
Natural strontium is an integral part of microorganisms, plants and animals. Strontium is an analogue of calcium, so it is most effectively deposited in bone tissue. Less than 1% is retained in soft tissues. Strontium accumulates with great speed in the body of children up to the age of four, when bone tissue is being actively formed. Strontium metabolism changes in some diseases of the digestive system and the cardiovascular system.
BARIUM
Being in nature
The content of barium in the earth's crust is 0.05% by weight; in seawater, the average barium content is 0.02 mg / liter. Barium is active, it belongs to the subgroup of alkaline-earth metals and is rather strongly bound in minerals. The main minerals are barite (BaSO4) and witherite (BaCO3).
Rare barium minerals: Celsian or barium feldspar (barium aluminosilicate), hyalofan (mixed aluminosilicate of barium and potassium), nitrobarite (barium nitrate), etc.
Isotopes
Natural barium consists of a mixture of seven stable isotopes: 130Ba, 132Ba, 134Ba, 135Ba, 136Ba, 137Ba, 138Ba. The latter is the most common (71.66%). Radioactive isotopes of barium are also known, the most important of which is 140Ba. It is formed by the decay of uranium, thorium and plutonium.
Receiving
The main raw material for barium production is barite concentrate (80-95% BaSO4), which in turn is obtained by barite flotation. Barium sulfate is further reduced with coke or natural gas:
BaSO4 + 4C = BaS + 4CO
BaSO4 + 2CH4 = BaS + 2C + 4H2O.
Then, when heated, sulfide is hydrolyzed to barium hydroxide Ba (OH) 2 or, under the action of CO2, is converted into insoluble barium carbonate BaCO3, which is then converted into barium oxide BaO (calcination at 800 ° C for Ba (OH) 2 and over 1000 ° C for BaCO3):
BaS + 2H2O = Ba (OH) 2 + H2S
BaS + H2O + CO2 = BaCO3 + H2S
Ba (OH) 2 = BaO + H2O
BaCO3 = BaO + CO2
Metallic barium is obtained from oxide by reduction with aluminum in vacuum at 1200-1250 ° C:
4BaO + 2Al = 3Ba + BaAl2O4.
Barium is purified by vacuum distillation or zone melting.
Chemical properties
Barium is an alkaline earth metal. In air, barium quickly oxidizes, forming a mixture of barium oxide BaO and barium nitride Ba3N2, and ignites with slight heating. Reacts vigorously with water to form barium hydroxide Ba (OH) 2:
Ba + 2H2O = Ba (OH) 2 + H2
Reacts actively with dilute acids. Many barium salts are insoluble or slightly soluble in water: barium sulfate BaSO4, barium sulfite BaSO3, barium carbonate BaCO3, barium phosphate Ba3 (PO4) 2. Barium sulfide BaS, unlike calcium sulfide CaS, is highly soluble in water.
Reacts easily with halogens to form halides.
When heated with hydrogen, it forms barium hydride BaH2, which in turn with lithium hydride LiH gives a Li complex.
Reacts when heated with ammonia:
6Ba + 2NH3 = 3BaH2 + Ba3N2
When heated, barium nitride Ba3N2 reacts with CO to form cyanide:
Ba3N2 + 2CO = Ba (CN) 2 + 2BaO
With liquid ammonia, it gives a dark blue solution, from which ammonia can be isolated, which has a golden sheen and readily decomposes with the elimination of NH3. In the presence platinum catalyst Ammonia decomposes to form barium amide:
= Ba (NH2) 2 + 4NH3 + H2
Barium carbide BaC2 can be obtained by heating in an arc furnace BaO with coal.
Forms phosphide Ba3P2 with phosphorus.
Barium reduces oxides, halides and sulfides of many metals to the corresponding metal.
Application
Anti-corrosion material
Barium is added together with zirconium to liquid metal coolants (alloys of sodium, potassium, rubidium, lithium, cesium) to reduce the aggressiveness of the latter to pipelines, and in metallurgy.
Ferro- and piezoelectric
Barium titanate is used as a dielectric in ceramic capacitors and as a material for piezoelectric microphones and piezoceramic emitters.
Optics
Barium fluoride is used in the form of single crystals in optics (lenses, prisms).
Pyrotechnics
Barium peroxide is used for pyrotechnics and as an oxidizing agent. Barium nitrate and barium chlorate are used in pyrotechnics to color the flame (green fire).
Atomic-hydrogen energy
Barium chromate is used in the production of hydrogen and oxygen by a thermochemical method (Oak Ridge cycle, USA).
High temperature superconductivity
Barium oxide, together with oxides of copper and rare earth metals, is used for the synthesis of superconducting ceramics operating at liquid nitrogen temperatures and above.
Nuclear energy
Barium oxide is used to melt a special type of glass - used to coat uranium rods. One of the widespread types of such glasses has the following composition - (phosphorus oxide - 61%, BaO - 32%, aluminum oxide - 1.5%, sodium oxide - 5.5%). Barium phosphate is also used in glass making for the nuclear industry.
Chemical power sources
Barium fluoride is used in solid-state fluorionic storage batteries as a component of fluoride electrolyte.
Barium oxide is used in powerful copper oxide batteries as a component of the active mass (barium oxide-copper oxide).
Barium sulfate is used as a negative electrode active mass expander in the production of lead-acid batteries.
Prices
Prices for barium metal in ingots with a purity of 99.9% fluctuate around $ 30 per kg.
Biological role
The biological role of barium has not been adequately studied. It is not included in the list of vital trace elements. All soluble barium salts are highly toxic.
RADIUM
Radium (lat. Radium), Ra, a radioactive chemical element of the II group of the periodic system of Mendeleev, atomic number 88. The isotopes of Ra with mass numbers 213, 215, 219-230 are known. The longest-lived is a-radioactive 226Ra with a half-life of about 1600 years. In nature, as members of natural radioactive series, 222Ra (the special name of the isotope is actinium-x, symbol AcX), 224Ra (thorium-x, ThX), 226Ra and 228Ra (mesotorium-I, MsThI) are found.
STORY
The discovery of Ra was reported in 1898 by the spouses P. and M. Curie, together with J. Bemont, shortly after A. Becquerel first (in 1896) discovered the phenomenon of radioactivity on uranium salts. In 1897, M. Sklodowska-Curie, who worked in Paris, established that the intensity of radiation emitted by a uranium resin (uraninite mineral) was much higher than could be expected given the content of uranium in the resin. Sklodowska-Curie suggested that this is due to the presence of yet unknown highly radioactive substances in the mineral. A thorough chemical study of the uranium resin made it possible to discover two new elements - first polonium, and a little later - and R. During the isolation of R., the behavior of the new element was monitored by its radiation, and therefore the element was named from lat. radius - ray. In order to isolate the pure compound R., the Curies in laboratory conditions processed about 1 ton of factory waste left after the extraction of uranium from uranium tar. In particular, at least 10,000 recrystallizations were carried out from aqueous solutions of a mixture of BaCl2 and RaCl2 (barium compounds serve as so-called isomorphic carriers in the extraction of R.). As a result, we managed to obtain 90 mg of pure RaCI2.
Ra is an extremely rare element. In uranium ores, which are its main source, for 1 ton of U there is no more than 0.34 g of Ra. R. belongs to the highly scattered elements and is found in very low concentrations in a wide variety of objects.
All Ra compounds exhibit a pale bluish glow in air. Due to the self-absorption of a- and b-particles emitted during the radioactive decay of 226Ra and its daughter products, each gram of 226Ra releases about 550 J (130 cal) of heat per hour, so the temperature of Ra preparations is always slightly higher than the ambient temperature.
PHYSICAL PROPERTIES
Ra is a silvery-white lustrous metal that tarnishes quickly in air. The lattice is cubic, body-centered, the calculated density is 5.5 g / cm3. According to various sources, tpl. is 700-960 ° С, t boiling point is about 1140 ° С. On the outer electron shell of the R. atom there are 2 electrons (configuration 7s2). Accordingly, R. has only one oxidation state, +2 (valence II). In terms of chemical properties, R. is most similar to barium, but more active. At room temperature, R. combines with oxygen, giving the oxide RaO, and with nitrogen, giving the nitride Ra3N2. R. reacts violently with water, releasing H2, and a strong base Ra (OH) 2 is formed. Chloride, bromide, iodide, nitrate, and sulfide are readily soluble in water; carbonate, sulfate, chromate, and oxalate are poorly soluble.
CHEMICAL PROPERTIES
By chem. St. you radium is similar to Va. Almost all radium compounds are isomorphic to the corresponding comp. Wha. In air, metallic radium quickly becomes covered with a dark film, which is a mixture of nitride and radium oxide. Metallic radium reacts violently with water to form hydroxide Ra (OH) 2 p-rim in water and release H2. The electrode potential of radium release from water solutions is -1.718V (in relation to the normal calomel electrode).
Compounds of radium have a holy self-luminescence-glow in the dark due to their own. radiation. Mn. colorless radium salts, but when decomposed under the action. own. radiation becomes yellow or brown. Good sol. in water RaCl2 (mp 900 ° C, density 4.91 g / cm3; see also table), RaBr2 (mp 728 ° C, density 5.79 g / cm3), RaI2 and Ra (NO3) 2. Better than other sol. in water RaBr2 (70 g in 100 g at 20 ° C). Radium chloride and bromide crystallize from water in the form of crystalline hydrates with two or six H2O molecules. Low soluble compounds - sulfate RaSO4 (about 2 10-4 g in 100 g of water at 20 ° С), iodate Ra (IO3) 2, fluoride RaF2, chromate RaCrO4, carbonate RaCO3 and oxalate RaC2O4. Known complexes of radium with lemon, wine, apple, milk, ethylenediaminetetraacetic to-tami and other ligands. Compared with other shch.-earth. metals, radium has a weaker tendency to complex formation.
Radium is isolated in the form of RaCl2 or other salts as a by-product of processing uranium ores (after extracting U from them) using methods of precipitation, fractional crystallization, and ion exchange; metallic radium is obtained by electrolysis of RaCl2 solution on a mercury cathode, reduction of RaO with aluminum at heating. in a vacuum.
APPLICATION
The study of the properties of Ra has played a huge role in the development of scientific knowledge, because made it possible to clarify many issues related to the phenomenon of radioactivity. For a long time, Ra was the only element whose radioactive properties found practical application (in medicine; for the preparation of luminous compounds, etc.). However, now in most cases it is more profitable to use not Ra, but cheaper artificial radioactive isotopes of other elements. Ra retained some importance in medicine as a source of radon in the treatment of radon baths. In small quantities, R. is used for the preparation of neutron sources (mixed with beryllium) and in the production of light compositions (mixed with zinc sulfide).
BIOLOGICAL ROLE
Radium in the body. Of the natural radioactive isotopes, the long-lived 226Ra has the greatest biological significance. R. is unevenly distributed in different parts of the biosphere. There are geochemical provinces with an increased content of R. The accumulation of R. in the organs and tissues of plants obeys the general laws of absorption of mineral substances and depends on the type of plant and the conditions of its growth. As a rule, there is more R. in the roots and leaves of herbaceous plants than in the stems and reproductive organs; most of all R. is in bark and wood. The average R. content in flowering plants is 0.3-9.0 × 10-11 curie / kg, in mor. algae 0.2-3.2 × 10-11 curie / kg.
It enters the body of animals and humans with food, in which it is constantly present (in wheat 20-26 × 10-15 g / g, in potatoes 67-125 × 10-15 g / g, in meat 8 × 10-15 g / g) , as well as with drinking water. The daily intake of 226Ra into the human body with food and water is 2.3 × 10-12 curies, and the losses with urine and feces are 0.8 × 10-13 and 2.2 × 10-12 curies. About 80% of R. entering the body (it is close in its chemical properties to Ca) accumulates in the bone tissue. The content of R. in the human body depends on the region of residence and the nature of the diet. Large concentrations of R. in the body have a harmful effect on animals and humans, causing painful changes in the form of osteoporosis, spontaneous fractures, and tumors. The content of R. in the soil in excess of 1 × 10-7-10-8 curie / kg noticeably inhibits the growth and development of plants.
Properties of alkaline earth metals
Physical properties
Alkaline earth metals (in comparison with alkali metals) have higher t╟pl. and t╟boil., ionization potentials, densities and hardness.
Chemical properties
1. Very reactive.
2. Possess a positive valence of +2.
3. React with water at room temperature (except Be) to liberate hydrogen.
4. They have a high affinity for oxygen (reducing agents).
5. With hydrogen form salt-like hydrides EH 2.
6. Oxides have the general formula EO. The tendency towards the formation of peroxides is less pronounced than for alkali metals.
Being in nature
3BeO ∙ Al 2 O 3 ∙ 6SiO 2 beryl
Mg
MgCO 3 magnesite
CaCO 3 ∙ MgCO 3 dolomite
KCl ∙ MgSO 4 ∙ 3H 2 O kainite
KCl ∙ MgCl 2 ∙ 6H 2 O carnallite
CaCO 3 calcite (limestone, marble, etc.)
Ca 3 (PO 4) 2 apatite, phosphorite
CaSO 4 ∙ 2H 2 O gypsum
CaSO 4 anhydrite
CaF 2 fluorspar (fluorite)
SrSO 4 celestine
SrCO 3 strontianite
BaSO 4 barite
BaCO 3 witherite
Receiving
Beryllium is obtained by reduction of fluoride:
BeF 2 + Mg t Be + MgF 2
Barium is obtained by oxide reduction:
3BaO + 2Al t 3Ba + Al 2 O 3
The rest of the metals are obtained by electrolysis of chloride melts:
CaCl 2 = Ca + Cl 2 ╜
cathode: Ca 2+ + 2ē = Ca 0
anode: 2Cl - - 2ē = Cl 0 2
MgO + C = Mg + CO
Metals of the main subgroup of group II are strong reducing agents; the compounds exhibit only the oxidation state +2. The activity of metals and their reducing ability increases in the following order: Be Mg Ca Sr Ba╝
1. Reaction with water.
Under normal conditions, the surface of Be and Mg is covered with an inert oxide film, so they are resistant to water. In contrast, Ca, Sr and Ba dissolve in water to form hydroxides, which are strong bases:
Mg + 2H 2 O t Mg (OH) 2 + H 2
Ca + 2H 2 O = Ca (OH) 2 + H 2 ╜
2. Reaction with oxygen.
All metals form oxides RO, barium peroxide BaO 2:
2Mg + O 2 = 2MgO
Ba + O 2 = BaO 2
3.Binary compounds are formed with other non-metals:
Be + Cl 2 = BeCl 2 (halides)
Ba + S = BaS (sulfides)
3Mg + N 2 = Mg 3 N 2 (nitrides)
Ca + H 2 = CaH 2 (hydrides)
Ca + 2C = CaC 2 (carbides)
3Ba + 2P = Ba 3 P 2 (phosphides)
Beryllium and magnesium react relatively slowly with non-metals.
4. All metals dissolve in acids:
Ca + 2HCl = CaCl 2 + H 2 ╜
Mg + H 2 SO 4 (dil.) = MgSO 4 + H 2 ╜
Beryllium also dissolves in aqueous solutions of alkalis:
Be + 2NaOH + 2H 2 O = Na 2 + H 2 ╜
5. Qualitative reaction to cations of alkaline earth metals - coloration of the flame in the following colors:
Ca 2+ - dark orange
Sr 2+ - dark red
Ba 2+ - light green
The Ba 2+ cation is usually opened by an exchange reaction with sulfuric acid or its salts:
Barium sulfate is a white precipitate, insoluble in mineral acids.
Alkaline earth metal oxides
Receiving
1) Oxidation of metals (except for Ba, which forms peroxide)
2) Thermal decomposition of nitrates or carbonates
CaCO 3 ═ t ═ CaO + CO 2 ╜
2Mg (NO 3) 2 t 2MgO + 4NO 2 ╜ + O 2 ╜
Chemical properties
Typical basic oxides. React with water (except BeO), acid oxides and acids
MgO + H 2 O = Mg (OH) 2
3CaO + P 2 O 5 = Ca 3 (PO 4) 2
BeO + 2HNO 3 = Be (NO 3) 2 + H 2 O
BeO - amphoteric oxide, soluble in alkalis:
BeO + 2NaOH + H 2 O = Na 2
Alkaline earth metal hydroxides R (OH) 2
Receiving
Reactions of alkaline earth metals or their oxides with water: Ba + 2H 2 O = Ba (OH) 2 + H 2
CaO (quicklime) + H 2 O = Ca (OH) 2 (slaked lime)
Chemical properties
Hydroxides R (OH) 2 - white crystalline substances, they are less soluble in water than alkali metal hydroxides (the solubility of hydroxides decreases with decreasing serial number; Be (OH) 2 - insoluble in water, soluble in alkalis). The basicity of R (OH) 2 increases with increasing atomic number:
Be (OH) 2 - amphoteric hydroxide
Mg (OH) 2 - weak base
the rest of the hydroxides are strong bases (alkalis).
1) Reactions with acidic oxides:
Ca (OH) 2 + SO 2 = CaSO 3 ¯ + H 2 O
Ba (OH) 2 + CO 2 = BaCO 3 ¯ + H 2 O
2) Reactions with acids:
Mg (OH) 2 + 2CH 3 COOH = (CH 3 COO) 2 Mg + 2H 2 O
Ba (OH) 2 + 2HNO 3 = Ba (NO 3) 2 + 2H 2 O
3) Reactions of exchange with salts:
Ba (OH) 2 + K 2 SO 4 = BaSO 4 ¯ + 2KOH
4) Reaction of beryllium hydroxide with alkalis:
Be (OH) 2 + 2NaOH = Na 2
Hardness of water
Natural water containing Ca 2+ and Mg 2+ ions is called hard. When boiled, hard water forms scale; food products do not boil down in it; detergents do not foam.
Carbonate (temporary) hardness is due to the presence of calcium and magnesium bicarbonates in water, non-carbonate (permanent) hardness is due to chlorides and sulfates.
The total hardness of water is considered as the sum of carbonate and non-carbonate.
Removal of water hardness is carried out by precipitation of Ca 2+ and Mg 2+ ions from a solution:
1) by boiling:
Ca (HCO 3) 2 t CaCO 3 ¯ + CO 2 + H 2 O
Mg (HCO 3) 2 t MgCO 3 ¯ + CO 2 + H 2 O
2) adding milk of lime:
Ca (HCO 3) 2 + Ca (OH) 2 = 2CaCO 3 ¯ + 2H 2 O
3) adding soda:
Ca (HCO 3) 2 + Na 2 CO 3 = CaCO 3 ¯ + 2NaHCO 3
CaSO 4 + Na 2 CO 3 = CaCO 3 ¯ + Na 2 SO 4
MgCl 2 + Na 2 CO 3 = MgCO 3 ¯ + 2NaCl
To remove temporary stiffness, all four methods are used, and for permanent - only the last two.
Thermal decomposition of nitrates.
E (NO3) 2 = t = EO + 2NO2 + 1 / 2O2
Features of chemistry and beryllium.
Be (OH) 2 + 2NaOH (g) = Na2
Al (OH) 3 + 3NaOH (g) = Na3
Be + 2NaOH + 2H2O = Na2 + H2
Al + 3NaOH + 3H2O = Na3 + 3 / 2H2
Be, Al + HNO3 (Conc) = passivation
