Group IIA contains only metals - Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium) and Ra (radium). The chemical properties of the first representative of this group, beryllium, differ most strongly from the chemical properties of the other elements of this group. Its chemical properties are in many ways even more similar to aluminum than to other group IIA metals (the so-called "diagonal similarity"). Magnesium, in terms of chemical properties, also differs markedly from Ca, Sr, Ba, and Ra, but still has much more similar chemical properties with them than with beryllium. Due to the significant similarity of the chemical properties of calcium, strontium, barium and radium, they are combined into one family, called alkaline earth metals.
All elements of group IIA belong to s-elements, i.e. contain all of their valence electrons s-sublevel. Thus, the electronic configuration of the outer electron layer of all chemical elements of this group has the form ns 2 , where n– number of the period in which the element is located.
Due to the peculiarities of the electronic structure of group IIA metals, these elements, in addition to zero, are capable of having only one single oxidation state, equal to +2. Simple substances formed by elements of group IIA, when participating in any chemical reactions, can only be oxidized, i.e. donate electrons:
Me 0 - 2e - → Me +2
Calcium, strontium, barium and radium are extremely reactive. The simple substances formed by them are very strong reducing agents. Magnesium is also a strong reducing agent. The reducing activity of metals obeys the general laws of the periodic law of D.I. Mendeleev and increases down the subgroup.
Interaction with simple substances
with oxygen
Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of a liquid inert to them, most often kerosene.
Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO 2):
2Mg + O 2 \u003d 2MgO
2Ca + O 2 \u003d 2CaO
2Ba + O 2 \u003d 2BaO
Ba + O 2 \u003d BaO 2
It should be noted that during the combustion of alkaline earth metals and magnesium in air, the reaction of these metals with atmospheric nitrogen also proceeds side by side, as a result of which, in addition to compounds of metals with oxygen, nitrides with the general formula Me 3 N 2 are also formed.
with halogens
Beryllium reacts with halogens only at high temperatures, while the rest of the Group IIA metals already at room temperature:
Mg + I 2 \u003d MgI 2 - magnesium iodide
Ca + Br 2 \u003d CaBr 2 - calcium bromide
Ba + Cl 2 \u003d BaCl 2 - barium chloride
with non-metals of IV–VI groups
All metals of group IIA react when heated with all non-metals of groups IV-VI, but depending on the position of the metal in the group, as well as the activity of non-metals, a different degree of heating is required. Since beryllium is the most chemically inert among all metals of group IIA, its reactions with nonmetals require significantly more about high temperature.
It should be noted that the reaction of metals with carbon can form carbides of various nature. There are carbides related to methanides and conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by a metal. They, like methane, contain carbon in the -4 oxidation state, and during their hydrolysis or interaction with non-oxidizing acids, methane is one of the products. There is also another type of carbides - acetylenides, which contain the C 2 2- ion, which is actually a fragment of the acetylene molecule. Carbides of the acetylenide type upon hydrolysis or interaction with non-oxidizing acids form acetylene as one of the reaction products. What type of carbide - methanide or acetylenide - will be obtained by the interaction of one or another metal with carbon depends on the size of the metal cation. As a rule, methanides are formed with metal ions having a small radius, and acetylides with larger ions. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:
The remaining metals of group II A form acetylenides with carbon:
With silicon, group IIA metals form silicides - compounds of the Me 2 Si type, with nitrogen - nitrides (Me 3 N 2), phosphorus - phosphides (Me 3 P 2):
with hydrogen
All alkaline earth metals react when heated with hydrogen. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, an increased pressure of hydrogen is also required. Beryllium does not react with hydrogen under any conditions.
Interaction with complex substances
with water
All alkaline earth metals actively react with water to form alkalis (soluble metal hydroxides) and hydrogen. Magnesium reacts with water only during boiling, due to the fact that when heated, the protective oxide film of MgO dissolves in water. In the case of beryllium, the protective oxide film is very resistant: water does not react with it either when boiling or even at a red heat temperature:
with non-oxidizing acids
All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the activity series to the left of hydrogen. In this case, a salt of the corresponding acid and hydrogen are formed. Reaction examples:
Be + H 2 SO 4 (razb.) \u003d BeSO 4 + H 2
Mg + 2HBr \u003d MgBr 2 + H 2
Ca + 2CH 3 COOH = (CH 3 COO) 2 Ca + H 2
with oxidizing acids
− dilute nitric acid
All Group IIA metals react with dilute nitric acid. In this case, the reduction products instead of hydrogen (as in the case of non-oxidizing acids) are nitrogen oxides, mainly nitric oxide (I) (N 2 O), and in the case of highly diluted nitric acid, ammonium nitrate (NH 4 NO 3):
4Ca + 10HNO 3 ( razb .) \u003d 4Ca (NO 3) 2 + N 2 O + 5H 2 O
4Mg + 10HNO3 (very disaggregated)\u003d 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O
− concentrated nitric acid
Concentrated nitric acid at ordinary (or low) temperature passivates beryllium, i.e. does not react with it. When boiling, the reaction is possible and proceeds mainly in accordance with the equation:
Magnesium and alkaline earth metals react with concentrated nitric acid to form a wide range of different nitrogen reduction products.
− concentrated sulfuric acid
Beryllium is passivated with concentrated sulfuric acid, i.e. does not react with it under normal conditions, however, the reaction proceeds during boiling and leads to the formation of beryllium sulfate, sulfur dioxide and water:
Be + 2H 2 SO 4 → BeSO 4 + SO 2 + 2H 2 O
Barium is also passivated by concentrated sulfuric acid due to the formation of insoluble barium sulfate, but reacts with it when heated, barium sulfate dissolves when heated in concentrated sulfuric acid due to its conversion to barium hydrogen sulfate.
The remaining metals of the main group IIA react with concentrated sulfuric acid under any conditions, including in the cold. Sulfur reduction can occur to SO 2, H 2 S and S, depending on the activity of the metal, the reaction temperature and the concentration of the acid:
Mg + H 2 SO 4 ( conc .) \u003d MgSO 4 + SO 2 + H 2 O
3Mg + 4H2SO4 ( conc .) \u003d 3MgSO 4 + S↓ + 4H 2 O
4Ca + 5H2SO4 ( conc .) \u003d 4CaSO 4 + H 2 S + 4H 2 O
with alkalis
Magnesium and alkaline earth metals do not interact with alkalis, and beryllium easily reacts both with alkali solutions and with anhydrous alkalis during fusion. Moreover, when the reaction is carried out in an aqueous solution, water is also involved in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and gaseous hydrogen:
Be + 2KOH + 2H 2 O \u003d H 2 + K 2 - potassium tetrahydroxoberyllate
When carrying out the reaction with solid alkali during fusion, beryllates of alkali or alkaline earth metals and hydrogen are formed.
Be + 2KOH \u003d H 2 + K 2 BeO 2 - potassium beryllate
with oxides
Alkaline earth metals, as well as magnesium, can reduce less active metals and some non-metals from their oxides when heated, for example:
The method of restoring metals from their oxides with magnesium is called magnesiumthermy.
S-elements 2 groups
GENERAL CHARACTERISTICS. To alkaline earth metals usually
include calcium, strontium and barium, since their oxides (earths) at
dissolving in water gives alkali. beryllium and magnesium oxides in water
dissolve. Sometimes all metals from group 2A are called
alkaline earth. At the outer level, atoms have 2 electrons (Be -
2s2, Mg - 3s2, Ca - 4s2, etc.).
When excited, s-electrons go to p-
sublevel and then the formation of two bonds is possible
(valency is two). In metal compounds
exhibit an oxidation state of +2.
1. Alkaline earth metals are strong reducing agents, although
inferior to alkali metals. Restorative properties are growing
from top to bottom, which coincides with an increase in atomic radii (Be - 0.113
nm, Ba - 0.221 nm) and the weakening of the bond between electrons and the nucleus. So, Be and Mg
decompose water very slowly, while Ca, Sr, Ba rapidly.
2. In air, Be and Mg are covered with a protective film and burn out when
only when ignited, while Ca, Sr, Ba self-ignite when
contact with air.
3. Be and Mg oxides are insoluble in water and Be and Mg hydroxides
are obtained indirectly, while the oxides Ca, Sr, Ba are combined with
water to form hydroxides. Beryllium oxide has amphoteric
properties, the remaining oxides are the main properties.
4. Be (OH) 2 and Mg (OH) 2 are almost insoluble in water (0.02 and 2 mg per 100 g).
The solubility of Ca(OH)2, Sr(OH)2, Ba(OH)2 is 0.1, 0.7 and 3.4 g.
this Be (OH) 2 is an amphoteric hydroxide, Mg (0H) 2, is a weak base,
Ca(OH)2, Sr(OH)2, Ba(0H)2 are strong bases.
5. The halides are highly soluble in water, but the solubility
sulfate falls from top to bottom. So, 35.6 g dissolves in 100 g of water.
MgSO4, but only 0.2 g CaSO4, 0.01 g SrSO4 and 0.0002 g BaSO4.
6. The solubility of carbonates decreases from top to bottom. MgCO3 - 0.06 g per
100 g of water, BaCO3 in total - 0.002 g. Thermal stability of carbonates
grows from top to bottom: If BeCO3 decomposes at 100o, MgCO3 - at 350o, then
CaCO3 - at 900o, SrCO3 - 1290o BaCO3 - at 1350o.
BERYLLIUM - has more pronounced covalent
(non-metallic) properties than other group 2A elements. And myself
beryllium, its oxide and hydroxide have amphoteric properties.
Be + 2HCl = BeCl2 + H2 Be + 2KOH + 2H2O = K2 + H2
BeO + 2HCl = BeCl2 + H2O BeO + 2KOH + H2O = K2
Be(OH)2 + 2HCl = BeCl2 + 2H2O Be(OH)2 + 2KOH = K2
Magnesium and calcium
GENERAL INFORMATION. The content of magnesium and calcium in the earth's crust 2.1
and 3.6%. Minerals magnesium- MgCO3. CaCO3 - dolomite, MgCO3 - magnesite, KCl.
6H2O - carnallite; MgSO4
KCl. 3H2O - kainite. Minerals calcium:
CaCO3 - calcite (limestone, chalk, marble), СaSO4
2H2O - gypsum, Ca3(PO4)2 -
phosphorite, 3Ca3(PO4)2
CaF2 - apatite.
Magnesium and calcium - silver-white metals melt at 651 and
851o C. Calcium and its salts color the flame brick red.
RECEIVING. Calcium and magnesium are obtained by electrolysis of the melt
calcium chloride or magnesium chloride or by the aluminothermic method.
electrolysis to
СaCl2 Ca + Cl2 4CaO + 2Al = 3Ca + CaO . Al2O3
Chemical properties of calcium and magnesium.
In compounds, both metals exhibit an oxidation state of +2. At
In this case, calcium is more active than magnesium, although it is inferior to strontium and
1. Interaction with oxygen comes with ignition and
release of heat and light.
Mg + O2 = 2MgO; 2Ca + O2 = 2CaO
2. Interaction with halogens. Fluorine combines with Ca and Mg
directly, the remaining halogens only when heated.
Mg + Cl2 = MgCl2; Ca + Br2 = CaBr2
3. When heated, Ca and Mg form hydrides with hydrogen, which
easily hydrolyzed and oxidized. to to
Mg + H2 = MgH2; Ca + H2 = CaH2
CaH2 + 2H2O = Ca(OH)2 + 2H2; CaH2 + O2 = CaO + H2O
4. When heated, both metals interact with others
non-metals:
Mg + S = MgS; 3Ca + N2 = Ca3N2; 3Mg + 2P = Mg3P2
3Ca + 2As = Mg3As2; Ca + 2C = CaC2; Mg + 2C = MgC2
Nitrides, sulfides and carbides of calcium and magnesium are susceptible to
hydrolysis:
Ca3N2 + 6H2O = 3Ca(OH)2 + 2NH3; CaC2 + 2H2O = Ca(OH)2 +
5. Beryllium and magnesium interact with water and alcohols only
when heated, while calcium violently displaces from them
Mg + H2O = MgO + H2; Ca + 2H2O = Ca(OH)2 + H2
Ca + 2C2H5OH \u003d Ca (C2H5O) 2 + H2
6. Magnesium and calcium take away oxygen from less active oxides
metals.
CuO + Mg = Cu + MgO; MoO3 + 3Ca = Mo + 3CaO
7. Magnesium and calcium displace hydrogen from non-oxidizing acids,
and oxidizing acids deeply reduce these metals.
Mg + 2HCl = MgCl2 + H2; Ca + 2CH3COOH = Ca(CH3COO)2 + H2
3Mg + 4H2SO4c = 3MgSO4 + S + 4H2O; 4Ca + 10HNO3c = 4Ca(NO3)2 + N2O
4Ca + 10HNO3 = 4Ca(NO3)2 + NH4NO3 + 3H2O
8. Calcium and magnesium are easily oxidized by solutions of oxidizing agents:
5Mg + 2KMnO4 + 8H2SO4 = 5MgSO4 + 2MnSO4 + K2SO4 + 8H2O
Ca + K2Cr2O7 + 7H2SO4 = 3CaSO4 + Cr2(SO4)3 + K2SO4 + 7H2O
Oxides calcium and magnesium hydroxides.
Magnesium oxide - MgO- white powder, refractory (refractory),
insoluble in water and acids and only amorphous oxide form
magnesium reacts slowly with acids. Get magnesium oxide
heating magnesium hydroxide.
MgO (amorphous) + 2HCl = MgCl2 + H2O; Mg(OH)2 = MgO + H2O
Magnesium hydroxide - Mg(OH)2- insoluble and
low dissociating base. Obtained by the action of alkalis on salts
magnesium. When carbon dioxide is passed through its solution,
precipitate of magnesium carbonate, which subsequently dissolves when
excess CO2.
MgCl2 + 2KOH = Mg(OH)2 + 2KCl MgCl2 + 2NH4OH = Mg(OH)2 + 2NH4Cl
Mg(OH)2 + CO2 = MgCO3 + H2O MgCO3 + CO2 + H2O = Mg(HCO3)2
Calcium oxide - Cao- quicklime. White refractory
a substance with pronounced basic properties (forms with water
hydroxide, reacts with acid oxides, acids and amphoteric
oxides).
CaO + H2O = Ca(OH)2 CaO + CO2 = CaCO3 CaO + 2HCl = CaCl2
CaO + Al2O3 = Ca(AlO2)2 CaO + Fe2O3 = Ca(FeO2)2
Obtained by roasting limestone or sulfate reduction
CaCO3 = CaO + CO2; 2СаSO4 + 2C = 2CaO + 2SO2 + CO2
calcium hydroxide Ca(OH)2- slaked lime (fluff), get
when calcium oxide reacts with water. Strong base except
It dissolves some non-metals and amphoteric metals.
Ca(OH)2 + 2HCl = CaCl2 + 2H2O Ca(OH)2 + SO3 = CaSO4 +
3Ca(OH)2 2FeCl3 = 2Fe(OH)3+ 3CaCl2 2NH4Cl + Ca(OH)2 = CaCl2 + NH3
2Ca(OH)2 + Cl2 = CaCl2 + Ca(ClO)2 + 2H2O Ca(OH)2 + 2Al + 2H2O =
Slaked lime is part of the mortar.
Solidification is based on the reactions:
Ca(OH)2 + CO2 = CaCO3 + H2O; Ca(OH)2 + SiO2 = CaSiO3 + H2O
sand from the air
When carbon dioxide is passed through a solution of Ca(OH)2
(lime water), calcium carbonate precipitates, which, when
further transmission of CO2 dissolves due to the formation
soluble calcium bicarbonate.
Ca(OH)2 + CO2 = CaCO3 + H2O; CaCO3 + CO2 + H2O = Ca(HCO3)2
Properties of elements of II A group.
|
Properties |
4Be |
12Mg |
20Ca |
38Sr |
56Ba |
88Ra |
|
Atomic mass |
9,012 |
24,305 |
40,80 |
87,62 |
137,34 |
226,025 |
|
Electronic configuration* |
||||||
|
0,113 |
0,160 |
0,190 |
0,213 |
0,225 |
0,235 |
|
|
0,034 |
0,078 |
0,106 |
0,127 |
0,133 |
0,144 |
|
|
Ionization energy |
9,32 |
7,644 |
6,111 |
5,692 |
5,21 |
5,28 |
|
Relative electro- |
1,5 |
1,2 |
1,0 |
1,0 |
0,9 |
0,9 |
|
Possible oxidation states |
||||||
|
clarke, at.% (distribu- |
1*10 -3 |
1,4 |
1,5 |
8*10 -3 |
5*10 -3 |
8*10 -12 |
|
State of aggregation (well.). |
S E R D E S E S T V A |
|||||
|
Colour |
Gray |
Silver |
S E R E B R I S T O - WHITE |
|||
|
1283 |
649,5 |
850 |
770 |
710 |
700 |
|
|
2970 |
1120 |
1487 |
1367 |
1637 |
1140 |
|
|
Density |
1,86 |
1,741 |
1,540 |
2,67 |
3,67 |
|
|
Standard electrode potential |
1,73 |
2,34 |
2,83 |
2,87 |
2,92 |
|
*The configurations of the external electronic levels of atoms of the corresponding elements are given. The configurations of the remaining electronic levels coincide with those for the noble gases that complete the previous period and are indicated in parentheses.
As follows from the data given in the table, the elements of group IIA have low (but still not the lowest: compare with IA gr.) ionization energy and relative electronegativity, and these values decrease from Be to Ba, which allows us to conclude that that these elements are typical reducing metals, and Ba is more active than Be.
Ve - exhibits, like aluminum, amphoteric properties. However, in Be, the metallic properties are still more pronounced than the non-metallic ones. Beryllium reacts, unlike the other elements of group IIA, with alkalis.
Chemical bonds in Be compounds are mainly covalent, while bonds in compounds of all other elements (Mg - Ra) are ionic in nature. At the same time, as with group IA elements, bonds with halogens and oxygen are very strong, and with hydrogen, carbon, nitrogen, phosphorus and sulfur they are easily hydrolyzed.
physical properties. These are silver-white metals, relatively light, soft (with the exception of beryllium), ductile, fusible (everything except beryllium), have good electrical and thermal conductivity.
Practical use. Be is used in nuclear technology as a neutron moderator and absorber. Alloys of beryllium with copper - bronze - are very resistant, and with nickel - they have high chemical resistance, due to which they are used in surgery.
Mg, Ca - are used as good reducing agents in metallothermy.
Ca, Sr, Ba - react quite easily with gases and are used as getters (absorbers from the air) in vacuum technology.
Receipt. Being highly reactive, alkaline earth metals do not occur in nature in a free state, they are obtained by electrolysis of halide melts or by metallothermy. In nature, alkaline earth elements are part of the following minerals: -beryl; - feldspar; - bischofite - used in medicine and for the production of magnesium by electrolysis. To obtain beryllium in metallurgy, fluoroberyllates are used: .
Chemical properties. Alkaline earth metals easily react with oxygen, halogens, non-metals, water and acids, especially when heated:
This reaction is especially easy for calcium and barium, so they are stored under special conditions.
Barium persulfide BaS is a phosphor.
Hydrolysis of acetylides produces acetylene:
It was not possible to obtain compounds of Be and Mg with hydrogen by direct interaction of simple substances: the reaction does not go whereas goes pretty easy. The resulting hydrides are strong reducing agents. passivation, no reaction
Oxides of alkaline earth metals. Oxides of alkaline earth elements are widely used in construction. They are obtained by decomposition of salts: - CaO - quicklime.
In the series of oxides from BeO to BaO from left to right, the solubility of oxides in water, their main properties and chemical activity increases, as follows: BeO is insoluble in water, amphoterene, MgO is slightly soluble in water, and CaO, SrO, BaO are highly soluble in water with the formation of hydroxides Me (OH) : .
The melting points of oxides decrease in the series BeO ® BaO. Melting points of oxides BeO and MgO » 2500 ° C, which allows them to be used as refractory materials.
Hydroxides of alkaline earth metals. In the series Be (OH) 2 ® Ba (OH) 2, the radius of Me 2+ ions increases, and, as a result, the probability of manifestation of the main properties of hydroxides, their solubility in water increases: Be (OH) 2 - is slightly soluble in water, due to its amphoteric exhibits weak acidic and basic properties, and Ba (OH) 2 is highly soluble in water and can be compared in strength with such a strong base as NaOH.
The amphotericity of beryllium hydroxide can be illustrated by the following reactions:
Salts of alkaline earth metals. Soluble salts Be and Ba - toxic, poisonous! CaF 2- sparingly soluble salt, occurs in nature as fluorite or fluorspar, is used in optics. CaCl 2 , MgCl 2- highly soluble in water, are used in medicine and chemical synthesis as desiccants. Carbonates are also widely used in construction: CaCO 3H MgCO 3- dolomite - used in construction and to obtain Vg and Ca. CaCO 3 - calcite, chalk, marble, Icelandic spar, MgCO 3- magnesite. The content of soluble carbonates in natural water determines its hardness: . Sulfates are also widespread natural compounds of alkaline earth metals: CaSO 4H 2H 2 O- gypsum - widely used in construction. MgSO 4H 7H 2 O- epsomite, "English bitter salt", BaSO 4- finds application with fluoroscopy. Phosphates: Ca 3 (RO 4) 2- phosphorite, Ca (H 2 RO 4) 2, CaHRO 4- precipitate - used for the production of fertilizers, Ca 5 (RO 4) 3H (OH -, F -, Cl -) - appatite - natural mineral Ca, NH 4 Mg (PO 4)- slightly soluble compound. Other salts are also known: Ca (NO 3) 2H 2H 2 O- Norwegian saltpeter, Mg(ClO 4) 2- Anhydrone is a very good desiccant.
Alkaline earth metals include metals of group IIA of the Periodic Table of D.I. Mendeleev - calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). In addition to them, the main subgroup of group II includes beryllium (Be) and magnesium (Mg). The outer energy level of alkaline earth metals has two valence electrons. The electronic configuration of the external energy level of alkaline earth metals is ns 2 . In their compounds, they exhibit a single oxidation state equal to +2. In OVR, they are reducing agents, i.e. donate an electron.
With an increase in the charge of the nucleus of atoms of elements that are part of the group of alkaline earth metals, the ionization energy of atoms decreases, and the radii of atoms and ions increase, the metallic features of chemical elements increase.
Physical properties of alkaline earth metals
In the free state, Be is a steel-gray metal with a dense hexagonal crystal lattice, rather hard and brittle. In air, Be is covered with an oxide film, which gives it a matte tint and reduces its chemical activity.
Magnesium in the form of a simple substance is a white metal, which, like Be, acquires a matte hue when exposed to air due to the formation of an oxide film. Mg is softer and more ductile than beryllium. The crystal lattice of Mg is hexagonal.
Free Ca, Ba and Sr are silver-white metals. When exposed to air, they are instantly covered with a yellowish film, which is the products of their interaction with the constituent parts of the air. Calcium is a rather hard metal, Ba and Sr are softer.
Ca and Sr have a cubic face-centered crystal lattice, barium has a cubic body-centered crystal lattice.
All alkaline earth metals are characterized by the presence of a metallic type of chemical bond, which causes their high thermal and electrical conductivity. The boiling and melting points of alkaline earth metals are higher than those of alkali metals.
Obtaining alkaline earth metals
Getting Be is carried out by the reduction reaction of its fluoride. The reaction proceeds when heated:
BeF 2 + Mg = Be + MgF 2
Magnesium, calcium and strontium are obtained by electrolysis of molten salts, most often chlorides:
CaCl 2 \u003d Ca + Cl 2
Moreover, when Mg is obtained by electrolysis of a dichloride melt, NaCl is added to the reaction mixture to lower the melting temperature.
To obtain Mg in industry, metal- and carbon-thermal methods are used:
2(CaO×MgO) (dolomite) + Si = Ca 2 SiO 4 + Mg
The main way to obtain Ba is oxide reduction:
3BaO + 2Al = 3Ba + Al 2 O 3
Chemical properties of alkaline earth metals
Since in n.a. the surface of Be and Mg is covered with an oxide film - these metals are inert with respect to water. Ca, Sr and Ba dissolve in water to form hydroxides exhibiting strong basic properties:
Ba + H 2 O \u003d Ba (OH) 2 + H 2
Alkaline earth metals are able to react with oxygen, and all of them, with the exception of barium, form oxides as a result of this interaction, barium - peroxide:
2Ca + O 2 \u003d 2CaO
Ba + O 2 \u003d BaO 2
Oxides of alkaline earth metals, with the exception of beryllium, exhibit basic properties, Be - amphoteric properties.
When heated, alkaline earth metals are capable of interacting with non-metals (halogens, sulfur, nitrogen, etc.):
Mg + Br 2 \u003d 2MgBr
3Sr + N 2 \u003d Sr 3 N 2
2Mg + 2C \u003d Mg 2 C 2
2Ba + 2P = Ba 3 P 2
Ba + H 2 = BaH 2
Alkaline earth metals react with acids - dissolve in them:
Ca + 2HCl \u003d CaCl 2 + H 2
Mg + H 2 SO 4 \u003d MgSO 4 + H 2
Beryllium reacts with aqueous solutions of alkalis - it dissolves in them:
Be + 2NaOH + 2H 2 O \u003d Na 2 + H 2
Qualitative reactions
A qualitative reaction to alkaline earth metals is the coloring of the flame with their cations: Ca 2+ colors the flame dark orange, Sr 2+ dark red, Ba 2+ light green.
A qualitative reaction to the barium cation Ba 2+ are SO 4 2- anions, resulting in the formation of a white precipitate of barium sulfate (BaSO 4), insoluble in inorganic acids.
Ba 2+ + SO 4 2- \u003d BaSO 4 ↓
Examples of problem solving
EXAMPLE 1
| Exercise | Carry out a series of transformations: Ca → CaO → Ca (OH) 2 → Ca (NO 3) 2 |
| Decision | 2Ca + O 2 → 2CaO CaO + H 2 O→Ca(OH) 2 Ca(OH) 2 + 2HNO 3 → Ca(NO 3) 2 + 2H 2 O |
