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Hydrogen compounds of group IV elements The IVth group of the periodic system contains the elements: C, Si, Ge, Sn, and Pb. Carbon. Man has known carbon, as charcoal from the combustion of wood, from prehistoric times. Its name derives from the Latin “carbo” meaning charcoal. The element is unique in the vast number of its compounds- there are probably more than one million known today. Carbon is widely distributed throughout the universe. Much of the energy of the sun and stars is due to the carbon cycle. Both black and transparent diamonds have been found in meteorites. The cosmic abundance is six times that of silicon. The name of this element derives from a corruption of the Greek word “adamas” meaning invincible. Combined with hydrogen, carbon occurs extensively as coal, petroleum and natural gas. In spite of its wide spread occurrence carbon still only constitutes 0.027% of the earth’s crust. The distribution of terrestrial carbon is approximately as follows: minerals-99.7%, atmosphere-0.2%, and living matter-0.01%. Even if combining with hydrogen is very difficult to achieve, a great number of hydrides of carbon is known which is the object of organic chemistry. The hydrides of carbon are covalent compounds. Carbon possesses in the valence shell four electrons and forms four bonds in almost all of its combinations. The bonds of carbon aren’t made with the pure orbitals 2s2p, but with some hybrid tetrahedral orbitals. The saturated hydrides of carbon don’t have an acid character, which means that don’t have the tendency to loose protons. In a period this tendency increases from the 4th to the 7th group. So in the series: H H H H:C:H H:N:H :O:H :F:H H HF gives easily the proton to water and forms ions like F¯ and H 3O+. Ammonia and methane don’t loose protons. The acid or basic properties of some combinations can be explained by the presence in its molecules of lone pairs, able to form coordinate bonds. The hydrides of carbon don’t have lone pairs, so they are neutral. The tendency to form covalent bonds and the neutral character allows carbon atoms to bond between themselves and to form chains of various forms. Regarding its health hazard carbon is itself relatively non-toxic. The radiological hazards of using radioactive carbon compounds have already been discovered. For 14C, the most widely used radioisotope, the tolerance doses are: Whole body burden 300µC Maximum permissible concentration in drinking water 8x10 -3µC ml-1 Maximum permissible concentration in air 4x10 -6µC ml-1 Silicon. 1 It is not possible to say when or by whom silicon and its early compounds were discovered, because man has used natural silica and silicates since dawn of the race. The name silica (and from it, the words silicon, silicide, silicate, and silicone by the usual etymology) comes from the Latin word silex, silicis for flint. Silica was found to be a relatively strong acid earth at high temperature, capable of forming salts of all the alkali and alkaline-earth metals and of most transition metals as well. In this way some understanding of the ancient ceramic arts became possible, for all involved silica and silicates. Those arts themselves were well developed on an empirical basis long, long before the alchemists; pottery in various forms was made in the Near East 6000 years ago, and elaborate techniques were used in China 5000 years ago. Vitreous enamels, “glass paste”, and silicate glazes in exquisite colors and superb workmanship were found in the tomb of Tutankhamen,1350 B.C. Silicon is found in free state or in anionic state in every natural water. In these water the quantity of silicon can vary between 1 and 40 ppm. The role of silicon on human body (which contains 7 grams of silicon) is not very well studied till now but it is admitted that the water shouldn’t contain more than 20 mg SiO 2/l. The overall picture of silicon in relation to biological systems is that of a ubiquitous, passive, and almost always benign element. It exhibits no general toxicity, and except for the special situation called silicosis, it does not interfere with the functioning of living organisms. On the positive side, silica and silicates appear to be beneficial in plant nutrition in that they increase the uptake of phosphorus from phosphate fertilizers, and of course potassium silicate is preferred as such an auxiliary because of its potassium content. The first who tried to obtain silicon was Van Helmot. Others who tried the same thing are: Otto Tache-Taehnetius (1644), Lavoisier, and Humphry Davy. In 1789 Sheele described the method of obtaining silicon tetra fluoride. Berzelius reduced silica, mixing it with Fe2S and charcoal at high temperatures, in 18101823. From ironsilica with hydrochloric acid he separated silicon. Thénard and Gay-Lussac discovered amorphous silicon. Saturated hydrates of silicon have the following formula: Si nH2n+2. These compounds are named silanes. They contain only simple bonds. There weren’t discovered compounds with double or triple Si-Si bonds. For silicon are known only terms with n=6 or less. Silanes are obtained through the action of hydrochloric acid (HCl), sulphuric acid (H2SO4) or phosphoric acid (H3PO4) on Mg2Si, in the absence of air or in different solvents (liquid NH3). The products of this reaction, besides hydrogen and monosilane, are: disilane Si2H6, trisilane Si3H8, tetrasilane Si4H10, pentasilane Si5H12 and hexasilane Si6H14. Mg2Si + 4 HCl 2 MgCl2 + SiH4 A general method for obtaining silanes involves the reduction of the corresponding chlorosilanes by lithium aluminium hydride or borohydride. Not only the simple binary silicon chlorides but also literally thousands of arganochloro-silanes and other halosilanes may similarly be reduced: 2 SiBr4 SiHCl3 (CH3)2SiCl2 C6H5SiHCl2 (C2H5)3SiBr SiH4 SiH4 (CH3)2SiH2 C6H5SiH3 (C2H5)3SiH LiAlH4 Et2O An economical way of reduction is with sodium hydride dissolved in an eutectic bath of melted lithium and potassium chloride: LiCl-KCl SiCl4 + 4 NaH 4 NaCl + SiH4 348oC Another method for establishing Si-H bonds involves the reaction of a hydrogen halide with elementary silicon or an alloy such as ferrosilicon: 350oC 3 HCl + Si SiHCl3 + H2 300oC 2 CH3Cl + Si CH3SiHCl2 + H2 + C Cu Another method of obtaining silane (in small quantity) is an American method, from plasma. In general, the plasma is a gaseous mixture of electrons, positive ions, photons, atoms, excited or unexcited molecules. If in the hydrogen plasma which leaves the anodic space of the generator is introduced SiCl4(gas) is obtained a mixture of chlorosilanes. If the quantity of hydrogen is big enough are obtained SiH2Cl2, SiH3Cl and silane, SiH4. But this method has a very small yield in silane. The principal product is SiH3Cl. Mono and disilane are colourless gases and the other terms are liquids. The smell of monosilane is like the smell of phosphorated hydrogen and arseniated hydrogen. Monosilane has a tetraedrical structure in which the distance Si-H is 1.55 A. The silanes are less stable than the paraffin hydrocarbons and when heated they decompose at a rate which increases with the complexity of the molecule. The comparative bond energies are 99 kcal/mole for C-H in CH4 and 76kcal/mole for Si-H in SiH4: 2 SiH4 = 2 SiH2 + 2 H2 = Si2 + 4 H2 In terms of chemical reactivity, the chief distinction between silanes and the normal alkanes is the readiness with which the silanes oxidise. They are oxidised even by methanol: SiH4 + 2 CH3OH = SiH2(OCH3)2 + 2 H2 SiH4 + 4 CH3OH = Si9(OCH4)4 + 4 H2 Si-H bonds have an ionic character Si+H- and react in consequence with halogens or halogen acids giving substitution products: 3 SiH4 + Br2 = SiH3Br + HBr SiH4 + HCl = SiH3Cl + H2 In the presence of chlorine and bromine, monosilane explodes. The silanes are strong reducing agents. They reduce Cu 2+ in aqueous solution to copper hydride, Hg+ to Hg, KMnO4 to MnO2 and Ag+ to Ag. Such reduction action leads to explosive reaction with chlorine or bromine, although the violence can be abated at low temperature. Transfer of hydrogen to carbon, with simultaneous transfer of chlorine to silicon, is a reaction discovered by Stock; it illustrates the high affinity of silicon for chlorine and the greater bond energy for C-H over Si-H. AlCl3 Si3H8 + 4 CH4Cl3 Si3H4Cl4 + 4 CH2Cl2 Or AlCl3 Si3H8 + 5 CHCl3 Si3H3Cl5 + 5 CH2Cl2 The reaction of alkali metals, their hydrides and their alkyls also are very interesting. These can be considered as metatheses, with hydride ion transferred: Si2H6 + KH = SiH4 + KSiH3 SiH4 + C2H5Li = LiH + C2H5-SiH3 SiH4 + 4C6H5Li = 4LiH + (C6H5)4Si The first reaction must be carried out in 1,2- dimethoxy-ethane to obtain a satisfactory yield, and the second in diethyl ether. If the familiar liquid alloy of sodium of potassium is used, also in 1,2- dimethoxy-ethane the potassium silyl can be obtained: Si2H6 + 2 NaK = 2 KSiH3 + 2 Na The potassium silyl so formed is a colorless crystalline substance with NaCl structure, which decomposes above 200oC to liberate hydrogen: KSiH3 + H2O = SiH4 + KOH And immediately: KOH SiH4 + 4 H2O Si(OH)4 + 4H2 The rapid and complete hydrolysis of the silanes in aqueous alkalis allows their quantitative determination by collecting and measuring the evolved by hydrogen. 4 Through action of the metals over silicon-chloroform or silicon-bromoform it is obtained the polymer (SiH)n named polysiline: SiHCl3 + 3 Na = 3 NaCl + 1/n (SiH)n 2 SiHBr3 + 3 Mg = 3 MgBr2 + 2/n (SiH)n An electric arc between silicon electrodes in a stream of hydrogen at 60 mm pressure gives rise to a monohydride of silicon that can be recognized from its band spectrum. An absorption band for the compound appears in the spectrum of the sun. Superior silanes are thermally decomposed to some solid products, which are polymers: Si5H12 = 2/n (SiH)n + Si2H6 + SiH4 Polysiline has a macromolecular structure: H H Si Si Si H H Si Si Si H H The action of gaseous hydrochloric acid, in alcohol, over SiCa determines the formation of another polymer, polysilene (SiH2)n. n SiCa + 2n HCl = (SiH2)n + n CaCl2 The structure of this compound is: H H H H SiSiSiSi H H H H (SiH2)n is a brown compound, which explodes spontaneously in air and is decomposed by mineral acids, liberating hydrogen. The compounds, which contain silicon, hydrogen and halogens, are named halogenosilanes. The following types of halogenosilanes are known: SiHX 3 (X= F, Cl, Br, I), SiH2X2 (X= F, Cl, Br, I), SiH3X (X= F, Cl, Br, I). O.Ruff prepared halogentrifluorosilane through action of SiHCl3 over SnF4, TiF4 or SbF3: 3 SnF4 + 4 SiHCl3 = 4 SiHF3 + 3 SnCl4 5 Other compounds of silicon are the deuterides. SiD 4 can be made by the same original method as SiH4 except that DCl and D2O are used for decomposing the magnesium silicide. Monosilane is used industrially as a source of hyper pure silicon for semiconductor applications, but there is no substantial use for the higher hydrides. Silicon in high purity is used in solar batteries, photo elements that supply the spaceships with electric current, for making diodes, triodes, for deoxidizing of copper alloys. Regarding the health hazard, the silicon halides all hydrolyze to produce hydrogen halides, which of course are irritating and corrosive when inhaled or ingested. Hence all compounds in which one or more halogen atoms are linked to silicon must be regarded as hazardous. Silicon hydrides have received the rating 3 (after Sax rates all common chemical substances may have: 0 (no hazard), 2 (moderate hazard), 3 (high hazard) and U (unknown hazard)) being considered toxic in the same way as boron hydrides; they also carry a fire hazard, of course. Germanium. In 1871, Mendeleev predicted the existence of germanium putting it under silicon with the name of “eka-silicon”. This element was named by his discoverer, Clemens Winkler (1838-1904), teacher at Freiberg University, in the honor of his country. Germanium is a ubiquitous component of living organisms, but it has no structural function. Sax reports that germanium has a low order of toxicity. Its compounds in general are much less poisonous than those of lead and tin, but GeH4 has a hemolytic effect and is dangerous at levels about 100 ppm. On the average, each one of us ingests 1500 µg of germanium daily in our food, so we already have in our tissues an appreciable level of this element for which no function or decisive effect is known. The hydrogen compounds of germanium correspond to the formula: GenH2n+2. Are known the following terms: (GeH)x, (GeH2)x, GeH4, Ge2H6, Ge3H8, Ge4H10, …, Ge9H20. The bond Ge-H is less stable than C-H bond. These compounds of germanium easy undergo in hydrohalides or in halides of germanium under the direct action of halogens. The compounds of hydrogen with germanium are less inflammable than those of silicon, but the heavier that are the easiest are oxidized to GeO 2 and H2O, and are more stable to hydrolysis than silanes. Monogermane GeH4. If a solution containing GeCl4 and H2SO4 is reduced, is obtained a gaseous mixture of hydrogen and a compound of germanium with hydrogen. If we introduce this mixture in a solution of sodium nitrate we obtain a black combination with the formula GeAg4. GeH4 + 4 AgNO3 = GeAg4 + 4 HNO3 Through heating with liquid air, from the mixture can be separated GeH 4. Through thermal dissociation is formed a mirror of germanium: 6 GeH4 = Ge + 2 H2 Monogermane was made by the action of acids on an alloy of magnesium with germanium. If aqueous acids are used, large quantities of hydrogen are produced and must be separated from the mixture of germanium hydrides; furthermore, the yield of hydrides based on germanium is disappointing. Ammonia system acids are more satisfactory: the action of NH4Br on powdered Mg-Ge alloy (3 parts Ge + 2 parts Mg fused at 800oC) suspended in liquid ammonia converts 60% to 70% of the Ge to various hydrides, which are separated by distillation in a vacuum chain. GeMg2 + 2 H2SO4 = GeH4 + 2 MgSO4 Some GeH4 also is obtained from the thermal decomposition of solid GeH, and also by the electrolysis of a solution of GeO2 in H2SO4. More modern and practical methods consist in the reduction of GeCl 4 by LiAlH4 in ether, and the still simpler and easier reduction of GeO 2 by NaBH4 in water solution. By the last method, GeO2 is dissolved in HBr and an excess of NaBH4 solution (5 g in 100 ml H2O) is added dropwise. The GeH4 (with about 1% Ge2H6) is condensed from the stream of hydrogen in a series of five traps cooled to –196oC; the yield is 98%. At ordinary temperature, GeH4 is a colourless gas, which burns in air with blue flame. It can be turned into a liquid at –900C and into a solid at –164.80C. Monogermane reacts with sodium dissolved in liquid ammonia to form NaGeH3, the same product obtained by splitting Ge 2H6 with sodium in the same medium. Treatment of the NaGeH3 with NH4Br in the same solvent evolves pure GeH4. 2 GeH4 + 2 Na = 2 Na2GeH3 + H2 In contrast to SiH4 and SnH4, GeH4 does not ignite when it meets air. It can be mixed with pure oxygen, at low pressure, and the oxidation begins slowly at 320oC. In the range 230oC to 330oC the reaction produces water and white deposit of GeO2, but at higher temperatures explosions may result and brown germanium is deposited as the thermal dissociation of GeH 4 precedes the oxidation of its products. In sharp distinction to SiH4 (which is decomposed rapidly by even exceedingly dilute alkali solutions), GeH4 is unaffected by 30% aqueous AgNO3 to evolve hydrogen and precipitate a black mixture of Ag and Ge. Oxidising agents convert it to GeO2 and water. It dissolves in liquid ammonia to form a conducting solution believed to contain NH4+ and GeH3- ions, and the solution dissolves P4 to form an ammonium phosphogermanide. Sodium and potassium in liquid ammonia convert GeH4 to NaGeH3 and KGeH3, white solids that are unstable at room temperature and decompose to the metal germanide and hydrogen. The substances MGeH3 react in liquid ammonia with CH3Cl to form CH3GeH3 but the same substances react with CH2Br2 in a reductive way, not to yield H3GeCH2GeH3: CH2Br2 + 2 NaGeH3 + NH3 7 CH3GeH3 + GeH3NH3 + 2 NaBr In the same solvent NaGeH3 reduces aromatic halides to hydrocarbons, rather than forming RGeH2: NaGeH3 + C6H5Br C6H6 + GeH2 + NaBr Digermane Ge2H6. This is obtained as a by-product in the preparation of GeH4 as described above, and also by the circulation of GeH4 through a silent electrical discharge at low pressures. Digermane decomposes at 200 oC in a manner that fits the scheme: GeH6 2 GeH3 GeH3 + Ge2H6 GeH4 + Ge2H5 Ge2H5 GeH2 + GeH3 GeH2 Ge + H2 2 GeH2 GeH4 + Ge At ordinary temperature is a liquid with boiling point 28.5 0C. Its melting point is –1090C. Through thermal decomposition is obtained monogermane, germanium, hydrogen and the polymer (GeH2)n. Trigermane Ge3H8. Is separated from the same mixture from which is obtained Ge2H6. It is a mobile colourless liquid, with boiling point –105.60C. When exposed to air, Ge3H8 soon changes to a white solid, but does not ignite. It does not dissolve in water, but appears to be oxidized by air in the water. Like Ge2H6, Ge3H8 dissolves in CCl4 and reacts with it, apparently to form GeCl4. Tetra and pentagermanium are obtained by fractionating the residue obtained after separation of inferior homologues. (GeH)x is obtained by treating the alloy GeNa with water(L.M.Dennis & N.A.Skow). x GeNa + x H2O = (GeH)x + x NaOH This substance is a dark brown powder, which explodes when is dried. Probably it has the following structure: H H │ │ │ │ ― Ge ― Ge ― Ge ― Ge ― │ │ H H H H │ │ ― Ge ― Ge ― Ge ― Ge ― │ │ │ │ H H 8 (GeH2)x was obtained by Glarum and Kraus in 1950 – 1952), through action of NaGeH3 over phenilbromide: NaGeH3 + Ca6H5Br = (GeH2)x + x CaCl2 It is an amorphous yellow powder, stable without having contact with air. It’s great stability and insolubility in organic solvents is explained by his chain structure: H H H H GeGeGeGe H H H H This compound reacts with alkali-hydroxides: KOH (50%) (GeH2)x K2GeO2 + H2 + GeH4 80˚C Some halogen derivatives of the compounds of hydrogen with germanium are: GeH3Cl, GeH2Cl2, GeHCl3, GeH3Br, GeH2Br2, GeHBr3, GeHI3. Some properties of the Ge (lV) hydrides are shown in the following table: Property Melting point Boiling point Density at m.p. Ge-H bond dist. Ge-Ge bond dist. Critical temperature, K Critical pressure GeH4 -164.8 -88.1 1.52 1.527 308 54.8 Ge2H6 -109 29 1.98 Ge3H8 -105.6 110.5 2.20 2.41 483 45.7 2.41 588 37.9 Ge4H10 Ge5H12 176.9 234 665 32.5 730 28.8 Germanium is used in radio technique, radar, as semiconductor and for obtaining special alloys. Tin In the book of Numbers in the Bible’s Old Testament, tin is mentioned as a metal of value under the name bedil. The ancient Indian author Veda refers to tin as “trapu”. Objects made of tin have been found in the tombs of ancient Egypt. Caesar, recording the presence of tin in Britain, referred to it as “plumbum album” as also did Pliny, to distinguish it from lead, which was plumbum nigrum. Tin is found in nature almost exclusively as the tin (lV) oxide as cassiterite or tinstone. Tin is estimated to be present in the earth’s crust as 4x10 -3% by weight, and to be in seawater at a concentration of 0.003 g/ton. The two binary hydrides of tin are known to date are SnH4 and Sn2H6. SnH4 was originally prepared in very low yield by the treatment of tin-magnesium alloy with dilute acids. The lithium aluminium hydride reduction of tin (lV) chloride has made SnH4 easy available; in the presence of a trace of oxygen this reaction gives 80-90% yields. In addition to a good yield of SnH4, the potassium 9 borohydride reduction of stannite ion in solution produced the hydride distannane Sn2H6. All physical data on SnH4 and SnD4 are in full accord with a regular tetrahedral structure. Stannane undergoes appreciable decomposition to tin and hydrogen even at ordinary temperatures, but about 100 oC it is very rapidly decomposed. The decomposition is first order with respect to stannane. Somewhat remarkably, oxygen inhibits decomposition at pressures above 1 mm, and stannane may be stored mixed with oxygen at room temperature. Stannane is toxic and is intermediate between silane and germane in its chemical properties; it is unattached by dilute acids and alkalis, but is decomposed by concentrated acids or alkali. It is a powerful reducing agent, and is rapidly decomposed by solutions of transitional metal salts. Stannane and hydrogen chloride evolve hydrogen to form stannylchloride, H 3SnCl, which is very unstable. Sodium in liquid ammonia may be titrated against SnH 4 to produce H3SnNa and H2SnNa2. It would appear, however, that these two sodium derivatives are stable as amines, as removal of ammonia, even at –63.3oC, causes decomposition. Stannane reduces nitrobenzene in 94% yield and reduces benzaldehyde to benzyl alcohol virtually quantitatively. Stannane reacts with boron trifluoride to form tin (lV) fluoride. Any extensive study of the chemistry of the tin-hydrogen bonds is extremely difficult in SnH4 due to its inherent instability. In the case of organotin hydrides, this difficulty is considerably allayed. The stability of the organotin hydrides increases with decreasing number of tin-hydrogen linkages in the molecule. The organotin hydrides are almost invariably synthesized by the reduction of the corresponding organotin halide. CH3SnCl3 + LiAlH4 CH3SnH3 (C2H5)2SnCl2 + LiAlH4 (C2H5)2SnH2 (C6H5)3SnCl + LiAlH4 (C6H5)3SnH (CH3)3SnNa + NH4Br (CH3)3SnH (C4H9)3SnCl (C4H9)3SnH Examples of tin hydrides are known containing two or more tin atoms: (C4H9)2Sn―Sn(C4H9)2 + LiAlH4 │ │ Cl Cl (C 4H9)2Sn―Sn(C4H9)2 │ │ H H [(C4H9)2Sn]n + (C4H9)2SnH2 Lead The very easy extraction of lead from its ores made it one of the few metals used exclusively from earliest times. Lead was in common usage in Ancient Egypt 10 for ornamental objects and solder, and lead salts were used to glaze pottery. The Hanging Gardens of Babylon was floored with sheet lead as a moisture retainer, and the Babylonians and other ancients used lead for caulking and for fattening of iron bolts and hooks in bridges, houses and other stone buildings. Over four centuries, it is estimated, the Roman Empire extracted and used six to eight million tons of lead, with a peak annual production of around sixty thousands tons. Lead hydride is the least well characterized of the group lV hydrides. It is formed along with hydrogen, on the electrolysis of dilute sulphuric acid with lead electrodes, and by dissolution of lead-magnesium alloy in dilute acid. The hydride of lead formed in small quantities is assumed to be PbH 4. In alkali or weakly acid solutions, lead cathodes disintegrate at high current densities. This is believed due to the formation of unstable hydride PbH 2 at the cathode, and at current densities over 10-50 mA/cm3 the formation of PbH2 was believed to be quantitative. The organohydrides of lead (lV) are not as robust as the corresponding compounds of tin, germanium and silicon but are nevertheless sufficiently stable for an extensive study of the lead –hydrogen bond to be made. Trialkyl-lead hydrides and dialkyl-lead dihydrides were made by reducing the corresponding chlorides with lithium aluminium hydride at low temperature (78oC). The exchange reaction between organotin hydrides and organolead salts has also proved a useful synthesis for organolead hydrides. (n-C4H9)3PbX + (C6H5)3SnH (n-C4H9)3PbH + (C6H5)3SnX Trimethyl-lead hydride and triethyl-lead hydride were found to decompose to the corresponding tetra-alkyl-lead, lead metal and hydrogen. Evolution of hydrogen begins on warming to about –30oC to –20oC, but even at 0oC the hydrides are not completely decomposed for several hours. 11 CH2=CHCOOCH3 R3PbX + RH C2H4 R3PbC2H5 R3PbCH2CH2COOCH3 RX CH2=CHCN R3PbCH2CH2CN HCl R3PbH CCl4 PhC≡CH CHCl3+CH2Cl2 CH≡CCN R3PbCl + H2 R3Pb H C=C H C6H5 R3Pb CN H C=C CH2=C H CN PbR3 H H C=C R3Pb CN Some of the reactions of trialkyl-lead hydride are outlined in the above scheme. All compounds of lead are toxic, and lead poisoning has been long known and exhaustively studied. Humans from contaminated food and drink ingest minute quantities of lead regularly but normal body processes easily eliminate such quantities. The specified maximum allowable atmosphere pollution is 0.15 mg of lead per cubic meter of air. Lead poisoning can be cured and recovery is usually 100%. 12 Bibliography: D.Negoiu ‘Tratat de chimie anorganică’, vol II – Nemetale C.D.Neniţescu ‘Chimie generală’, M.Mironescu, C.Albu ‘Din istoria descoperii elementelor chimice’ F.Winter, D.Becherescu ‘Plasma în chimie’ C.Patroescu ‘Analiza apelor’ J.C.Bailar JR., H.J. Emeleus, Sir Ronald Nyholm, A.F. Trotman-Dikenson, ‘Comprehensive inorganic chemistry’, vol.l, vol.ll. GOOD LUCK ! from [email protected] 13