BEING IN NATURE
The earth's crust contains beryllium - 0.00053%, magnesium - 1.95%, calcium - 3.38%, strontium - 0.014%, barium - 0.026%, radium is an artificial element.
They are found in nature only in the form of compounds - silicates, aluminosilicates, carbonates, phosphates, sulfates, etc.
RECEIVING
1. Beryllium is obtained by reduction of fluoride:
BeF 2 + Mg t ˚ C → Be + MgF 2
2. Barium is obtained by reduction of the oxide:
3BaO + 2Al t ˚ C → 3Ba + Al 2 O 3
3. The remaining metals are obtained by electrolysis of chloride melts:
Because Since the metals of this subgroup are strong reducing agents, their production is possible only by electrolysis of molten salts. In the case of Ca, CaCl 2 is usually used (with the addition of CaF 2 to lower the melting point)
CaCl 2 =Ca+Cl 2
PHYSICAL PROPERTIES
Alkaline earth metals (compared to alkali metals) have higher temperatures. and temperature, density and hardness.
APPLICATION
| Beryllium (Amphoterene) | Magnesium | Ca, Sr, Ba, Ra |
| 1. Manufacturing of heat-protective structures for space. ships (heat resistance, heat capacity of beryllium) 2. Beryllium bronzes (lightness, hardness, heat resistance, anti-corrosion of alloys, tensile strength higher than steel, can be rolled into strips 0.1 mm thick) 3. In nuclear reactors, X-ray engineering, radio electronics 4. Be alloy , Ni, W- they make watch springs in Switzerland But Be is fragile, poisonous and very expensive | 1. Production of metals - magnesium thermia (titanium, uranium, zirconium, etc.) 2. For the production of ultra-light alloys (aircraft manufacturing, automobile production) 3. In organic synthesis 4. For the production of lighting and incendiary rockets. | 1. Production of lead-cadmium alloys necessary for the production of bearings. 2. Strontium is a reducing agent in uranium production. Phosphors are strontium salts. 3. Used as getters, substances for creating a vacuum in electrical appliances. Calcium Production of rare metals, part of alloys. Barium Absorber in cathode ray tubes. Radium X-ray diagnostics, research work. |
CHEMICAL PROPERTIES
1. Very reactive, strong reducing agents. The activity of metals and their reducing ability increases in the series: Be–Mg–Ca–Sr–Ba
2. They have an oxidation state of +2.
3. React with water at room temperature (except Be) to release hydrogen.
4. With hydrogen they form salt-like hydrides EH 2.
5. Oxides have the general formula EO. The tendency to form peroxides is less pronounced than for alkali metals.
Reaction with water.
Under normal conditions, the surface of Be and Mg is covered with an inert oxide film, so they are stable against water, but with hot water, magnesium forms the base Mg(OH) 2.
In contrast, Ca, Sr and Ba dissolve in water to form hydroxides, which are strong bases:
Be + H 2 O → BeO+ H 2
Ca + 2H 2 O → Ca(OH) 2 + H 2
Reaction with oxygen.
All metals form oxides RO, barium forms 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 nonmetals.
4. All metals dissolve in acids:
Ca + 2HCl → CaCl 2 + H 2
Mg + H 2 SO 4 (diluted) → 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 - coloring 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 discovered by an exchange reaction with sulfuric acid or its salts:
BaCl 2 + H 2 SO 4 → BaSO 4 ↓ + 2HCl
Ba 2+ + SO 4 2- → BaSO 4 ↓
Barium sulfate is a white precipitate, insoluble in mineral acids.
Alkaline earth metal oxides
Receipt
1) Oxidation of metals (except Ba, which forms peroxide)
2) Thermal decomposition of nitrates or carbonates
CaCO 3 t ˚ C → CaO + CO 2
2Mg(NO 3) 2 t˚C → 2MgO + 4NO 2 + O 2
Chemical properties
Typical basic oxides. Reacts with water (except BeO and MgO), acid oxides and acids
CaO + H 2 O → Ca(OH) 2
3CaO + P 2 O 5 → Ca 3 (PO 4) 2
BeO + 2HNO 3 → Be(NO 3) 2 + H 2 O
BeO is an amphoteric oxide, soluble in alkalis:
BeO + 2NaOH + H 2 O → Na 2
Alkaline earth metal hydroxides R(OH) 2
Receipt
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 are white crystalline substances, less soluble in water than alkali metal hydroxides ( the solubility of hydroxides decreases with decreasing atomic 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
Ca(OH) 2 - alkali
the remaining hydroxides are strong bases (alkalis).
1) Reactions with acid oxides:
Ca(OH) 2 + CO 2 → CaСO 3 ↓ + H 2 O! Qualitative reaction to carbon dioxide
Ba(OH) 2 + SO 2 → BaSO 3 ↓ + H 2 O
2) Reactions with acids:
Ba(OH) 2 + 2HNO 3 → Ba(NO 3) 2 + 2H 2 O
3) Exchange reactions 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
Water hardness
Natural water containing Ca 2+ and Mg 2+ ions is called hard water. Hard water forms scale when boiled and food products cannot be cooked in it; Detergents do not produce foam.
Carbonate (temporary) hardness due to the presence of calcium and magnesium bicarbonates in water, non-carbonate (constant) hardness – chlorides and sulfates.
Total water hardness is considered as the sum of carbonate and non-carbonate.
Removing Hardness water is carried out by precipitation of Ca 2+ and Mg 2+ ions from solution
To the family alkaline earth elements include calcium, strontium, barium and radium. D.I. Mendeleev included magnesium in this family. Alkaline earth elements are called because their hydroxides, like alkali metal hydroxides, are soluble in water, i.e., they are alkalis. “...They are called earthy because in nature they are found in the state of compounds that form an insoluble mass of earth, and they themselves, in the form of oxides RO, have an earthy appearance,” Mendeleev explained in “Fundamentals of Chemistry.”
General characteristics of elements of group IIa
Metals of the main subgroup of group II have the electronic configuration of the outer energy level ns², and are s-elements.
Easily donate two valence electrons, and in all compounds have an oxidation state of +2
Strong reducing agents
The activity of metals and their reducing ability increases in the series: Be–Mg–Ca–Sr–Ba
Alkaline earth metals include only calcium, strontium, barium and radium, less often magnesium
Beryllium is closer to aluminum in most properties
Physical properties of simple substances
Alkaline earth metals (compared to alkali metals) have higher temperatures. and boiling point, ionization potentials, densities and hardness.
Chemical properties of alkaline earth metals + Be
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 alkalis:
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. They form binary compounds 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 nonmetals.
4. All alkaline earth metals dissolve in acids:
Ca + 2HCl → CaCl 2 + H 2
Mg + H 2 SO 4 (diluted) → MgSO 4 + H 2
5. Beryllium dissolves in aqueous solutions of alkalis:
Be + 2NaOH + 2H 2 O → Na 2 + H 2
6. Volatile compounds of alkaline earth metals give the flame a characteristic color:
calcium compounds are brick red, strontium compounds are carmine red, and barium compounds are yellowish green.
Beryllium, like lithium, is one of the s-elements. The fourth electron appearing in the Be atom is placed in the 2s orbital. The ionization energy of beryllium is higher than that of lithium due to the higher nuclear charge. In strong bases it forms the beryllate ion BeO 2-2. Consequently, beryllium is a metal, but its compounds are amphoteric. Beryllium, although a metal, is significantly less electropositive compared to lithium.
The high ionization energy of the beryllium atom is noticeably different from other elements of the PA subgroup (magnesium and alkaline earth metals). Its chemistry is largely similar to that of aluminum (diagonal similarity). Thus, this is an element with amphoteric qualities in its compounds, among which the basic ones still predominate.
The electronic configuration of Mg: 1s 2 2s 2 2p 6 3s 2 compared to sodium has one significant feature: the twelfth electron is placed in the 2s orbital, where there is already a 1e - .
Magnesium and calcium ions are irreplaceable elements of the life of any cell. Their ratio in the body must be strictly defined. Magnesium ions are involved in the activity of enzymes (for example, carboxylase), calcium - in the construction of the skeleton and metabolism. Increasing calcium levels improves food absorption. Calcium stimulates and regulates the functioning of the heart. Its excess sharply increases the activity of the heart. Magnesium plays part of the role of a calcium antagonist. The introduction of Mg 2+ ions under the skin causes anesthesia without a period of excitation, paralysis of muscles, nerves and heart. Getting into the wound in the form of metal, it causes long-term non-healing purulent processes. Magnesium oxide in the lungs causes what is called foundry fever. Frequent contact of the skin surface with its compounds leads to dermatitis. The most widely used calcium salts in medicine are CaSO 4 sulfate and CaCL 2 chloride. The first is used for plaster casts, and the second is used for intravenous infusions and as an internal remedy. It helps fight swelling, inflammation, allergies, relieves spasms of the cardiovascular system, and improves blood clotting.
All barium compounds, except BaSO 4, are poisonous. They cause menegoencephalitis with damage to the cerebellum, damage to smooth cardiac muscles, paralysis, and in large doses - degenerative changes in the liver. In small doses, barium compounds stimulate bone marrow activity.
When strontium compounds are introduced into the stomach, stomach upset, paralysis, and vomiting occur; the symptoms of lesions are similar to lesions from barium salts, but strontium salts are less toxic. Of particular concern is the appearance in the body of the radioactive isotope of strontium 90 Sr. It is excreted extremely slowly from the body, and its long half-life and therefore long duration of action can cause radiation sickness.
Radium is dangerous to the body due to its radiation and huge half-life (T 1/2 = 1617 years). Initially, after the discovery and production of radium salts in a more or less pure form, it began to be used quite widely for fluoroscopy, the treatment of tumors and some serious diseases. Now, with the advent of other more accessible and cheaper materials, the use of radium in medicine has practically ceased. In some cases, it is used to produce radon and as an additive to mineral fertilizers.
In the calcium atom, the filling of the 4s orbital is completed. Together with potassium, it forms a pair of s-elements of the fourth period. Calcium hydroxide is a fairly strong base. Calcium, the least active of all alkaline earth metals, has an ionic bond in its compounds.
According to its characteristics, strontium occupies an intermediate position between calcium and barium.
The properties of barium are closest to the properties of alkali metals.
Beryllium and magnesium are widely used in alloys. Beryllium bronzes are elastic alloys of copper with 0.5-3% beryllium; Aviation alloys (density 1.8) contain 85-90% magnesium (“electron”). Beryllium differs from other group IIA metals - it does not react with hydrogen and water, but it dissolves in alkalis because it forms an amphoteric hydroxide:
Be+H 2 O+2NaOH=Na 2 +H 2.
Magnesium reacts actively with nitrogen:
3 Mg + N 2 = Mg 3 N 2.
The table shows the solubility of hydroxides of group II elements.
Traditional technical problem - water hardness, associated with the presence of Mg 2+ and Ca 2+ ions in it. From bicarbonates and sulfates, magnesium and calcium carbonates and calcium sulfate settle on the walls of heating boilers and pipes with hot water. They especially interfere with the operation of laboratory distillers.
S-elements perform an important biological function in a living organism. The table shows their contents.
Extracellular fluid contains 5 times more sodium ions than inside cells. An isotonic solution (“physiological fluid”) contains 0.9% sodium chloride, it is used for injections, washing wounds and eyes, etc. Hypertonic solutions (3-10% sodium chloride) are used as lotions in the treatment of purulent wounds (“pulling " pus). 98% of potassium ions in the body are found inside cells and only 2% in extracellular fluid. A person needs 2.5-5 g of potassium per day. 100 g of dried apricots contain up to 2 g of potassium. 100 g of fried potatoes contains up to 0.5 g of potassium. ATP and ADP participate in intracellular enzymatic reactions in the form of magnesium complexes.
Every day a person needs 300-400 mg of magnesium. It enters the body with bread (90 mg of magnesium per 100 g of bread), cereals (100 g of oatmeal contains up to 115 mg of magnesium), and nuts (up to 230 mg of magnesium per 100 g of nuts). In addition to building bones and teeth based on hydroxylapatite Ca 10 (PO 4) 6 (OH) 2, calcium cations are actively involved in blood clotting, transmission of nerve impulses, and muscle contraction. An adult needs to consume about 1 g of calcium per day. 100 g of hard cheese contains 750 mg of calcium; 100 g of milk – 120 mg of calcium; in 100 g of cabbage – up to 50 mg.
Alkaline earth metals- chemical elements of the 2nd group of the periodic table of elements: calcium, strontium, barium and radium.
- 1 Physical properties
- 2 Chemical properties
- 2.1 Simple substances
- 2.2 Oxides
- 2.3 Hydroxides
- 3 Being in nature
- 4 Biological role
- 5 Notes
Physical properties
Alkaline earth metals include only calcium, strontium, barium and radium, and less commonly magnesium. The first element of this subgroup, beryllium, in most properties is much closer to aluminum than to the higher analogues of the group to which it belongs. The second element in this group, magnesium, differs in some respects significantly from the alkaline earth metals in a number of chemical properties. All alkaline earth metals are gray substances that are solid at room temperature. Unlike alkali metals, they are significantly harder and generally cannot be cut with a knife (the exception is strontium. An increase in the density of alkaline earth metals is observed only starting with calcium. The heaviest is radium, comparable in density to germanium (ρ = 5.5 g/cm3) .
| Atomic number |
Name, symbol |
Number of natural isotopes | Atomic mass | Ionization energy, kJ mol−1 | Electron affinity, kJ mol−1 | EO | Metal. radius, nm | Ionic radius, nm | tpl, °C |
tboiling, °C |
ρ, g/cm³ |
ΔHpl, kJ mol−1 | ΔHboiling, kJ mol−1 |
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| 4 | Beryllium Be | 1+11a | 9,012182 | 898,8 | 0,19 | 1,57 | 0,169 | 0,034 | 1278 | 2970 | 1,848 | 12,21 | 309 |
| 12 | Magnesium Mg | 3+19a | 24,305 | 737,3 | 0,32 | 1,31 | 0,24513 | 0,066 | 650 | 1105 | 1,737 | 9,2 | 131,8 |
| 20 | Calcium Ca | 5+19a | 40,078 | 589,4 | 0,40 | 1,00 | 0,279 | 0,099 | 839 | 1484 | 1,55 | 9,20 | 153,6 |
| 38 | Strontium Sr | 4+35a | 87,62 | 549,0 | 1,51 | 0,95 | 0,304 | 0,112 | 769 | 1384 | 2,54 | 9,2 | 144 |
| 56 | Barium Ba | 7+43a | 137,327 | 502,5 | 13,95 | 0,89 | 0,251 | 0,134 | 729 | 1637 | 3,5 | 7,66 | 142 |
| 88 | Radium Ra | 46a | 226,0254 | 509,3 | - | 0,9 | 0,2574 | 0,143 | 700 | 1737 | 5,5 | 8,5 | 113 |
a Radioactive isotopes
Chemical properties
Alkaline earth metals have the electronic configuration of the outer energy level ns², and are s-elements, along with the alkali metals. Having two valence electrons, alkaline earth metals easily give them up, and in all compounds they have an oxidation state of +2 (very rarely +1).
The chemical activity of alkaline earth metals increases with increasing atomic number. Beryllium in its compact form does not react with oxygen or halogens even at red heat temperatures (up to 600 °C; reactions with oxygen and other chalcogens require an even higher temperature, fluorine is an exception). Magnesium is protected by an oxide film at room temperature and higher temperatures (up to 650 °C) and does not oxidize further. Calcium oxidizes slowly and deeply at room temperature (in the presence of water vapor), and burns with slight heating in oxygen, but is stable in dry air at room temperature. Strontium, barium and radium quickly oxidize in air, giving a mixture of oxides and nitrides, so they, like alkali metals and calcium, are stored under a layer of kerosene.
Also, unlike alkali metals, alkaline earth metals do not form superoxides and ozonides.
Oxides and hydroxides of alkaline earth metals tend to increase their basic properties with increasing atomic number.
Simple substances
Beryllium reacts with weak and strong acid solutions to form salts:
however, it is passivated by cold concentrated nitric acid.
The reaction of beryllium with aqueous solutions of alkalis is accompanied by the release of hydrogen and the formation of hydroxoberyllates:
When carrying out a reaction with an alkali melt at 400-500 °C, dioxoberyllates are formed:
Magnesium, calcium, strontium, barium and radium react with water to form alkalis (except for magnesium, which reacts with water only when hot magnesium powder is added to water):
Also, calcium, strontium, barium and radium react with hydrogen, nitrogen, boron, carbon and other non-metals to form the corresponding binary compounds:
Oxides
Beryllium oxide is an amphoteric oxide, dissolves in concentrated mineral acids and alkalis to form salts:
but with less strong acids and bases the reaction no longer occurs.
Magnesium oxide does not react with dilute and concentrated bases, but reacts easily with acids and water:
Oxides of calcium, strontium, barium and radium are basic oxides that react with water, strong and weak solutions of acids and amphoteric oxides and hydroxides:
Hydroxides
Beryllium hydroxide is amphoteric, in reactions with strong bases it forms beryllates, and with acids - beryllium salts of acids:
Magnesium, calcium, strontium, barium and radium hydroxides are bases, the strength increases from weak to very strong, being the strongest corrosive substance, exceeding potassium hydroxide in activity. They are highly soluble in water (except for magnesium and calcium hydroxides). They are characterized by reactions with acids and acid oxides and with amphoteric oxides and hydroxides:
Being in nature
All alkaline earth metals are found (in varying quantities) in nature. Due to their high chemical activity, all of them are not found in a free state. The most common alkaline earth metal is calcium, the amount of which is 3.38% (by weight of the earth’s crust). It is slightly inferior to magnesium, the amount of which is 2.35% (of the mass of the earth’s crust). Barium and strontium are also common in nature, accounting for 0.05 and 0.034% of the mass of the earth's crust, respectively. Beryllium is a rare element, the amount of which is 6·10−4% of the mass of the earth's crust. As for radium, which is radioactive, it is the rarest of all alkaline earth metals, but it is always found in small quantities in uranium ores. in particular, it can be isolated from there chemically. Its content is 1·10−10% (of the mass of the earth's crust).
Biological role
Magnesium is found in the tissues of animals and plants (chlorophyll), is a cofactor in many enzymatic reactions, is necessary in the synthesis of ATP, is involved in the transmission of nerve impulses, and is actively used in medicine (bischophytotherapy, etc.). Calcium is a common macronutrient in the body of plants, animals and humans. In the human body and other vertebrates, most of it is found in the skeleton and teeth. bones contain calcium in the form of hydroxyapatite. The “skeletons” of most groups of invertebrates (sponges, coral polyps, mollusks, etc.) are made from various forms of calcium carbonate (lime). Calcium ions are involved in blood clotting processes, and also serve as one of the universal second messengers inside cells and regulate a variety of intracellular processes - muscle contraction, exocytosis, including the secretion of hormones and neurotransmitters. Strontium can replace calcium in natural tissues, since it is similar in properties to it. In the human body, the mass of strontium is about 1% of the mass of calcium.
At the moment, nothing is known about the biological role of beryllium, barium and radium. All compounds of barium and beryllium are poisonous. Radium is extremely radiotoxic. In the body, it behaves like calcium - about 80% of radium entering the body accumulates in bone tissue. Large concentrations of radium cause osteoporosis, spontaneous bone fractures and malignant tumors of bones and hematopoietic tissue. Radon, a gaseous radioactive decay product of radium, also poses a danger.
Notes
- According to the new IUPAC classification. According to the outdated classification, they belong to the main subgroup of group II of the periodic table.
- Nomenclature of Inorganic Chemistry. IUPAC Recommendations 2005. - International Union of Pure and Applied Chemistry, 2005. - P. 51.
- Group 2 - Alkaline Earth Metals, Royal Society of Chemistry.
- Gold Fund. School encyclopedia. Chemistry. M.: Bustard, 2003.
| Periodic table of chemical elements by D. I. Mendeleev | ||||||||||||||||||||||||||||||||
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |||||||||||||||
| 1 | H | He | ||||||||||||||||||||||||||||||
| 2 | Li | Be | B | C | N | O | F | Ne | ||||||||||||||||||||||||
| 3 | Na | Mg | Al | Si | P | S | Cl | Ar | ||||||||||||||||||||||||
| 4 | K | Ca | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | ||||||||||||||
| 5 | Rb | Sr | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | ||||||||||||||
| 6 | Cs | Ba | La | Ce | Pr | Nd | Pm | Sm | Eu | Gd | Tb | Dy | Ho | Er | Tm | Yb | Lu | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn |
| 7 | Fr | Ra | Ac | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | MD | No | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Uut | Fl | Uup | Lv | Uus | Uuo |
| 8 | Uue | Ubn | Ubu | Ubb | Ubt | Ubq | UBP | Ubh | ||||||||||||||||||||||||
alkaline earth metals in, alkaline earth metals and, alkaline earth metals chemistry, alkaline earth metals
Properties of alkaline earth metals
Physical properties
Alkaline earth metals (compared to alkali metals) have higher t╟pl. and boiling point, ionization potentials, densities and hardness.
Chemical properties
1. Very reactive.
2. They have a positive valence of +2.
3. React with water at room temperature (except Be) to release hydrogen.
4. They have a high affinity for oxygen (reducing agents).
5. With hydrogen they form salt-like hydrides EH 2.
6. Oxides have the general formula EO. The tendency to form 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
Receipt
Beryllium is obtained by reduction of fluoride:
BeF 2 + Mg═ t ═ Be + MgF 2
Barium is obtained by reduction of the oxide:
3BaO + 2Al═ t ═ 3Ba + Al 2 O 3
The remaining 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; compounds exhibit only the +2 oxidation state. The activity of metals and their reducing ability increases in the series: 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 + O2 = 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 (diluted) = 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 - coloring 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 discovered by an exchange reaction with sulfuric acid or its salts:
Barium sulfate is a white precipitate, insoluble in mineral acids.
Alkaline earth metal oxides
Receipt
1) Oxidation of metals (except 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. Reacts 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 is an amphoteric oxide, soluble in alkalis:
BeO + 2NaOH + H 2 O = Na 2
Alkaline earth metal hydroxides R(OH) 2
Receipt
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 are white crystalline substances, less soluble in water than hydroxides of alkali metals (the solubility of hydroxides decreases with decreasing atomic number; Be(OH) 2 is 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 remaining hydroxides are strong bases (alkalis).
1) Reactions with acid 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) Exchange reactions 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
Water hardness
Natural water containing Ca 2+ and Mg 2+ ions is called hard water. Hard water forms scale when boiled and food products cannot be cooked in it; Detergents do not produce foam.
Carbonate (temporary) hardness is caused by the presence of calcium and magnesium bicarbonates in water, non-carbonate (permanent) hardness is caused by chlorides and sulfates.
The total hardness of water is considered as the sum of carbonate and non-carbonate.
Water hardness is removed by precipitation of Ca 2+ and Mg 2+ ions from solution:
1) boiling:
Сa(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 lime milk:
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 hardness, all four methods are used, and for permanent hardness, only the last two are used.
Thermal decomposition of nitrates.
E(NO3)2 =t= EO + 2NO2 + 1/2O2
Features of the chemistry of 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
Part one. General characteristics II A groups of the Periodic Table of elements.
The following elements are located in this group: Be, Mg, Ca, Sr, Ba, Ra. They have a common electronic configuration: (n-1)p 6 ns 2, except Be 1s 2 2s 2. Due to the latter, the properties of Be are slightly different from the properties of the subgroup as a whole. The properties of magnesium also differ from the properties of the subgroup, but to a lesser extent. In the series Ca – Sr – Ba – Ra, the properties change sequentially. The relative electronegativity in the Be – Ra series decreases because As the size of the atom increases, valence electrons are given up more readily. The properties of subgroup IIA elements are determined by the ease of losing two ns electrons. In this case, E 2+ ions are formed. When studying X-ray diffraction, it turned out that in some compounds elements of the IIA subgroup exhibit monovalency. An example of such compounds are EG, which are obtained by adding E to the EG 2 melt. All elements of this series are not found in nature in a free state due to their high activity.
Part two. Beryllium and magnesium.
History of beryllium
Beryllium compounds in the form of precious stones have been known since ancient times. For a long time, people have been searching for and developing deposits of blue aquamarines, green emeralds, greenish-yellow beryls and golden chrysoberyls. But it was only at the end of the 18th century that chemists suspected that beryl contained some new unknown element. In 1798, the French chemist Lewis Nicolas Vauquelin isolated the oxide "La terree du beril" from beryl, which was different from aluminum oxide. This oxide gave the salts a sweet taste, did not form alum, was dissolved in a solution of ammonium carbonate, and was not precipitated by potassium oxalate. Beryllium metal was first obtained in 1829 by the famous German scientist Weller and at the same time by the French scientist Bussy, who obtained beryllium metal powder by reducing beryllium chloride with potassium metal. The beginning of industrial production dates back to the 30-40s. last century.
History of magnesium
The element got its name from the area of Magnesia in Ancient Greece. Natural magnesium-containing materials magnesite and dolomite have long been used in construction.
The first attempts to isolate the metallic base of magnesia in its pure form were made at the beginning of the 19th century. the famous English physicist and chemist Humphry Davy (1778–1829) after he subjected molten potassium and caustic soda to electrolysis and obtained metallic Na and K. He decided to try in a similar way to carry out the decomposition of the oxides of alkaline earth metals and magnesia. In his initial experiments, Davy passed a current through wet oxides, protecting them from contact with air by a layer of oil; however, the metals were fused with the cathode and could not be separated.
Davy tried many different methods, but all of them were unsuccessful for various reasons. Finally, in 1808, he was lucky - he mixed wet magnesia with mercury oxide, placed the mass on a platinum plate and passed a current through it; The amalgam was transferred into a glass tube, heated to remove the mercury, and a new metal was obtained. Using the same method, Davy managed to obtain barium, calcium and strontium. Industrial production of magnesium using the electrolytic method began in Germany at the end of the 19th century. Theoretical and experimental work on the production of magnesium by the electrolytic method in our country was carried out by P.P. Fedotiev; the process of reduction of magnesium oxide by silicon in vacuum was studied by P.F. Antipin.
Spreading
Beryllium is one of the less common elements: its content in the earth’s crust is 0.0004 wt. %. Beryllium in nature is in a bound state. The most important minerals of beryllium are beryl-Be 3 Al 2 (SiO 3) 6, chrysoberyl-Be(AlO 2) 2 and phenacite-Be 2 SiO 4. The main part of beryllium is dispersed as impurities into the minerals of a number of other elements, especially aluminum. Beryllium is also found in deep sea sediments and the ash of some coals. Some varieties of beryl, colored by impurities in different colors, are classified as precious stones. These are, for example, green emeralds and bluish-green aquamarines.
Magnesium is one of the most common elements in the earth's crust. The magnesium content is 1.4%. The most important minerals include, in particular, carbonaceous carbonate rocks, which form huge massifs on land and even entire mountain ranges - magnesite MgCO 3 and dolomite MgCO 3 -CaCO 3 . Under layers of various alluvial rocks, together with deposits of rock salt, colossal deposits of another easily soluble magnesium-containing mineral are known - carnallite MgCl 2 -KCl-6H 2 O. In addition, in many minerals magnesium is closely associated with silica, forming, for example, olivine[(Mg, Fe) 2 SiO 4 ] and less common forsterite(Mg 2 SiO 4). Other magnesium containing minerals are brucite Mg(OH)2 , kieserite MgSO4 , epsonite MgSO 4 -7H 2 O , Cainite MgSO 4 -KCl-3H 2 O . On the surface of the Earth, magnesium easily forms aqueous silicates (talc, asbestos, etc.), an example of which is serpentine 3MgO-2SiO 2 -2H 2 O. Of the known minerals, about 13% contain magnesium. However, natural magnesium compounds are widely found in dissolved form. In addition to various minerals and rocks, 0.13% magnesium in the form of MgCl 2 is constantly found in ocean waters (its reserves here are inexhaustible - about 6-10 16 tons) and in salt lakes and springs. Magnesium is also included in chlorophyll in an amount of up to 2% and acts here as a complexing agent. The total content of this element in the living matter of the Earth is estimated at about 10 11 tons.
Receipt
The main (about 70%) method of producing magnesium is electrolysis of molten carnallite or MgCl 2 under a layer of flux to protect against oxidation. The thermal method of obtaining magnesium (about 30%) involves the reduction of calcined magnesite or dolomite. Beryllium concentrates are processed into beryllium oxide or hydroxide, from which fluoride or chloride is obtained. When obtaining metallic beryllium, electrolysis of the melt of BeCl 2 (50 wt.%) and NaCl is carried out. This mixture has a melting point of 300 o C versus 400 o C for pure BeCl 2. Beryllium is also obtained magnesium- or aluminothermally at 1000-1200 0 C from Na 2: Na 2 + 2Mg = Be + 2Na + MgF 2. Particularly pure beryllium (mainly for the nuclear industry) is obtained by zone smelting, vacuum distillation and electrolytic refining.
Peculiarities
Beryllium is a “pure” element. In nature, magnesium occurs in the form of three stable isotopes: 24 Mg (78.60%), 25 Mg (10.11%) and 26 Mg (11.29%). Isotopes with masses 23, 27 and 28 were artificially obtained.
Beryllium has an atomic number of 4 and an atomic weight of 9.0122. It is in the second period of the periodic table and heads the main subgroup of group 2. The electronic structure of the beryllium atom is 1s 2 2s 2. During a chemical interaction, the beryllium atom is excited (which requires a cost of 63 kcal/g×atom) and one of the 2s electrons moves to the 2p orbital, which determines the specifics of the chemistry of beryllium: it can exhibit a maximum covalency of 4, forming 2 bonds according to the exchange mechanism, and 2 for donor-acceptor. On the ionization potential curve, beryllium occupies one of the top places. The latter corresponds to its small radius and characterizes beryllium as an element that is not particularly willing to give up its electrons, which primarily determines the low degree of chemical activity of the element. From the point of view of electronegativity, beryllium can be considered as a typical transition element between electropositive metal atoms, which easily donate their electrons, and typical complexing agents, which tend to form covalent bonds. Beryllium exhibits a diagonal analogy with aluminum to a greater extent than LicMg and is a kinosymmetric element. Beryllium and its compounds are highly toxic. MPC in the air is 2 μg/m 3 .
In the periodic table of elements, magnesium is located in the main subgroup of group II; Magnesium serial number is 12, atomic weight is 24.312. The electronic configuration of an unexcited atom is 1s 2 2s 2 2p 6 3s 2 ; the structure of the outer electron shells of the Mg (3s 2) atom corresponds to its zero-valent state. Excitation to divalent 3s 1 3p 1 requires an expenditure of 62 kcal/g-atom. The ionization potentials of magnesium are lower than those of beryllium, therefore magnesium compounds are characterized by a higher proportion of bond ionicity. In terms of complexing ability, magnesium is also inferior to beryllium. Interaction with elements of group IIIB with unfinished d-shells has some features. This group includes Sc, Y, Ln, and Th. These elements form a number of intermediate phases with magnesium and dissolve well in it in the liquid state. State diagrams of mixtures of these elements with magnesium are eutectic in nature. The solubility of these elements in magnesium in the solid state is not high (2 – 5% by weight). With alkaline earth and especially with alkali metals, magnesium does not form a significant solubility region in the solid state, which is due to the large difference in atomic radii. The exception is lithium, whose atomic radius differs from the atomic radius of magnesium by 2%. Systems of magnesium with copper, silver and gold are of the eutectic type. The solubility of silver at the eutectic temperature is –16% by weight.
Physical properties
Beryllium - silver-white metal. Quite hard and fragile. Has diamagnetic properties. In air, it becomes covered with a thin oxide film, giving the metal a gray, matte color and protecting it from further corrosion. The compressibility of beryllium is very low. The least amount of metals (17 times less than Al) blocks X-ray radiation. It crystallizes in an hcp structure with periods a=0.228 nm and c=0.358 nm, CN=6. At 1254 o C, the hexagonal a modification transforms into cubic b. Beryllium forms eutectic alloys with Al and Si.
