Alkaline earth metals in the periodic table. Alkaline earth metals (9th grade). Special properties of beryllium and magnesium

The concept of alkaline earth metals includes a part 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 the alkaline earth classification, therefore, in some sources, beryllium and magnesium are not included in the list, limited to four elements.

The metal got its name due to the fact that when their oxides interact with water, an alkaline environment is formed. Physical properties of alkaline earth metals: all elements have a gray metallic color, under normal conditions they have a solid structure, their density increases with increasing serial number, and they have a very high melting point. Unlike alkali metals, the 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 atomic number, they act as a reducing agent in reactions.

The characteristic of alkaline earth metals indicates their high activity. In particular, this applies to elements with a large serial number. For example, under normal conditions, beryllium does not interact with oxygen and halogens. To start the response mechanism, it must be heated to a temperature of over 600 degrees Celsius. Magnesium under normal conditions has an oxide film on the surface and also does not react with oxygen. Calcium is oxidized, but rather slowly. But strontium, barium and radium oxidize 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 ordinal number of the metal. Beryllium hydroxide is amphoteric compound, which does not react with water, but is highly soluble in acids. Magnesium hydroxide is a weak alkali, insoluble in water but reactive with strong acids. Calcium hydroxide is a strong, water-soluble base that 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, which reacts well with water and almost all types of acids.

How to get

Alkaline earth metal hydroxides are obtained by exposing a pure element to water. The reaction proceeds at room conditions (except for beryllium, which requires an increase in temperature) with the release of hydrogen. When heated, all alkaline earth metals react with halogens. The resulting compounds are used in the production of a wide range of products from chemical fertilizers to ultra-precise microprocessor parts. Alkaline earth metal compounds exhibit the same high activity as pure elements, so they are used in many chemical reactions.

Most often this occurs during exchange reactions, when it is necessary to displace less active metal. They take part in redox reactions as a strong reducing agent. Divalent calcium and magnesium cations give water the so-called hardness. This phenomenon is overcome by the precipitation of ions with the help of physical impact or adding special emollients to the water. Alkaline earth metal salts 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 found in their pure form, as they quickly interact with the environment, forming chemical compounds. They are part of the minerals and rocks contained in the thickness 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 has 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 mined by reducing it from fluoride at high temperatures. Barium recovers from its oxide. Calcium, magnesium and strontium are obtained by electrolysis of their chloride melt. The most difficult thing is to synthesize pure radium. It is mined by impact on uranium ore. According to scientists, on average, one ton of ore accounts for 3 grams of pure radium, although there are also rich deposits that contain as much as 25 grams per ton. The methods of precipitation, fractional crystallization and ion exchange are used to isolate the metal.

Application of alkaline earth metals

The range of application of alkaline earth metals is very extensive and covers many industries. Beryllium is in most cases used as an alloying additive in various alloys. It increases the hardness and strength of materials, well protects the surface from corrosion. Also due to weak absorption 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 the production of titanium. Its alloys are distinguished by high strength and lightness, therefore they are used in the manufacture of aircraft, cars, 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 process of vital activity of the body, therefore it is part of some drugs.

Calcium in its pure form is practically not used. It is needed to restore other metals from their compounds, as well as in the production of drugs to strengthen bone tissue. Strontium is used to reduce 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 aggressive environments, as it has excellent protective properties. Radium is used in medicine for short-term skin irradiation in the treatment of malignant tumors.

Consider the chemical properties of alkaline earth metals. Let's define the features of their structure, obtaining, being in nature, application.

Regulations in PS

To begin with, we determine the location of these elements in Mendeleev. They are located in the second group of the main subgroup. These include calcium, strontium, radium, barium, magnesium, beryllium. They all contain two valence electrons. In general, beryllium, magnesium and alkaline earth metals have ns2 electrons at the outer level. AT chemical compounds they exhibit an oxidation state of +2. During interaction with other substances, they exhibit reducing properties, donating electrons from an external energy level.

Change properties

As the nucleus of the atom increases, beryllium, magnesium and enhance their metallic properties, since an increase in the radius of their atoms is observed. Consider physical properties alkaline earth metals. Beryllium in its normal state is a gray metal with a steely sheen. It has a dense hexagonal crystal lattice. Upon contact with atmospheric oxygen, beryllium immediately forms an oxide film, as a result of which its chemical activity decreases, and a dull coating is formed.

Physical properties

magnesium as a simple substance is a white metal that forms an oxide coating in air. It has a hexagonal crystal lattice.

The physical properties of the alkaline earth metals calcium, barium, strontium are similar. They are metals with a characteristic silvery sheen, covered with a yellowish film under the influence of atmospheric oxygen. Calcium and strontium have a cubic face-centered lattice, while barium has a body-centered structure.

The chemistry of alkaline earth metals is based on the fact that they have a metallic bond. That is why they are distinguished by high electrical and thermal conductivity. Their melting and boiling points are greater than those of alkali metals.

How to get

The production of beryllium on an industrial scale is carried out by reducing the metal from fluoride. The condition for the flow of this chemical reaction is preheating.

Given that alkaline earth metals are found in nature in the form of compounds, electrolysis of their salt melts is carried out to obtain magnesium, strontium, and calcium.

Chemical properties

The chemical properties of alkaline earth metals are associated with the need to first remove an oxide film layer from their surface. It is she who determines the inertness of these metals to water. Calcium, barium, strontium, when dissolved in water, form hydroxides with pronounced basic properties.

The chemical properties of alkaline earth metals suggest their interaction with oxygen. For barium, the interaction product is peroxide, for all the others, oxides are formed after the reaction. In all representatives of this class, oxides exhibit basic properties, only beryllium oxide is characterized by amphoteric properties.

The chemical properties of alkaline earth metals are also manifested in the reaction with sulfur, halogens, and nitrogen. When reacting with acids, the dissolution of these elements is observed. Considering that beryllium belongs to amphoteric elements, it is able to enter into chemical interaction with alkali solutions.

Qualitative reactions

Basic formulas of alkaline earth metals covered in the course inorganic chemistry associated with salts. To identify representatives of this class in a mixture with other elements, you can use a qualitative definition. When salts of alkaline earth metals are introduced into the flame of an alcohol lamp, the coloring of the flame with cations is observed. The strontium cation gives a dark red tint, the calcium cation - Orange color, and the barium cation is green.

To detect the barium cation in a qualitative analysis, sulfate anions are used. As a result of this reaction, barium sulfate is formed. white color, which is insoluble in inorganic acids.

Radium is a radioactive element found in nature in trace amounts. When magnesium interacts with oxygen, a blinding flash is observed. This process has been used for some time when photographing in dark rooms. Magnesium flares have now been replaced by electrical systems. Beryllium belongs to the alkaline earth metal family and reacts with many chemicals. Calcium and magnesium, like aluminum, can reduce such rare metals as titanium, tungsten, molybdenum, niobium. The data are called calciumthermia and magnesiumthermia.

Application features

What is the use of alkaline earth metals? Calcium and magnesium are used to make light alloys and rare metals.

For example, magnesium is contained in the composition of duralumin, and calcium is a component of lead alloys used to obtain cable sheaths and create bearings. Widespread use of alkaline earth metals in technology in the form of oxides. (calcium oxide) and burnt magnesium (magnesium oxide) are required for the construction industry.

When calcium oxide interacts with water, a significant amount of heat is released. (calcium hydroxide) is used for construction. A white suspension of this substance (milk of lime) is used in the sugar industry for the purification of beet juice.

Salts of metals of the second group

Salts of magnesium, beryllium, alkaline earth metals can be obtained by reacting their oxides with acids. Chlorides, fluorides, iodides of these elements are white crystalline substances, mostly highly soluble in water. Among sulfates, only compounds of magnesium and beryllium have solubility. Its decrease from beryllium salts to barium sulfates is observed. Carbonates are practically insoluble in water or have minimal solubility.

Sulfides of alkaline earth elements are found in small amounts in heavy metals. If you direct light at them, you can get different colors. Sulfides are included in the luminous compounds called phosphors. Apply similar paints to create luminous dials, road signs.

Common alkaline earth metal compounds

Calcium carbonate is the most common earth's surface element. It is an integral part of compounds such as limestone, marble, chalk. Among them, the main application is limestone. This mineral is indispensable in construction, it is considered an excellent building stone. In addition, quicklime and slaked lime, glass, and cement are obtained from this inorganic compound.

The use of crushed lime helps to strengthen roads, and thanks to the powder, soil acidity can be reduced. represents the shells of the most ancient animals. This compound is used to make rubber, paper, and create school crayons.

Marble is in demand among architects and sculptors. It was from marble that many were created. unique creations Michelangelo. Some of the Moscow metro stations are lined with marble tiles. Magnesium carbonate is used in large volumes in the manufacture of bricks, cement, glass. It is needed in the metallurgical industry to remove waste rock.

Calcium sulfate, found in nature in the form of gypsum (calcium sulfate crystalline hydrate), is used in the construction industry. In medicine, this compound is used to make casts, as well as to create plaster casts.

Alabaster (semiaquatic gypsum) releases a huge amount of heat when interacting with water. This is also applied in industry.

Epsom salt (magnesium sulfate) is used medicinally as a laxative. This substance has a bitter taste, it is found in sea water.

Barite porridge (barium sulfate) does not dissolve in water. That is why this salt is used in X-ray diagnostics. Salt delays X-rays which allows to detect diseases of the gastrointestinal tract.

Phosphorites (rock) and apatites contain calcium phosphate. They are needed to obtain calcium compounds: oxides, hydroxides.

Calcium plays for living organisms special meaning. It is this metal that is necessary for the construction of the bone skeleton. Calcium ions are necessary to regulate the work of the heart, increase blood clotting. Its deficiency causes malfunctions nervous system, loss of clotting, loss of the ability of hands to hold various objects normally.

In order to avoid health problems, every day a person should consume approximately 1.5 grams of calcium. The main problem is that in order for the body to absorb 0.06 grams of calcium, you need to eat 1 gram of fat. Maximum amount This metal is found in lettuce, parsley, cottage cheese, and cheese.

Conclusion

All representatives of the second group of the main subgroup of the periodic table are necessary for life and activity modern man. For example, magnesium is a stimulant of metabolic processes in the body. He must be present in nervous tissue, blood, bones, liver. Magnesium is an active participant in photosynthesis in plants, as it is an integral part of chlorophyll. Human bones make up about a fifth of the total weight. They contain calcium and magnesium. Oxides, salts of alkaline earth metals have found a variety of applications in the construction industry, pharmaceuticals and medicine.

The fresh surface of E quickly darkens due to the formation of an oxide film. This film is relatively dense - over time, the entire metal slowly oxidizes. The film consists of EO, as well as EO 2 and E 3 N 2 . The normal electrode potentials of the reactions E-2e = E 2+ 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 (200-300 o C) calcium interacts with water vapor according to the scheme:

2Ca + H 2 O \u003d CaO + CaH 2.

The secondary reactions are:

CaH 2 + 2H 2 O \u003d Ca (OH) 2 + 2H 2 and CaO + H 2 O \u003d Ca (OH) 2.

In strong sulfuric acid, E almost do not dissolve due to the formation of a film of poorly soluble ESO 4 . With dilute mineral acids, E react violently with the release of hydrogen. When heated above 800 ° C, calcium reacts with methane according to the scheme:

3Ca + CH 4 \u003d CaH 2 + CaC 2.

When heated, they react with hydrogen, with sulfur and with 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 those of its counterparts. All radium compounds slowly decompose under the action of their own radiation, while 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 higher than the temperature environment by 1.5 deg. It is also known that 1 g of radium per day emits 1 mm 3 of radon (226 Ra = 222 Rn + 4 He), on which its use as a source of radon for radon baths is based.

hydrides E - white, crystalline salt-like substances. They are obtained directly from the elements by heating. The start temperatures of the reaction E + H 2 = EN 2 are 250 o C (Ca), 200 o C (Sr), 150 o C (Ba). Thermal dissociation of EN 2 begins at 600 o C. CaH 2 does not decompose in a hydrogen atmosphere at the melting point (816 o C). In the absence of moisture, alkaline earth metal hydrides are stable in air at ordinary temperatures. They do not react with halogens. However, when heated, the chemical activity of EN 2 increases. They are able to reduce oxides to metals (W, Nb, Ti, Ce, Zr, Ta), for example

2CaH 2 + TiO 2 \u003d 2CaO + 2H 2 + Ti.

The reaction of CaH 2 with Al 2 O 3 takes place at 750 o C:

3CaH 2 + Al 2 O 3 \u003d 3CaO + 3H 2 + 2Al,

CaH 2 + 2Al \u003d CaAl 2 + H 2.

CaH2 reacts with nitrogen at 600°C according to the scheme:

3CaH 2 + N 2 \u003d Ca 3 N 2 + 3H 2.

When EN 2 is ignited, they slowly burn out:

EN 2 + O 2 \u003d H 2 O + CaO.

Explosive when mixed with solid oxidizers. Under the action of water on EN 2, hydroxide and hydrogen are released. This reaction is highly exothermic: EN 2 wetted with water in air ignites spontaneously. EN 2 reacts with acids, for example, according to the scheme:

2HCl + CaH 2 \u003d CaCl 2 + 2H 2.

EN 2 is used to obtain pure hydrogen, as well as to determine traces of water in organic solvents. Nitride E are colorless refractory substances. They are obtained directly from the elements at elevated temperatures. They decompose in water according to the scheme:

E 3 N 2 + 6H 2 O \u003d 3E (OH) 2 + 2NH 3.

E 3 N 2 react when heated with CO according to the scheme:

E 3 N 2 + 3CO \u003d 3EO + N 2 + 3C.

The processes that occur when E 3 N 2 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. Phosphides E 3 R 2 are formed directly from the elements or by calcining trisubstituted phosphates with coal:

Ca 3 (RO 4) 2 + 4C \u003d Ca 3 P 2 + 4CO

They are hydrolyzed by water according to the scheme:

E 3 R 2 + 6H 2 O \u003d 2RN 3 + 3E (OH) 2.

With acids, alkaline earth metal phosphides give the corresponding salt and phosphine. This is the basis for their use for the production of phosphine in the laboratory.

Complex ammonia composition E (NH 3) 6 - solids with a metallic luster and high electrical conductivity. They are obtained by the action of liquid ammonia on E. They ignite spontaneously in air. Without air access, they decompose into the corresponding amides: E (NH 3) 6 \u003d E (NH 2) 2 + 4NH 3 + H 2. When heated, they vigorously decompose according to the same pattern.

Carbides alkaline earth metals, which are obtained by calcining E with coal, are decomposed by water with the release of acetylene:

ES 2 + 2H 2 O \u003d E (OH) 2 + C 2 H 2.

The reaction with BaC 2 is so violent that it ignites on contact with water. The heats of formation of ES 2 from the elements for Ca and Ba are 14 and 12 kcalmol. When heated with nitrogen, ES 2 give CaCN 2 , Ba(CN) 2 , SrCN 2 . known silicides (ESi and ESi 2). They can be obtained by heating directly from the elements. They hydrolyze with water and react with acids to give H 2 Si 2 O 5 , SiH 4 , the corresponding E compound, and hydrogen. known borides EV 6 obtained from the elements when heated.

Oxides calcium and its analogues are white refractory (T bp CaO = 2850 o C) substances that actively 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 + 2NH 4 Cl \u003d SrCl 2 + 2NH 3 + H 2 O.

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 o C reacts with potassium cyanide:

KCN + SrO = Sr + KCNO.

Strontium oxide dissolves in methanol to form Sr(OCH 3) 2 . During magnesium-thermal reduction of BaO, an intermediate oxide Ba 2 O can be obtained, which is unstable and disproportionate.

Hydroxides alkaline earth metals - white substances soluble in water. They are strong bases. In the Ca-Sr-Ba series, the basic nature and solubility of hydroxides increase. rPR(Ca(OH) 2) = 5.26, rPR(Sr(OH) 2) = 3.5, rPR(Ba(OH) 2) = 2.3. Ba(OH) 2 is usually isolated from hydroxide solutions. 8H 2 O, Sr (OH) 2. 8H 2 O, Ca (OH) 2. H 2 O. EOs add water to form hydroxides. The use of CaO in construction is based on this. A close mixture of Ca(OH) 2 and NaOH in a 2:1 weight ratio is called soda lime and is widely used as a CO 2 scavenger. Ca (OH) 2, when standing in air, absorbs CO 2 according to the scheme:

Ca(OH)2 + CO2 = CaCO3 + H2O.

About 400 o C Ca (OH) 2 reacts with carbon monoxide:

CO + Ca (OH) 2 \u003d CaCO 3 + H 2.

Barite water reacts with CS 2 at 100 o C:

CS 2 + 2Ba (OH) 2 \u003d BaCO 3 + Ba (HS) 2 + H 2 O.

Aluminum reacts with barite water:

2Al + Ba (OH) 2 + 10H 2 O \u003d Ba 2 + 3H 2. E(OH) 2

used to open carbonic anhydride.

E form peroxides white. They are significantly less stable than oxides and are strong oxidizers. Of practical importance is the most stable BaO 2, which is a white, paramagnetic powder with a density of 4.96 g1cm 3 etc. pl. 450°. BaO 2 is stable at normal temperature (it can be stored for years), it is poorly soluble in water, alcohol and ether, it dissolves in dilute acids with the release of salt and hydrogen peroxide. Thermal decomposition barium peroxides accelerate oxides, Cr 2 O 3 , Fe 2 O 3 and CuO. Barium peroxide reacts when heated with hydrogen, sulfur, carbon, ammonia, ammonium salts, potassium ferricyanide, etc. With concentrated hydrochloric acid barium peroxide reacts releasing chlorine:

BaO 2 + 4HCl = BaCl 2 + Cl 2 + 2H 2 O.

It oxidizes water to hydrogen peroxide:

H 2 O + BaO 2 \u003d Ba (OH) 2 + H 2 O 2.

This reaction is reversible and in the presence of even carbonic acid the equilibrium is shifted to the right. ВаО 2 is used as a starting product for the production of Н 2 О 2 and also as an oxidizing agent in pyrotechnic compositions. However, BaO 2 can also act as a reducing agent:

HgCl 2 + BaO 2 \u003d Hg + BaCl 2 + O 2.

BaO 2 is obtained by heating BaO in air flow to 500 ° C according to the scheme:

2ВаО + О 2 = 2ВаО 2.

As the temperature rises, the reverse process takes place. Therefore, when Ba burns, only oxide is released. SrO 2 and CaO 2 are less stable. common method obtaining EO 2 is the interaction of E(OH) 2 with H 2 O 2, while EO 2 is released. 8H 2 O. Thermal decomposition of EO 2 begins at 380 o C (Ca), 480 o C (Sr), 790 o C (Ba). When EO 2 is heated with concentrated hydrogen peroxide, yellow unstable substances, EO 4 superoxides, can be obtained.

E salts are usually colorless. Chlorides, bromides, iodides and nitrates are highly soluble in water. Fluorides, sulfates, carbonates and phosphates are poorly soluble. Ion Ba 2+ - toxic. Halids E are divided into two groups: fluorides and all the rest. Fluorides are almost insoluble in water and acids and do not form crystalline hydrates. On the contrary, chlorides, bromides, and iodides are highly soluble in water and are isolated from solutions in the form of crystalline hydrates. Some properties of EG 2 are presented below:

When obtained by exchange decomposition in solution, fluorides are released in the form of voluminous mucous precipitates, which quite easily form colloidal solutions. EG 2 can be obtained by acting with the corresponding halogens on the corresponding E. EG 2 melts are capable of dissolving up to 30% E. When studying the electrical conductivity of chloride melts of elements of the second group of the main subgroup, it was found that their molecular-ionic composition is very different. The degrees of dissociation according to the scheme ESl 2 = E 2+ + 2Cl- are equal: BeCl 2 - 0.009%, MgCl 2 - 14.6%, CaCl 2 - 43.3%, SrCl 2 - 60.6%, BaCl 2 - 80, 2%. Halides (except fluorides) E contain water of crystallization: CaCl 2 . 6H 2 O, SrCl 2. 6H 2 O and BaCl 2. 2H 2 O. X-ray diffraction analysis established the structure of E[(OH 2) 6 ]G 2 for Ca and Sr crystalline hydrates. With slow heating of EG 2 crystalline hydrates, anhydrous salts can be obtained. CaCl 2 readily forms supersaturated solutions. Natural CaF 2 (fluorite) is used in the ceramic industry, and is also used to produce HF and is a fluorine mineral. Anhydrous CaCl 2 is used as a desiccant due to its hygroscopicity. Calcium chloride hydrate is used for the preparation of refrigeration mixtures. BaCl 2 - used in cx and for opening

SO 4 2- (Ba 2+ + SO 4 2- \u003d BaSO 4).

Fusion of EG2 and EN2 hydrohalides can be obtained:

EG 2 + EN 2 = 2ENG.

These substances melt without decomposition but are hydrolyzed by water:

2ENG + 2H 2 O \u003d EG 2 + 2H 2 + E (OH) 2.

Solubility in water chlorates , bromates and iodates in water it decreases in the series Ca - Sr - Ba and Cl - Br - I. Ba (ClO 3) 2 - is used in pyrotechnics. Perchlorates E are highly soluble not only in water but also in organic solvents. The most important of the E(ClO 4) 2 is Ba(ClO 4) 2 . 3H 2 O. Anhydrous barium perchlorate is a good drying agent. Its thermal decomposition begins only at 400 o C. Hypochlorite calcium Ca (ClO) 2. nH 2 O (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. bleach can be obtained by acting with chlorine on solid slaked lime. It decomposes with water and smells like chlorine in the presence of moisture. Reacts with CO 2 of air:

CO 2 + 2CaOCl 2 \u003d CaCO 3 + CaCl 2 + Cl 2 O.

Chlorine lime is used as an oxidizing agent, bleach and as a disinfectant.

For alkaline earth metals, azides E(N 3) 2 and thiocyanates E(CNS) 2 . 3H 2 O. Azides are much less explosive than lead azide. The thiocyanates easily lose water when heated. They are highly soluble in water and organic solvents. Ba(N 3) 2 and Ba(CNS) 2 can be used to obtain azides and thiocyanates of other metals from sulfates by an exchange reaction.

Nitrates calcium and strontium usually exist in the form of Ca(NO 3) 2 crystalline hydrates. 4H 2 O and Sr(NO 3) 2 . 4H 2 O. For barium nitrate, the formation of a crystalline hydrate is not characteristic. When heated, Ca (NO 3) 2. 4H 2 O and Sr(NO 3) 2 . 4H 2 O easily lose water. In an inert atmosphere, nitrates E are thermally stable up to 455 o C (Ca), 480 o C (Sr), 495 o C (Ba). The hydrated melt of calcium nitrate 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 exhibits a tendency to complex formation, for which an unstable complex K 2 is known. Calcium nitrate is soluble in alcohols, methyl acetate, acetone. Strontium and barium nitrates are almost insoluble there. The melting points of nitrates E are estimated at 600 o C, however, at the same temperature, decomposition begins:

E (NO 3) 2 \u003d E (NO 2) 2 + O 2.

Further decomposition occurs at a higher temperature:

E (NO 2) 2 \u003d EO + NO 2 + NO.

E nitrates have long been used in pyrotechnics. Highly volatile salts E color the flame in the appropriate colors: Ca - orange-yellow, Sr - red-carmine, Ba - yellow-green. Let's understand the essence of this using the example of Sr: Sr 2+ has two HAOs: 5s and 5p or 5s and 4d. We will inform the energy of this system - we will heat it. Electrons from orbitals closer to the nucleus will move to these HAOs. But such a system is not stable and will release energy in the form of a quantum of light. Just Sr 2+ emits quanta with a frequency corresponding to the lengths of red waves. When obtaining 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 oxidizer, a solid reducing agent, and some organic matter, bleaching the flame of the reducing agent, and being a binding agent. Calcium nitrate is used as a fertilizer.

All phosphates and hydrophosphates E are poorly soluble in water. They can be obtained by dissolving an appropriate amount of CaO or CaCO 3 in phosphoric acid. They are also precipitated during exchange reactions such as:

(3-x) Ca 2+ + 2H x PO 4 - (3-x) \u003d Ca (3-x) (H x PO 4) 2.

Of practical importance (as a fertilizer) is monosubstituted calcium orthophosphate, which, along with Ca (SO 4), is part of superphosphate. It is received according to the scheme:

Ca 3 (PO 4) 2 + 2H 2 SO 4 \u003d Ca (H 2 PO 4) 2 + 2CaSO 4

Oxalates 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:

CaC 2 O 4 \u003d CaCO 3 + CO.

Acetates E are isolated in the form of crystalline hydrates and are highly soluble in water.

With sulfates E - white, poorly soluble substances in water. Solubility CaSO 4 . 2H 2 O per 1000 g of water at normal temperature is 8. 10 -3 mol, SrSO 4 - 5. 10 -4 mol, BaSO 4 - 1. 10 -5 mol, RaSO 4 - 6. 10 -6 mol. In the Ca - Ra series, the solubility of sulfates decreases rapidly. Ba 2+ is a sulfate ion reagent. Calcium sulfate contains water of crystallization. Above 66 ° C, anhydrous calcium sulfate is released from the solution, below - CaSO 4 gypsum. 2H 2 O. Heating of 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 as early as 2000 years ago. The solubility of ESO 4 in strong sulfuric acid is much higher than in water (BaSO 4 up to 10%), which indicates complex formation. The corresponding complexes are ESO 4 . H 2 SO 4 can be obtained in the free state. Double salts with alkali metal and ammonium sulfates are known only for Ca and Sr. (NH 4) 2 is soluble in water and is used in analytical chemistry to separate Ca from Sr, because (NH 4) 2 is slightly soluble. Gypsum is used for the combined production of sulfuric acid and cement, because. when heated with a reducing agent (charcoal), gypsum decomposes:

CaSO 4 + C \u003d CaO + SO 2 + CO.

At a higher temperature (900 o C), sulfur is reduced even more according to the scheme:

CaSO 4 + 3C \u003d CaS + CO 2 + 2CO.

A similar decomposition of Sr and Ba sulfates begins at higher temperatures. BaSO 4 is non-toxic and is used in medicine and in the production of mineral paints.

Sulfides E are white solids that crystallize as NaCl. The heats of their formation and the energies of the crystal lattices are (kcalmol): 110 and 722 (Ca), 108 and 687 (Sr), 106 and 656 (Ba). They can be obtained by synthesis from elements when heated or by calcining sulfates with coal:

ES04 + 3C = ES + CO2 + 2CO.

Less soluble CaS (0.2 hl). ES enters into the following reactions when heated:

ES + H 2 O \u003d EO + H 2 S; ES + G 2 \u003d S + EG 2; ES + 2O 2 \u003d ESO 4; ES + xS = ES x+1 (x=2.3).

Sulfides of alkaline earth metals in a neutral solution are completely hydrolyzed according to the scheme:

2ES + 2H 2 O \u003d E (HS) 2 + E (OH) 2.

Acid sulfides can also be obtained in a free state by evaporating a solution of sulfides. They react with sulfur:

E (NS) 2 + xS \u003d ES x + 1 + H 2 S (x \u003d 2.3.4).

Of the crystalline hydrates, BaS are known. 6H 2 O and Ca(HS) 2 . 6H 2 O, Ba (HS) 2. 4H 2 O. Ca(HS) 2 is used to remove hair. ES are subject to the phenomenon of phosphorescence. known polysulfides E: ES 2, ES 3, ES 4, ES 5. They are obtained by boiling a suspension of ES in water with sulfur. In air, ES are oxidized: 2ES + 3O 2 \u003d 2ESO 3. By passing air through the CaS suspension, one can obtain thiosulfate Sa according to the scheme:

2CaS + 2O 2 + H 2 O \u003d Ca (OH) 2 + CaS 2 O 3

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 subject to hydrolysis, but to a lesser extent than sulfides.

Solubility chromates E in the series Ca - Ba falls 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:

E 2+ + CrO 4 2- = ECrO4.

Calcium chromate is released in the form of a crystalline hydrate - CaCrO 4 . 2H 2 O (rPR CaCrO 4 = 3.15). Even before the melting point, it loses water. SrCrO 4 and ВаCrO 4 do not form crystalline hydrates. pPR SrCrO 4 = 4.44, pPR BaCrO 4 = 9.93.

Carbonates E white, poorly soluble substances in water. When heated, ESO 3 pass into EO, splitting off CO 2 . In the Ca-Ba series, the thermal stability of carbonates increases. The most practically important of them 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 production of lime from limestone amounts to tens of millions of tons. Thermal dissociation of CaCO 3 is endothermic:

CaCO 3 \u003d CaO + CO 2

and requires an expenditure of 43 kcal per mole of limestone. Calcination of CaCO 3 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 schemes:

CaO + H 2 O \u003d Ca (OH) 2 and Ca (OH) 2 + CO 2 \u003d CaCO 3 + H 2 O.

A colossally important practical role is played by cement, a greenish-gray powder consisting of a mixture of various silicates and calcium aluminates. When mixed with water, it hardens due to hydration. In its production, a mixture of CaCO 3 with clay is fired before sintering (1400-1500 about C). The mixture is then ground. The composition of cement can be expressed as a percentage of the components Cao, SiO 2, Al 2 O 3, Fe 2 O 3, with Cao being a base, and everything else being acid anhydrides. The composition of silicate (Portland) cement is composed mainly of Ca 3 SiO 5 , Ca 2 SiO 4 , Ca 3 (AlO 3) 2 and Ca (FeO 2) 2 . His grasp goes according to the schemes:

Ca 3 SiO 5 + 3H 2 O \u003d Ca 2 SiO 4. 2H 2 O + Ca (OH) 2

Ca 2 SiO 4 + 2H 2 O \u003d Ca 2 SiO 4. 2H 2 O

Ca 3 (AlO 3) 2 + 6H 2 O \u003d Ca 3 (AlO 3) 2. 6H 2 O

Ca (FeO 2) 2 + nH 2 O \u003d Ca (FeO 2) 2. nH2O.

Natural chalk is introduced into the composition of various putties. Fine-crystalline, precipitated from a solution of CaCO 3 is part of tooth powders. BaO is obtained from BaCO 3 by calcination with coal according to the scheme:

VaCO 3 + C \u003d BaO + 2CO.

If the process is carried out at a higher temperature in a stream of nitrogen, cyanide barium:

VaCO 3 + 4C + N 2 \u003d 3CO + Ba (CN) 2.

Ba(CN) 2 is highly soluble in water. Ba(CN) 2 can be used to produce other metal cyanides by exchange decomposition with sulfates. Bicarbonates E are soluble in water and can only be obtained in solution, for example, by passing carbon dioxide into a suspension of CaCO 3 in water:

CO 2 + CaCO 3 + H 2 O \u003d Ca (HCO 3) 2.

This reaction is reversible and shifts to the left when heated. The presence of calcium and magnesium bicarbonates in natural waters causes water hardness.

Properties of alkaline earth metals

Physical properties

Alkaline earth metals (compared to alkali metals) have higher t╟pl. and t╟bp., ionization potentials, densities and hardness.

Chemical properties

1. Very reactive.

2. Have a positive valence of +2.

3. React with water at room temperature (except for Be) with evolution of hydrogen.

4. They have a high affinity for oxygen (reducing agents).

5. They form salt-like hydrides EH 2 with hydrogen.

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

Receipt

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 remaining metals are obtained by electrolysis of chloride melts:

CaCl 2 \u003d 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; in compounds, they 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 \u003d Ca (OH) 2 + H 2 ╜

2. Reaction with oxygen.

All metals form oxides RO, barium peroxide BaO 2:

2Mg + O 2 \u003d 2MgO

Ba + O 2 \u003d BaO 2

3. Binary compounds are formed with other non-metals:

Be + Cl 2 = BeCl 2 (halides)

Ba + S = BaS (sulfides)

3Mg + N 2 \u003d 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 \u003d CaCl 2 + H 2 ╜

Mg + H 2 SO 4 (razb.) \u003d MgSO 4 + H 2 ╜

Beryllium also dissolves in aqueous solutions of alkalis:

Be + 2NaOH + 2H 2 O \u003d Na 2 + H 2 ╜

5. Qualitative reaction to alkaline earth metal cations - 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 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

Receipt

1) Oxidation of metals (except Ba, which forms a 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 \u003d Mg (OH) 2

3CaO + P 2 O 5 \u003d Ca 3 (PO 4) 2

BeO + 2HNO 3 \u003d Be (NO 3) 2 + H 2 O

BeO - amphoteric oxide, soluble in alkalis:

BeO + 2NaOH + H 2 O \u003d Na 2

Alkaline earth metal hydroxides R(OH) 2

Receipt

Reactions of alkaline earth metals or their oxides with water: Ba + 2H 2 O \u003d Ba (OH) 2 + H 2

CaO (quicklime) + H 2 O \u003d Ca (OH) 2 (slaked lime)

Chemical properties

Hydroxides R (OH) 2 - white crystalline substances, soluble in water worse 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 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) The 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. Hard water when boiled forms a scale, it does not boil soft food products; detergents do not produce foam.

Carbonate (temporary) hardness is due to the presence of calcium and magnesium bicarbonates in water, non-carbonate (permanent) hardness - 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 the solution:

1) 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) by 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 \u003d CaCO 3 ¯+ 2NaHCO 3

CaSO 4 + Na 2 CO 3 \u003d CaCO 3 ¯ + Na 2 SO 4

MgCl 2 + Na 2 CO 3 \u003d MgCO 3 ¯ + 2NaCl

All four methods are used to remove temporary stiffness, and only the last two are used for permanent hardness.

Thermal decomposition of nitrates.

E (NO3) 2 \u003d t \u003d 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

alkaline earth metals and, alkaline earth metals chemistry
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, less often 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 of this group, magnesium, in some respects differs significantly from the alkaline earth metals in a number of chemical properties. All alkaline earth metals are gray solids at room temperature. unlike alkali metals, they are much harder, and they are mostly not 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) .

Some atomic and physical properties of alkaline earth metals

Atomic room	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	tboil, °C	ρ, g/cm³	ΔHpl, kJ mol−1	ΔHboil, 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 an electronic configuration of the external energy level ns², and are s-elements, along with alkali metals. Having two valence electrons, alkaline earth metals easily donate them, and in all compounds they have an oxidation state of +2 (very rarely +1).

The chemical activity of alkaline earth metals increases with increasing serial number. Beryllium in a compact form does not react with either oxygen or halogens even at a red heat temperature (up to 600 ° C, to react with oxygen and other chalcogens, even more heat, fluorine is an exception). Magnesium is protected by an oxide film at room temperature and higher (up to 650 °C) temperatures and does not oxidize further. Calcium oxidizes slowly and at room temperature in depth (in the presence of water vapor), and burns out with slight heating in oxygen, but is stable in dry air at room temperature. Strontium, barium, and radium rapidly oxidize in air to give 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 in basic properties with increasing serial number.

Simple substances

Beryllium reacts with weak and strong acid solutions to form salts:

however, it is passivated with cold concentrated nitric acid.

The reaction of beryllium with aqueous solutions alkali is accompanied by the release of hydrogen and the formation of hydroxoberyllates:

When carrying out the 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 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 - amphoteric oxide, dissolves in concentrated mineral acids and alkalis with the formation of salts:

but with less strong acids and bases, the reaction no longer proceeds.

Magnesium oxide does not react with dilute and concentrated bases, but easily reacts 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, with acids - beryllium salts of acids:

Hydroxides of magnesium, calcium, strontium, barium and radium are bases, the strength increases from weak to very strong, which is the strongest corrosive substance, exceeding potassium hydroxide in activity. They dissolve well in water (except 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 amounts) in nature. Due to their high chemical activity, all of them are not found in the free state. The most common alkaline earth metal is calcium, the amount of which is 3.38% (of the mass of the earth's crust). Magnesium is slightly inferior to it, the amount of which is 2.35% (of the mass of the earth's crust). Barium and strontium are also common in nature, which, respectively, are 0.05 and 0.034% of the mass of the earth's crust. 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 separated from there by chemical means. 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 for the synthesis of ATP, is involved in the transfer nerve impulses, is actively used in medicine (bischophytotherapy, etc.). Calcium is a common macronutrient in plants, animals and humans. in the human body and other vertebrates, most of it is in the skeleton and teeth. Calcium is found in bones in the form of hydroxyapatite. The “skeletons” of most groups of invertebrates (sponges, coral polyps, mollusks, etc.) are composed of various forms of calcium carbonate (lime). Calcium ions are involved in blood coagulation 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, as 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 the radium that enters the body accumulates in bone tissue. Large concentrations radium cause osteoporosis, spontaneous bone fractures and malignant tumors of bones and hematopoietic tissue. Radon, a gaseous radioactive decay product of radium, is also dangerous.

Notes

1. According to the new IUPAC classification. According to the outdated classification, they belong to the main subgroup of group II of the periodic table.
2. Nomenclature of Inorganic Chemistry. IUPAC Recommendations 2005. - International Union of Pure and Applied Chemistry, 2005. - P. 51.
3. Group 2 - Alkaline Earth Metals, Royal Society of Chemistry.
4. Golden fund. School Encyclopedia. Chemistry. M.: Bustard, 2003.

Periodic system chemical elements 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	Up	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