Download The IVth group of the periodic system contains the

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




SiSiSiSi




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




GeGeGeGe




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