معرفی عناصر گروه 3

Introduction group IIIa

Group 3a will include the following elements: boron B, aluminium Al, gallium Ga, indium In and thallium Tl.

Boron is a unique and exciting element. Over the years it has proved a constant challenge and stimulus not only to preparative chemists and theoreticians, but also to industrial chemists and technologists. It is the only non-metal in Group 3a of the periodic table and shows many similarities to its nelghbour, carbon, and its diagonal relative, silicon. Thus, like carbon and silicon, it shows a marked propensity to form covalent, molecular compounds, but it differs sharply from them in having one less valence electron than the number of valence orbitals, a situation sometimes referred to as "electron deficiency". This has a dominant effect on its chemistry. Borax was known in the ancient world where it was used to prepare glazes and hard (borosilicate) glasses. Sporadic investigations during the eighteenlh century led ultimately to the isolalion of very impure boron by H. Davy and by J. L. Gay Lussac and L. J. Thenard in 1808, but it was not until 1892 that H. Moissan obtained samples of 95-98% purity by reducing B2O3 with Mg.

Aluminium derives its name from alum, the double sulfate KAl(SO4)2-12H20, which was used medicinally as an astringent in ancient Greece and Rome . Humphry Davy was unable to isolate the metal but proposed the name "alumium" and then "aluminum"; this was soon modified to aluminium and this form is used throughout the world except in North America where the ACS decided in 1925 to adopt "aluminum" in its publications. The impure metal was first isolated by the Danish scientist H. C. Oersted using the reaction of dilute potassium amalgam on AlCl3. This method was improved in 1827 by H. Wohler who used metallic potassium, but the first commercially successful process was devised by H. St.C. Deville in 1854 using sodium. In the same year both he and R. W. Bunsen independently obtained metallic aluminium by electrolysis of fused NaAlCl4.

Gallium was predicted as eka-aluminium by D. I. Mendeleev in 1870 and was discovered by P. E. Lecoq de Boisbaudran in 1875 by means of the spectroscope; de Boisbaudran was, in fact, guided at the time by an independent theory of his own and had been searching for the missing element for some years. The first indications came with the observation of two new violet lines in the spark spectrum of a sample deposited on zinc, and within a month he had isolated 1 g of the metal starting from several hundred kilograms of crude zinc blende ore. The element was named in honour of France (Latin Gallia) and the striking similarity of its physical and chemical properties to those predicted by Mendeleev did much to establish the general acceptance of the Periodic Law; indeed, when de Boisbaudran first stated that the density of Ga was 4.7g/cm3 rather than the predicted 5.9g/cm3, Mendeleev wrote to him suggesting that he redetermine the figure (the correct value is 5.904 g/cm3).

Indium and thallium were also discovered by means of the spectroscope as their names indicate. Indium was first identified in 1863 by F. Reich and H.T. Richter and named from the brilliant indigo blue line in its flame spectrum (Latin indicum). Thallium was discovered independently by W. Crookes and by C. A. Lamy in the preceding year 1861-1862 and named after the characteristic bright green line in its flame spectrum.

کل واکنش های عنصر بور

BORON - B

Properties of Boron:

Nonmetal. Gray-black (crystal) or brown (amorphous). Refractory, very hard, brittle. Chemically passive. Does not react with hydrogen, water, diluted acids, diluted alkalis into the solution. Reacts with water steam, concentrated nitric acid, halogen, nitrogen, hydrogen fluoride and hydrogen sulfide, ammonia and alkalis when heated.

 

Molar mass	g/mol	10.811
Density	g/cm3	2.340
Melting point	°C	2075
Boiling point	°C	3700

 

Methods for the preparation of Boron:

B2H6 = 2B + 3H2 (300-500°C)

B2O3 + 2Al = Al2O3 + 2B (800-900°C)

2BCl3 + 3H2 = 2B + 6HCl (800-1200°C)

2BI3 = 2B + 3I2 (over 700°C or on light)

 

Сhemical reactions with Boron:

2B + 3H2O(steam) = B2O3 + 3H2 (700-800°C).

B + 3HNO3(hot, conc.) = B(OH)3↓ + 3NO2↑.

2B(amorphous) + 2NaOH(conc.) + 6H2O = 2Na[B(OH)4] + 3H2↑.

4B + 4NaOH + 3O2 = 4NaBO2 + 2H2O (350-400°C).

4B + 3O2 = 2B2O3 (700°C, burning on air).

2E + 3E2 = 2BE3 (30°C, E = F; over 400°C E = Cl, Br, I).

2B + 3S = B2S3 (over 600°C).

2B + N2 = 2BN (900-1000°C).

B + P(red) = BP (900-1200°C).

4B + C(graphite) = B4C (over 2000°C, impurity B13C2).

2B + 6HE = BE3 + 3H2 (400-500°C; E = F, Cl).

2B + H2S = B2S3 + 3H2 (800-900°C).

2B + 2NH3 = 2BN + 3H2 (1000-1200°C).

5B + 3NO = B2O3 + 3BN (800°C).

2B + 3CO = B2O3 + 3C(graphite) (1400°C).

4B + 3CS2 = 2B2S3 + 3C(graphite) (930°C).

4B + 3SiO2 = 2B2O3 + 3Si (1300-1500°C).

کل واکنش های عنصر آلومینیوم

ALUMINUM - Al

Properties of Aluminum:

White light plastic metal. Passivated in water, in the concentrated nitric acid and in solution of a potassium bichromate because of formation of a steady oxidic film. The amalgamated metal reacts with water. The strong reducer, high reactivity. Shows amphoteric properties, reacts with diluted acids and alkalis.

 

Molar mass	g/mol	26.982
Density	g/cm3	2.702
Melting point	°C	660.37
Boiling point	°C	2500

 

Methods for the preparation of Aluminum:

2Al2O3→Electrolysis(in melt of Na3[AlF6])→4Al(cathode)+3O2(anode)↑ (900°C).

2AlCl3(liquid) → Electrolysis → 2Al(on cathode) + 3Cl2(on anode)↑.

 

Сhemical reactions with metal Aluminum:

2(Al, Hg) + 6H2O = 2Al(OH)3↓ + 3H2↑ + 2Hg↓ (amalgam, normal temp.)

2Al + 6HCl(diluted) = 2AlCl3 + 3H2↑.

8Al + 30HNO3(diluted) = 8Al(NO3)3 + 3NO2 + 15H2O.

8Al + 30HNO3(high diluted) = 8Al(NO3)3 + 3NH4NO3 + 9H2O.

2Al + 2(NaOH·H2O) = 2NaAlO2 + 3H2O (400-500°C).

2Al + 2NaOH(conc.) + 6H2O(hot) = 2Na[Al(OH)4] + 3H2↑.

8Al + 18H2O + 3KNO3 + 5KOH = 8K[Al(OH)4] + NH3↑ (boiling).

4Al(powder) + 3O2 = 2Al2O3 (burning on air).

2Al + 3F2 = 2AlF3 (600°C).

2Al(powder) + 3E2 = 2AlE3 (25°C; E = Cl, Br).

2Al(powder) + 3I2 = 2AlI3 (25°C; catalyst - water drop).

2Al + 3S = Al2S3 (150-200°C).

2Al(powder) + N2 = 2AlN (800-1200°C).

4Al + P₄ = 4AlP (500-800°C, in the atm. H2).

4Al + 3C(graphite) = Al4C3 (1500-1700°C).

2Al + 6HF(gas) = 2AlF3 + 3H2 (450-500°C).

2Al + 3H2S = Al2S3 + 3H2 (600-1000°C).

2Al + 2NH3 = 2AlN + 3H2 (over 600°C).

8Al + 3(Fe^IIFe^III₂)O₄ = 4Al2O3 + 9Fe (over 2000°C).

کل واکنش های عنصر گالیم

GALLIUM - Ga

Properties of Gallium:

Silver-white with a bluish shade, fusible, very soft, plastic metal. In solid and liquid states is formed by molecules of Ga₂, gas - monatomic. It is passivated in cold water (is formed steady oxidic film). Strong reducer. Reacts with hot water, strong acids, alkalis, hydrate of ammonia, nonmetals.

 

Molar mass	g/mol	69.723
Density	g/cm3	5.904
Melting point	°C	29.78
Boiling point	°C	2403

 

Methods for the preparation of Gallium:

Ga2O3 + 3H2 = 2Ga + 3H2O (700°C).

2GaCl3(melt) → Electrolysis → 2Ga↓(on cathode) + 3Cl2↑(on anode).

 

Сhemical reactions with Gallium:

2Ga + 6H2O(hot) = 2Ga(OH)3↓ + 3H2↑.

2Ga + 4H2O(steam) = 2GaO(OH) + 3H2 (350°C).

2Ga + 6HCl(diluted) = 2GaCl3 + 3H2↑.

Ga + 6HNO3(conc.) = Ga(NO3)3 + 3NO2↑ + 3H2O.

2Ga + 2NaOH(conc., hot) + 6H2O = 2Na[Ga(OH)4] + 3H2↑.

2Ga + 2(NH3·H2O)(conc., hot) + 6H2O = 2NH4[Ga(OH)4] + 3H2↑.

2Ga + 2Na2CO3(conc.) + 8H2O = 2Na[Ga(OH)4] + 3H2↑ + 2NaHCO3.

2Ga + O2 = 2GaO (burning on air).

2Ga + 3Cl2 = 2GaCl3 (80-200°C).

2Ga + 3S = Ga2S3 (800°C).

2Ga + 3H2S = Ga2S3 + 3H2 (250-350°C).

2Ga + 2NH3 = 2GaN + 3H2 (1050-1200°C).

کل واکنش های عنصر ایندیم

INDIUM - In

Properties of Indium:

Silver-white, very soft, plastic, fusible metal. Doesn't change in damp air. Doesn't react with water, alkalis, hydrate of ammonia. Reducer. Is oxidized by acids, oxygen, and other nonmetals.

 

Molar mass	g/mol	114.82
Density	g/cm3	7.30
Melting point	°C	156.634
Boiling point	°C	2024

 

Methods for the preparation of Indium:

In2O3 + 3H2 = 2In + 3H2O (700°C).

In2O3 + 3C(graphite) = 2In + 3CO (800-900°C).

2In2(SO4)3 + 6H2O → Electrolysis → 4In↓(on cathode) + 3O2↑(on anode) + 6H2SO4.

 

Сhemical reactions with Indium:

2In + 6HCl(diluted) = 2InCl3 + 3H2↑.

In + 2HCl(gas) = InCl2(gas) + H2 (700-970°C).

In + 4HNO3(diluted, hot) = In(NO3)3 + NO↑ + 2H2O.

4In + 3O2 = 2In2O3 (800°C, burning on air).

2In + 3Cl2 = 2InCl3 (120-150°C).

2In + 3S = In2S3 (1050-1100°C).

2In + CO2 = In2O + CO (850°C).

2In + H2S = In2S + H2 (700-800°C).

کل واکنش های عنصر تالیم

THALLIUM - Tl

Properties of Thallium:

Silver-white metal, plastic, very soft, fusible. On air becomes covered by an oxidic film. In a compact form doesn't react with water, hydrochloric acid, alkalis, hydrate of ammonia. Is oxidized by sulfuric and nitric acids, hydrogen peroxide, and chlorine.

 

Molar mass	g/mol	204.383
Density	g/cm3	11.84
Melting point	°C	303.6
Boiling point	°C	1457

 

Methods for the preparation of Thallium:

Tl2O(solid) = Tl2O(gas) &harr Tl + O2 (over 1300°C).

Tl2O + H2 = 2Tl + H2O (over 500°C).

Tl2O + CO = 2Tl + CO2 (250-325°C).

2Tl2SO4 + 2H2O → Electrolysis → 4Tl↓(on cathode) + O2↑(on anode) + 2H2SO4.

2TlCl + H2 = 2Tl + 2HCl (650-750°C).

 

Сhemical reactions with Thallium:

2Tl + H2SO4(diluted, cold) = Tl2SO4 + H2↑.

3Tl + 4HNO3(diluted, hot) = 3TlNO3 + NO↑ + 2H2O.

Tl + 6HNO3(conc., hot) = Tl(NO3)3 + 3NO2↑ + 3H2O.

4Tl + 2O2 = Tl2O + Tl2O3 (400°C, burning on air).

4Tl + 2H2O = + O2 = 4TlOH (50-70°C).

2Tl + 3H2O2(conc.) = Tl2O3↓ + 3H2O.

2Tl + Cl2 = 2TlCl (normal temp.).

2Tl + 2HCl(conc.) + 3Cl2 = 2H[TlCl4].

2Tl + S = Tl2S (320°C, in the atm. of hydrogen).

2Tl + 3S = Tl2S3 (200-250°C).

معرفی عناصر گروه 2

Introduction group IIa

The Group 2a (beryllium Be, magnesium Mg, calcium Ca, strontium Sr, barium Ba, radium Ra) or alkaline earth metals exemplify and continue the trends in properties noted for the alkali metals.

The discovery of beryllium in 1798 followed an unusual train of events. The mineralogist R.-J. Haily had observed the remarkable similarity in external crystalline structure, hardness and density of a beryl from Limoges and an emerald from Peru, and suggested to L.N. Vauquelin that he should analyse them to see if they were chemically identical. As a result, Vauquelin showed that both minerals contained not only alumina and silica as had previously been known, but also a new earth, beryllia, which closely resembled alumina but gave no alums, apparently did not dissolve in an excess of KOH and had a sweet rather than an astringent taste. Caution: beryllium compounds are now known to be extremely toxic, especially as dusts or smokes; it seems likely that this toxicity results from the ability of Be to displace Mg from Mg-activated enzymes due to its stronger coordinating ability.

Both beryl and emerald were found to be essentially Be3Al2Si6O18, the only difference between them being that emerald also contains about 2% Cr, the source of its green color. The combining weight of Be was 4.7 but the similarity between Be and Al led to considerable confusion concerning the valency and atomic weight of Be (2 x 4.7 or 3 x 4.7); this was not resolved until Mendeleev 70 years later stated that there was no room for a tervalent element of atomic weight 14 near nitrogen in his periodic table, but that a divalent element of atomic weight 9 would fit snugly between Li and B. Beryllium metal was first prepared by F. Wishler in 1828 (the year he carried out his celebrated synthesis of urea from NH4CNO); he suggested the name by allusion to the mineral. The metal was independently isolated in the same year by A.B. Bussy using the same method  reduction of BeCl2 with metallic K. The first electrolytic preparation was by P. Lebeau in 1898 and the first commercial process (electrolysis of a fused mixture of BeF2 and BaF2) was devised by A. Stock and H. Goldschmidt in 1932.

Compounds of Mg and Ca, like those of their Group 1a neighbours Na and K, have been known from ancient times though nothing was known of their chemical nature until the seventeenth century. Magnesian stone was the name given to the soft white mineral steatite (otherwise called soapstone or talc) which was found in the Magnesia district of Thessally, whereas calcium derives from the Latin calx, calcis - lime. The Romans used a mortar prepared from sand and lime (obtained by heating limestone, CaCO3) because these lime mortars withstood the moist climate of Italy better than the Egyptian mortars based on partly dehydrated gypsum (CaSO4 - 2H20); these had been used, for example, in the Great Pyramid of Gizeh. The names of the elements themselves were coined by H. Davy in 1808 when he isolated Mg and Ca, along with Sr and Ba by an electrolytic method following work by J. J. Berzelius and M. M. Pontin: the moist earth (oxide) was mixed with one-third its weight of HgO on a platinum plate which served as anode; the cathode was a platinum wire dipping into a pool of mercury and electrolysis gave an amalgam from which the desired metal could be isolated by distilling off the mercury.

A mineral found in a lead mine near Strontian, Scotland, in 1787 was shown to be a compound of a new element by A. Crawford in 1790. This was confirmed by T. C. Hope the following year and he clearly distinguished the compounds of Ba, Sr and Ca, using amongst other things their characteristic flame colorations: Ba yellow-green, Sr bright red, Ca orange-red. Barium-containing minerals had been known since the seventeenth century but the complex process of unravelling the relation between them was not accomplished until the independent work of C. W. Scheele and J. G. Gahn between 1774 and 1779: heavy spar was found to be BaSO4 and called barite or barytes, whence Scheele's new base baryta (BaO) from which Davy isolated barium in 1808.

Radium, the last element in the group, was isolated in trace amounts as the chloride by Pierre and Marie Curie in 1898 after their historic processing of tonnes of pitchblende. It was named by Mariee Curie in allusion to its radioactivity; the element itself was isolated electrolytically via an amalgam by M. Curie and A. Debierne in 1910 and its compounds give a carmine-red flame test.

کل واکنشهای عنصر برلیم

BERYLLIUM Be

Light gray, floaty, fairly hard, brittle metal. On the air is covered by an oxide film. Passivated in cold water, concentrated sulfuric and nitric acids. Reducing agent. Reacts with boiling water, dilute acids, concentrated alkalis, non-metals, ammonia, metal oxides.

 

Obtaining metallic beryllium Be:

BeO + Mg = MgO + Be (700–800°C).

BeF2 + Mg = Be + MgF2 (700–750°C).

BeF2(liquid) → (Electrolysis) Be↓(cathode) + F2↑(anode).

BeCl2(liquid) → (Electrolysis) Be↓(cathode) + Cl2↑(anode).

 

Сhemical reactions with metallic beryllium Be:

2 Be + 3 H2O = BeO↓ +Be(OH)2↓ + 2 H2↑ (boiling).

Be + 2 HCl(diluted) = BeCl2 + H2↑.

3 Be + 8 HNO3(diluted, hot) = 3 Be(NO3)2 + 2 NO↑ + 4 H2O.

Be + 2 NaOH(conc.) + 2 H2O = Na2[Be(OH)4] + H2↑.

Be + 2 NaOH = Na2BeO2 + H2 (400–500°C).

2 Be + O2 = 2BeO (900°C, combustion on air).

Be + E2 = BeE2 (normal temp., E = F; 250°C, E = Cl; 480°C, E = Br, I).

Be + S = BeS (1150°C).

3 Be + N2 = Be3N2 (700–900°C),

2 Be + C(graphite) = Be2C (1700-1900°C, vacuum).

Be + 4 HF(conc.) = H2[BeF4] +H2↑.

3 Be + 2 NH3 = Be3N2 + 3 H2 (500–700° C).

Be + C2H2 = BeC2 + H2 (400–450°C).

Be + MgO = BeO + Mg (1075°C).

کل واکنشهای عنصر منگنز

MAGNESIUM - Mg

Properties of Magnesium:

The silver-white, relatively soft, ductile, malleable metal. On air covered with an oxide film. Passivated in cold water, concentrated sulfuric acid and hydrofluoric acid. Does not react with alkalis. A strong reducing agent, reacts with hot water, diluted acids, nonmetals. Transferred to solution by ammonium salts.

 

 

Molar mass	g/mol	24.305
Density	g/cm3	1.737
Melting point	°C	648
Boiling point	°C	1095

 

 

Methods for the preparation of Magnesium:

MgO + C(coke) = Mg + CO (over 2000°C).

MgO + Ca = CaO + Mg (1300°C).

Mg3N2 = 3Mg + N2 ( 700-1500°C).

MgCl2(melt) → Electrolysis → Mg(cathode) + Cl2↑(anode).

 

 

Сhemical reactions with Magnesium:

Mg + 2H2O(hot) = Mg(OH)2↓ + H2↑.

Mg + 2HCl(diluted) = MgCl2 + H2↑.

4Mg + 10HNO3(diluted) = 4Mg(NO3)2 + 2N2O↑ + 5H2O.

Mg + 2NH4Cl(conc., hot) = MgCl2 + 2NH3↑ + H2↑.

Mg + H2 = MgH2 (175°C, pressure, catalyst MgI2).

2Mg + O2 = 2MgO (600-650°C, combustion on air),

3Mg + N2 = Mg3N2 (780-800°C, combustion on air).

Mg + Cl2(moist) = MgCl2 (normal temp.).

Mg + H2S = MgS + H2 (500°C).

3Mg + 2NH3 = Mg3N2 + 3H2 (600-850°C).

Mg + 2N2O4 = Mg(NO3)2↓ + 2NO (150°C, vaccum, in ethyl acetate).

4Mg + SiO2 = Mg2Si + MgO (below 800°C, in the atmosphere of H2),

2Mg + SiO2 = Si + 2MgO (1000°C).

کل واکنشهای عنصر کلسیم

CALCIUM - Ca

Properties of Calcium:

Alkaline earth metal, silver-white, ductile, fairly hard. In humid air is covered with a film of oxide-hydroxide. Colors the flame of gas burner to brown-red. Reactive, reacts with oxygen, nitrogen, hydrogen, halogens and other non-metals by heating. A strong reducing agent, reacts with water, dilute acids, ammonia.

 

Molar mass	g/mol	40.078
Density	g/cm3	1.54
Melting point	°C	842
Boiling point	°C	1495

 

Methods for the preparation of Calcium:

CaH2 = Ca + H2 (over 1000°C).

4CaO + 2Al = 3Ca + (CaAl2)O4 (1200°C).

3CaCl2 + 2Al = 3Ca + 2AlCl3 (600-700°C).

CaCl2(liquid) → Electrolysis → Ca(on cathode) + Cl2↑(on anode).

 

Сhemical reactions with metal Calcium:

Ca + 2H2O = Ca(OH)2↓ + H2↑ (normal temp.),

2Ca + H2O(vapor) = CaO+CaH2 (200-300°C).

Ca + 2HCl(diluted) = CaCl2 + H2↑.

4Ca + 10HNO3(diluted) = 4Ca(MO3)2 + N2O↑ + 5H2O.

4Ca + 10HNO3(highly diluted) = 4Ca(NO3)2 + NH4NO3 + 3H2O.

Ca + H2 = CaH2 (500-700°C).

2Ca + O2 = 2CaO (over 300°C, burning on air).

Ca + E2 = CaE2 (normal temp., E = F; 200-400°C, E = Cl, Br, I).

Ca + S = CaS (150°C).

3Ca + N2 = Ca3N2 (200-450°C, burning on air),

3Ca + 2P(red) = Ca3P2 (350–450°C).

Ca + 2C(graphite) = CaC2 (550°C).

Ca + 6NH3(gas) = [Ca(NH3)6](e−)2(yellow) [normal temp.],

6Ca + 2NH3(gas) = Ca3N2 + 3CaH2 (600-650°C).

Ca + 6NH3(liquid) = [Ca(NH3)6](blue) [-40°C, in the atm. of Ar],

Ca + 2NH3(liquid) = Ca(NH2)2↓ + H2↑ (catalyst Pt).

کل واکنشهای عنصر استرانسیم

STRONTIUM - Sr

Properties of Strontium:

Alkaline earth metal. Light yellow, malleable. In the air is covered with oxide-nitride film. Colors the flame of a gas burner in a bright red color. High reactivity. Reacts with oxygen, nitrogen, hydrogen, halogen with heating. Strong reducing agent, oxidised by water, diluted acids, ammonia.

 

Molar mass	g/mol	87.62
Density	g/cm3	2.63
Melting point	°C	768
Boiling point	°C	1390

 

Methods for the preparation of Strontium:

4SrO+ 2Al = 3Sr + (SrAl2)O4 (1200°C).

3SrCl2 + 2Al = 3Sr + 2AlCl3 (600-700°C).

SrCl2(liquid)→ Electrolysis → Sr(on cathode) + Cl2↑(on anode).

 

Сhemical reactions with metal Strontium:

Sr + 2H2O = Sr(OH)2↓ + H2↑ (normail temp.),

2Sr + H2O(vapor) = SrO + SrH2 (200-300°C).

Sr + 2HCl(diluted) = SrCl2 + H2.

4Sr + 10HNO3(diluted) = 4Sr(NO3)2 +N2O↑ + 5H2O,

4Sr + 10HNO3(highly diluted) = 4Sr(NO3)2 + NH4NO3 + 3H2O.

Sr + H2 = SrH2 (200-500°C).

Sr + O2 = 2SrO (over 250°C, burning on air).

Sr + Cl2 = SrCl2 (200-400°C).

3Sr + N2 = Sr3N2 (450-500°C, burning on air).

Sr + 2C(graphite) = SrC2 (500°C).

6Sr + 2NH3(gas) = Sr3N2 + 3SrH2 (600-650°C).

Sr + 6NH3(liquid) = [Sr(MH3)6](blue) [-40°C, in the atm. of Ar],

Sr + 2NH3(liquid) = Sr(NH2)2↓ + H2 (catalyst Pt).

کل واکنشهای عنصر باریم

BARIUM - Ba

Properties of Barium:

Alkaline earth metal. Silvery-white, malleable, ductile. In the air is covered with a dark oxide-nitride film. Colors the flame of a gas burner in a yellow-green color. High reactivity. Reacts with oxygen, nitrogen, hydrogen, halogens and other non-metals. Strong reducing agent. Reacts with water, acids, hydrogen sulfide, ammonia.

 

Molar mass	g/mol	137.327
Density	g/cm3	3.60
Melting point	°C	727
Boiling point	°C	1860

 

Methods for the preparation of Barium:

BaH2 = Ba + H2 (over 675°C).

3BaO + Si = BaSiO3 + 2Ba (1200°C),

3BaO + 2Al = 2Ba + (BaAl2)O4 (1100-1200°C).

 

 

Сhemical reactions with metal Barium:

Ba + 2H2O = Ba(OH)2 + H2↑ (normal temp.).

Ba + 2HCl(diluted) = BaCl2 + H2↑.

4Ba + 10HNO3(diluted) = 4Ba(NO3)2 + N2O↑ + 5H2O,

4Ba + 10HNO3(highly diluted) = 4Ba(NO3)2 + NH4NO3 + 3H2O.

Ba + H2 = BaH2 (150-300°C).

3Ba + 2O2 = 2BaO + BaO2 (until 500°C, burning on air),

2Ba + O2 = 2BaO (over 800°C).

Ba + E2 = BaE2 (100-150◦ C; E = F, Cl, Br, I).

Ba + S = BaS (150°C).

3Ba + N2 = Ba3N2 (200-460°C, burning on air).

Ba + 2C(graphite) = BaC2 (500°C).

Ba + H2S = BaS+H2 (over 350°C).

6Ba + 2NH3(gas) = Ba3N2 + 3BaH2 (600-650°C).

Ba + 6NH3(liquid) = [Ba(NH3)6](blue) [-40°C, in the atm. of Ar],

Ba + 2NH3(liquid) = Ba(NH2)2 + H2 (catalyst Pt).

2Ba + 3CO2 = 2BaCO3 + C(graphite) [normal temp.].

کل واکنشهای عنصر رادیم

RADIUM Ra

Properties of Radium:

Alkaline earth metal. White, shiny, soft. Radioactive, the most long-lived isotope is Ra-226. The high reactivity. In the air is covered with a dark oxide-nitride film. Colors the flame of a gas burner in dark red color. A strong reducing agent. Reacts with water, acids, chlorine, sulfur. Milligram quantities of radium is isolated in the processing of uranium ore in the form of RaCl2. Produced by electrolysis of solution RaCl2 on a mercury cathode.

 

Molar mass	g/mol	226.025
Density	g/cm3	6.0
Melting point	°C	969
Boiling point	°C	1536

 

Сhemical reactions with metal Radium:

Ra + 2H2O = Ra(OH)2 + H2↑.

Ra + 2HCl(diluted) = RaCl2 + H2↑.

Ra + H2SO4 = RaSO4↓ + H2↑.

4Ra + 10HNO3(diluted) = 4Ra(NO3)2 + N2O + 5H2O.

2Ra + O2 = 2RaO (100°C, burning on air).

Ra + Cl2 = RaCl2 (normal temp.).

3Ra + N2 = Ra3N2 (100°C, burning on air).

Ra + S = RaS (150°C).

Ra + 2H2O + Na2CO3 = RaCO3↓ + H2↑ + 2NaOH.

معرفی عناصر گروه 1

Introduction group Ia

The alkali metals (Lithium Li, Natrium Na, Potassium K, Rubidium Rb, Cesium Cs, Francium Fr) form a homogeneous group of extremely reactive elements which illustrate well the similarities and trends to be expected from the periodic classification. Their physical and chemical properties are readily interpreted in terms of their simple electronic configuration, and for this reason they have been extensively studied by the full range of experimental and theoretical techniques.

Compounds of sodium and potassium have been known from ancient times and both elements are essential for animal life. They are also major items of trade, commerce and chemical industry.

Lithium was first recognized as a separate element at the beginning of the nineteenth century but did not assume major industrial importance until about 40 years ago.

Rubidium and cesium are of considerable academic interest but so far have few industrial applications.

Francium, the elusive element 87, has only fleeting existence in nature due to its very short radioactive half-life, and this delayed its discovery until 1939.

کل واکنشهای عنصر لیتیم

LITHIUM - Li

An alkali metal. A silvery-white. The easiest of the metals, soft, low melting. Reactive, air-oxide-nitride coated film. Ignited at a moderate heat, stains flame of a gas burner in a dark-red color. A strong reducing agent, reacts with water, acids, non-metals, ammonia.

 

Molar mass	g/mol	6.941
Density	g/cm3	0.534
Melting point	°C	180.5
Boiling point	°C	1336.6

 

Obtaining Lithium:

2LiH = 2Li + H2 (450°C).

2LiH (liquid) → Electrolysis → 2Li(on cathode) + H2↑(on anode).

2Li2O + Si = 4Li + SiO2 (1000°C).

Li2O + Mg = 2Li + MgO (t> 800°C).

3Li2O + 2Al = 6Li + Al2O3 (t> 1000°C).

2LiCl(liquid) → Electrolysis 2Li(on cathode) + Cl2↑(on anode).

2LiCl(melt) (on Hg-cathode) → Electrolysis → 2Li(on cathode) + Cl2↑(on anode).

2Li3N = 6Li + N2 (300-500°C, vacuum).

 

Reactions with Lithium:

2Li + 2H2O = 2LiOH + H2↑.

2Li + 2HCl(conc.) = 2LiCl + H2↑.

2Li + 3H2SO4(conc.) = 2LiHSO4 + SO2↑ + 2H2O.

3Li + 4HNO3(dilute) = 3LiNO3 + NO↑ + 2H2O.

2Li + H2 = 2LiH (500-700°C).

2Li + E2 = 2LiE (normal temp., E = F, Cl, Br; t>200°C, E = I).

4Li + O2 = 2Li2O (t>200°C, impurity Li2O2)

2Li + S = Li2S (t>130° C)

6Li + N2(moist) = 2Li3N (normal temp.)

6Li + N2 = 2Li3N (200-250°C, pressure)

2Li + 2C = Li2C2 (t>200°C, vacuum)

4Li + Si = Li4Si (600-700°C, impurity Li2Si)

2Li + 2NH3 = 2LiNH2 + H2 (220°C)

2Li + NH3 = Li2NH + H2 (400°C)

Li + 4NH3 (liquid) = [Li(NH3)4]0 (blue) [-40°C]

[Li(NH3)4]⁰ + nNH3(liquid) ↔ [Li(NH3)4]⁺ + e^- + nNH3.

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کل واکنشهای عنصر سدیم

NATRIUM (SODIUM) - Na

Propertiesof sodium Na:

An alkali metal. A silvery-white (in a thin layer - with a purple tint), lightweight, very soft, low melting. Dark red sodium vapor is composed of atoms Na (predominant) and molecules Na2. Under special conditions is formed a violet-blue colloidal solution of sodium in ether. Chemically dissolves in liquid NH3 (blue solution), molten NaOH. Very reactive, air-covered oxide film (tarnish), ignited at a moderate heat. Stable in an atmosphere of argon and nitrogen. A strong reducing agent, reacts vigorously with water, acids, non-metals. With nitrogen only reacts when heated (as opposed to Li). With the mercury forms an amalgam, amalgam - a strong reducing agent, but (in contrast to the pure sodium), the reaction with water flowing quietly.

Does not react with air, is well preserved under a layer of gasoline or kerosene, is easily suspended in boiling inert solvents (toluene, xylene, heptane, octane, etc.). In an inert atmosphere, the molten sodium is rapidly distributed over the surface of certain solids (NaCl, Na2CO3, coal, iron, Al2O3, SiC, ZrO2), forming a gray-black monatomic coverage. Сolors the flame of a gas burner in yellow. The most widespread metal in seawater.

 

Molar mass	g/mol	22.99
Density in solid state	g/cm3	0.968
Melting point	°C	97.83
Boiling point	°C	886

 

Obtaining sodium Na:

4 NaOH (liquid) →Electrolysis 4Na (cathode) + O2↑ (anode) + 2H2O.

Na2CO3 + 2C(coke) = 2Na + 3CO (900-1000°C).

2NaCl (liquid) →Electrolysis 2Na (cathode) + Сl2↑ (anode).

2NaCl + 2H2O →Electrolysis→ H2↑ (cathode) + Сl2↑ (anode) + 2NaOH.

2NaCl (melt) (on Hg-cathode)→Electrolysis 2Na (cathode) + Сl2↑ (anode).

 

Reactions with sodium Na:

2Na + 2 H2O = 2 NaOH + H2↑.

2Na + 2 HСl (diluted) = 2 NaCl + H2↑.

2Na + 2 NaOH = 2 Na2O + H2 (600°C).

2Na + H2 = 2 NaH (250-400°C, pressure).

2Na + O2 (air) = Na2O2 (burning, impurity Na2O)

2Na + O2 = Na2O2 (250-400°C).

4Na + O2 + 2 H2O = 4 NaOH.

2Na + E2 = 2 NaE (normal temp., E = F, Cl; 150-250°C, E = Вr, I).

2Na + E = Na2E (t>130° С, E = S, Se, Те),

2Na + nS = Na2(Sn) [-40° С, in the liquid NH3, n = 1, 2, 4, 5].

6Na + N2 = 2 Na3N (100° С, electric discharge),

3Na + P (red) = Na3P (green) [200°C, in the atmosphere of Аr].

2Na + 2C (graphite) = Na2C2 (150-200°C).

2Na + 2H2S (saturated) = 2NaHS↓ + H2↑ (in benzene).

Na + 4NH3(liquid) = [Na(NH3)4] (blue) (at - 40°C),

[Na(NH3)4] + nNH3(liquid) ↔ [Na(NH3)4]+ + e- - nNH3.

2Na + 2NH3(gas) = 2NaNH2 + H2 (350°C).

2Na + B2O3 + 7H2 = 2Na[BH4] + 3H2O (250-300°C).

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کل واکنشهای عنصر پتاسیم

POTASSIUM - K

Properties of Potassium :

An alkali metal. A silvery-white (in a thin layer - with a purple tinge), soft, low melting. Blue-green pairs of potassium consist from atoms of K (mainly) and K2 molecules. Chemically dissolves in liquid ammonia (dark blue solution), in the molten hydroxide potassium. It is extremely reactive, the strongest reducing agent, reacts with O2 of air, water (there is ignition of the evolved hydrogen), dilute acids, non-metals, ammonia, hydrogen sulfide. Practically does not react with nitrogen (in contrast to Li and Na). It is well preserved under a layer of gasoline or kerosene. With the mercury forms an amalgam. Do not fused with Li, Mg, Zn, Cd, Al and Ga. Forms intermetallic compounds with Na, Tl, Sn, Pb and Bi. Colors the flame of a gas burner in a purple color. The fifth element of the spread in nature.

 

Molar mass	g/mol	39.098
Density in solid state	g/cm3	0.8629
Melting point	°C	63.51
Boiling point	°C	760

 

Obtaining metallic potassium:

2 KH = 2K + H2 (400° C, vacuum).

4 KOH (liquid) → Electrolysis 4 K (cathode) + O2 ↑(anode) + 2 H2O.

2 KCl (liquid) → Electrolysis 2 K (cathode) + Cl2 ↑(anode).

2 KCl + 2H2O → Electrolysis H2↑(cathode) + Cl2 ↑(anode) + 2KOH

2 KCl (melt) (on Hg-cathode)→ Electrolysis  2 K(cathode) + Cl2 ↑(anode).

 

Chemical reactions with metallic potassium:

2 K + 2H2O = 2 KOH + H2↑.

2 K + 2 HCl (diluted) = 2KCl + H2↑.

8 K + 6 H2SO4 (diluted) = 4 K2SO4 + SO2 + S ↓ + 6 H2O (impurityH2S)

21 K + 26 HNO3 (diluted) = 21 KNO3 + NO ↑ + N2O ↑ + N2 ↑ + 13H2O.

2 K + 2 KOH = 2K2O + H2 (450° C).

2 K + H2 = 2 KH (200—350° C).

K + O2 (air) = KO2 (burning, impurity of K2O2)

K →( O2 ) K2O2↓ →O2, (time) KO2↓ (- 50° C, in the liquid NH3).

4 K + O2 + 2 H2O = 4KOH.

2 K + E2 = 2 KE (normal temp.; E = F, Cl, Вr, I).

2 K + E = K2E (100—200°C ;E = S, Se, ТE).

З K + Р (red) = K3P (green) [200° C, in the atmosphere Аr].

2 K + 2 H2S (saturated)= 2 KHS ↓ + H2↑ (in benzene).

2 K + 2 NH3 (gas) = 2KNH2 + H2 (65—105° C).

K + 6 NH3 (liquid) = [K(NH3)6] (dark-blue) [- 50° C]

[K(NH3)6] + n NH3 (liquid) ↔ [K(NH3)6]+ + E-nNH3.

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کل واکنشهای عنصر روبیدیم

RUBIDIUM - Rb

Properties Rubidium Rb:

An alkali metal. White, soft, very low melting. The vapor of rubidium is painted in a greenish-blue color. Chemically dissolves in liquid NH3 (dark-blue solution), in the melt RbOH. Extremely reactive, strong reducing agent. Vigorously reacts with oxigen of air, water (there is inflammation of the metal and released hydrogen), dilute acids, non-metals, ammonia, hydrogen sulfide. Does not react with nitrogen. It is well preserved only under a layer of paraffin oil or vaseline. With the mercury forms an amalgam. Colors the flame of a gas burner in a purple color.

 

Obtaining rubidium Rb:

2RbH = 2Rb + H2 (t > 200° С).

2Rb2O = Rb2O2 + 2 Rb (400-550° С).

RbOH (liquid) → Electrolysis → 4Rb (cathode) + O2↑(anode) + 2 H2O.

2RbCl (liquid) → Electrolysis → 2Rb (cathode) + Cl2↑(anode).

2RbCl + 2 H2O → Electrolysis → H2↑ (cathode) + Cl2↑(anode) + 2 RbOH,

2RbCl(melt)(on Hg-cathode) → Electrolysis → 2Rb (cathode) + Cl2↑ (anode).

 

Сhemical reactions with rubidium Rb:

2Rb + 2H2O = 2 RbOH + H2↑.

2Rb + 2HCl (diluted) = 2RbCl + H2↑.

8Rb + 6H2SO4 (diluted, cold) = 4 Rb2SO4 + SO2 + S↓ + 6 H2O (impurity H2S).

21Rb + 26HNO3 (diluted, cold) = 21 RbNO3 + NO↑ + N2O↑ + N2↑ + 13 H2O.

2Rb + 2RbOH = 2 Rb2O + H2 (400° С).

2Rb + H2 = 2 RbH (300-350° С, pressure).

Rb + O2 (air) = RbO2 (burning).

4Rb + O2 = 2Rb2O (in the cold),

Rb → + O2 → Rb2O2 → + O2 → time → RbO2↓(-50° С, in the liquid NH3).

4Rb + O2 + 2 H2O = 4 RbOH (normal temp.).

2Rb + E2 = 2RbE (normal temp.; E = F, Cl, Br, I).

2Rb + S = Rb2S (100-130° С).

2Rb + 2 H2S (saturated) = 2 RbHS ↓ + H2↑ (in benzene).

2Rb + 2 NH3 (gas) = 2 RbNH2 + H2 (40-60° С).

Rb + 6 NH3 (liquid) = [Rb(NH3)6]      (-40° С),

[Rb(NH3)6] + nNH3 (liquid) ↔ [Rb(NH3)6]⁺ + e^-·nNH3.

4Rb + 3SiO2 = Rb2SiO3 + Si (t > 300° C).

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کل واکنشهای عنصر سزیم

CESIUM - Cs

Properties of metallic cesium Cs:

An alkali metal. White (on the cut - light yellow), soft, very low melting. The vapor of cesium is painted in a greenish-blue color. Chemically dissolves in liquid ammonia (dark blue solution), the melt of CsOH. It is extremely reactive, the strongest reducing agent, reacts with oxygen in the air, water (there is inflammation of the metal and hydrogen is released), dilute acids, non-metals, ammonia, hydrogen sulfide. Does not react with nitrogen. It is well preserved only under a layer of paraffin oil or vaseline. With the mercury forms an amalgam. Colors the flame of a gas burner in the blue.

 

Obtaining metallic cesium Cs:

2 Cs2O = Cs2O2 + 2 Cs (300—500° C).

4 CsOH (liquid) → Electrolysis 4 Cs (cathode) + O2↑ (anode) + 2 H2O.

2 CsCl + 2H2O → Electrolysis H2↑ (cathode) + Cl2↑ (anode) + 2 CsOH,

2 CsCl (melt) (on Hg-cathode)→ Electrolysis 2 Cs (cathode) + Cl2↑(anode).

 

Сhemical reactions with metallic cesium Cs:

2 Cs + 2 H2O = 2CsOH + H2↑.

2 Cs + 2 HCl (diluted) = 2 CsCl + H2↑.

8 Cs + 6 H2SO4 (diluted, cold) = 4 Cs2SO4 + SO2 + S↓ + 6H2O (impurity H2S),

21 Cs + 26 HNO3 (diluted, cold) = 21 CsNO3 + NO↑ + N2O↑ + N2↑ + 13H2O.

2 Cs + 2 CsOH = 2 Cs2O + H2↑ (300—350° C).

2 Cs + H2 = 2 CsH (300—350° C, pressure).

Cs + O2(air) = CsO2 (combustion),

4 Cs + O2 = 2 Cs2O (in the cold), Cs →(O2) Cs2O2↓ →(O2), (time) CsO2↓ (-50° C, in the liquid NH3).

4 Cs + O2 + 2 H2O = 4 CsOH.

2 Cs + E2 = 2 CsE (normal temp.; E = F, Cl, Br, I).

2 Cs + S = Cs2S (100—130° C).

2 Cs + 2 H2S (saturated) = 2 CsHS ↓ + H2 ↑ (in benzene).

2 Cs + 2 NH3 (gas) = 2 CsNH2 + H2 (30—45° C).

Cs + 6 NH3 (liquid) = [Cs(NH3)6](time, blue) [-40° C],

[Cs(NH3)6] + nNH3 (liquid) ↔ [Cs(NH3)6]+ + e-·nNH3.

4 Cs + 3 SiO2 = 2 Cs2SiO3 + Si (t>300° C).

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کل واکنشهای عنصر فرانسیم

FRANCIUM Fr

Alkali metal. White, very fusible. Radioactive, most lived isotope 223Fr (half-life of 22 min.). The most reactive of all metals, chemical behavior similar to cesium. On the air is covered by an oxide film. A strong reducing agent, reacts vigorously with water, and the acid releasing hydrogen. Cation of Fr+ in the aqueous solution is colorless. In the Earth's lithosphere is formed by the radioactive decay of uranium and actinium. Synthesized by bombardment of uranium nuclei by protons or neutrons nuclei of radium. Isolated compounds FrClO4 and Fr2[PtCl6] by method of coprecipitation with appropriate poorly soluble salts of Rb and Cs.