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L I B RA FLY
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UN IVERSITY
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546
1^52-53
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INORGANIC SEMINARS
1952 - 1953
TABLE OF CONTENTS
1952 - 1953
Page
COMPOUNDS CONTAINING THE SILICON-SULFUR LINKAGE 1
Stanley Kirschner
ANALYTICAL PROCEDURES USING ACETIC ACID AS A SOLVENT 5
Donald H . Wilkins
THE SOLVENT PHOSPHORYL CHLORIDE, POCl3 12
S.J. Gill
METHODS FOR PREPARATION OF PURE SILICON 17
Alex Beresniewicz
IMIDODISULFINAMIDE 21
G.R. Johnston
FORCE CONSTANTS IN POLYATOMIC MOLECILES 28
Donn D. Darsow
METATHESIS IN LIQUID ARSENIC TRICHLORIDE 32
Harold H. Matsuguma
THE RHENI DE OXIDATION STATE 40
Robert L. Rebertus
HALOGEN CATIONS 45
L.H. Diamond
REACTIONS OF THE NITROSYL ION 50
M.K. Snyder
THE OCCURRENCE OF MAXIMUM OXIDATION STATES AMONG THE
FLUOROCOMPLEXES OF THE FIRST TRANSITION SERIES 56
D.H. Busch
POLY- and METAPHOSPHATES 62
V.D. Aftandilian
PRODUCTION OF SILICON CHLORIDES BY ELECTRICAL DISCHARGE AND HIGH
TEMPERATURE TECHNIQUES 67
VI. £, Cooley
FLUORINE CONTAINING OXYHALIDES OF SULFUR 72
E.H. Grahn
PREPARATION AND PROPERTIES OF URANYL CARBONATES 76
Richard *• Rowe
THE NATURE OF IODINE SOLUTIONS 80
Ervin colton
SOME REACTIONS OF OZONE 84
Barbara H. Weil
' HYDRAZINE BY ELECTROLYSIS IN LIQUID AMMONIA 89
Robert N. Hammer
NAPHTHAZARIN COMPLEXES OF THORIUM AND RARE EARTH METAL IONS 93
Melvin Tecotzky
THESIS REPORT 97
Perry Kippur
ION-PAIR FORMATION IN ACETIC ACID 101
M.M. Jones
IRON VERSENATES 105
R.L. Rebertus
AMPHOTERISM IN NON-AQUEOUS SYSTEMS 110
Harold J. Matsuguma
PLACE IN THE PERIODIC SYSTEM OF THE HEAVIEST ELEMENTS 115
Ervin Col ton
COMPLEX ANION DETERMINATION BY ION EXCHANGE 121
M.K. Snyder
THE NATURE OF THE FRIEDEL-CRAFTS COMPLEX 124
Clayton T. Els ton
ACID-BASE BEHAVIOR IN INERT SOLVENTS 127
Daryle H. Busch
METAL-DIAMINE CHELATE COMPOUNDS CONTAINING HIGHER-MEHBERED
RINGS 135
William E. Cooley
CYCLOPENTADIENYL COMPLEXES 139
Victor D. Af tandilian
DISULFUR HEPTOXIDE 144
A.B. Galun
SULFANURIC CHLORIDE 147
L.H. Diamond
ROLL CALL - THALLIUM TRI IODIDE 149a
Robert H. Marshall
PREPARATION OF COPPER HYDRIDE IN AQUEOUS MEDIUM 150
N.rl. Bojars
INORGANIC PAPER CHROMATOGRAPHY 154
Mary Joan Sirotek
INORGANIC COLUMN CHROMATOGRAPHY 161
S.a. Bartkiewicz
AMMONOLYSIS AND AMINOLYSIS OF SOME SUBSTITUTED SILICON HALIDES 164
A.S. Hay
THESIS REPORT 167
Niels C. Nielsen
THESIS REPORT - THE SYNTHESIS OF HYDRAZINE 168
Roger W, Sanftner
-3-
P. A. Horrigan 170
THE LITERATURE OP JNOHJAMIC CHEMISTRY
COMPOUNDS CONTAINING THE SILICON-SULFUR LINKAGE
Stanley Kirschner September 30, 1952
1« Nomenclature
At a recent conference of the International Union of Pure
and Applied Chemistry (5) , it was decided that since the term
"silicone" had been taken over as a trade name by industry, it
no longer had a place in systematic nomenclature. At this same
meeting a nomenclature for organo-silicon compounds was presented,.
It was decided that prefixes such as di- and tri- would be used
to designate the number of silicon atoms present in a compound.
Some examples are:
A - Silanes
1* SiH* is silane*
2» H3Si-SiH3 is d.isilane*
B - Silazanes
1. H3Si-PJH-SiH3 is disilazane.
C - Silthianes
1. H3Si-S-SiH3 is disilthiane*
1) - Siloxanes
1, H3Si-0~SiH3 is dilsiloxane.
E - Hydroxy Derivatives
1. H3SiOH is silanol.
2. H2Si('0H)2 is silandiol.
3. H3Si-0~SiH20H' is dislloxanol.
For a more comprehensive survey of organo-silicon nomenclature,
the reader is referred to reference (5) . .
2« Historical Summary
Until as recently as ten years ago there were less than
about a dozen known compounds which contained the silicon-sulfur
linkage. The best characterized of these were: (a) silicon
disulfide, SiS2, which was prepared (9) by heating a mixture of
silicon, carbon, and carbon disulfide; (b) silicon monosulfide,
SiS, and (c) silicon oxy sulfide, SiOS, which were prepared
simultaneously (4,12) by heating silicon and sulfur vapor in a
porcelain reaction tube; (d) silicon thiochloricle , SiSCl2, which"
was formed (2) along with SiS2 and SiCl4 by passing sulfur mono-
chloride over crystallized silicon in a heated tube; (e) silicon
thiobromide, SiSBr2, which was prepared (1,11) by the reaction
between SiBr* and H2S at 150oC. with an A13r3 catalyst; "(f) sili-
con chlorohydrosulfide, SiCl3SH (today called trichloro-thio-
silanol), which was formed (11,8) by heating a mixture of SiCl4
and H2S in a porcelain tube; and (g) silicothiourea, SiS(NH2)2,
4 ■ •
'.'," '.. .. .'....., . .
; -*»
-2-
which was prepared (l) by the reaction between SiSBr2 and gaseous
NH3 in anhydrous benzene; the M-I4Br produced was removed by
washing with liquid ammonia.
3* decent Developments in the Chemistry of Compounds Containing
the Silicon-Sulfur Linkage
In 1950, Eaborn (6,7) described what he believed to be the
first organosilicon sulfides to be reported. (However, Friedel
and Ladenburg (8) described the preparation of Si(0C2H5)3SH in
1872). These compounds are hexaethyl- and hexamethyldisilthiane,
which were prepared as products in a conversion series developed
by Eaborn, The series is:
R3SiI -> (R3Si) 2S -> R3SiBr -^SiH)-^3^^-* R3SiNCS -> R3SiNC0 ->
( (R3Si)2o)
J H3SiP J
Any compound in the series may be converted to any other on its
right by heating with the appropriate silver salt. The yields
are frequently better than 90$, and the reverse conversions do
not seem to occur. For example, the hexa-alkyldi silthianes were
prepared by heating the corresponding trialkyliodosilane with
Ag2S.
Eaborn states that the silthianes are the sulfur analogues
of the siloxanes, but they are (a) less stable (thermally) than
the siloxanes, and (b) completely hydrolyzed when brought into
solution. He feels that this is due largely to a weaker bond
between silicon and sulfur than between silicon and oxygen, and
he cites the values for the heats of formation of the Sl-S bond
and the Si-0 bond given in Sidgwick (13) in support of his ideas:
Heats of Formation of Links from Atoms (Ha)
Si-0 89.3 kcal/mole
Si~S 60,9 kcal/mole
However, Eaborn prepared these di silthianes from trialkyl-
iodosilanes, and, according to Larsson and Marin (10), these are
rather difficult to obtain. They, therefore, devised a method
for the preparation of hexaethyldisilthiane which does not
involve any of the iodosilanes. They report that, after passing
H2S through (C2Hs)3SirJH2 (163 grams) for several hours," they
obtained 34.5 g of (C2H5)3SiSH and 95,2 g of (C3H5 )3S5>S"~Si(C2H5) 3 .
In addition, they found that similar treatment on tr'i-n-propyl-
aminosilane gave tri-n-propylthiosilanol and hexa-n-propyldisil-
thiane . The methyl derivatives were not prepared because the
authors were unable to prepare the starting material,
trimethylaminosilane©
. I
'.)•'■ .
-3-
Champetier, Etienne, and Kullman (3) continued the study of-
silicon-sulfur compounds and have recently described the prepara-
tion of some silan thiols and " thiosilicones"* Although Friedel
and Ladenburg (8) had been able to prepare trichlorothiosilanol
by heating SiCl4 and H2S above 600°, a method utilizing so^'high a
temperature could not be used in the preparation of organo-
silicon compounds, so the authors reacted alkali and alkaline
earth hydro sulfides with alkylhalosilanes to prepare the alkyl-
thiosilanols. For example, the authors reacted HSS with a Grignard
Reagent to get a magnesium hydrosulf ide:
2H2S + 2RMgX -> (HS)2Mg'MgXa + 2RH
which was reacted with the trie thy lhalosilane (C2H5)3SiY, The
products were (C2H5)3SiX (when the halogen Y was classified before
the halogen X in the series F, Ql, Br, I) and hexaethyldisilthiane!
1, 2(CsH5)3SiY +,MgX2 -+2(C2H5)3SiX + MgY2
2'(2(C2H")3iiY)°r + %^H)s-^2(C2H5)3SiSH + ^J or
3. 2(C2H5)3SiSH -> (G2H5)3Si-S-Si(G2H5)3 + H2S
In order to obtain trimethyl thiosilanol with only a small
percentage of the hexamethyldisilthiane, LiSH was used in place
of a magnesium compound:
LiR + H2S *hJ LiSH + RH
06H50H3 —
When 0,75 moles of LiSH was treated with C%5 moles of trimethyl-
chlorosilane at room temperature, a 48^ yield of (CH3)3SiSH and
only a small quantity of hexamethyldisilthiane were obtained upon
fractional distillation*
These authors also report that a convenient way to obtain a
high yield of various disilthianes is to react an alkylhalosilane
with HaS in the presence of pyridine:
2(C2Hs)3SiCl + H2S + 2G6H5k-+ (C2H5 ) 3Si-S-Si (G2H5) 3 + 2C6H4N.HC1
Diethyldichlorosilane leads to tetraethylcyclodisilthiane 5
room C2H5 ^Sv ^G2H5
2(G2H5)2SiCl3 + 2H2S + 4C6H5N ^L Si ^Siv + 4G6H5N.HC1
(in petroleum dry ' ' C2HS ^S"" G2H5
ether)
in a 78% yield, which the authors claim is the first example of
a " thiosilicone" , and which is similar in structure to silicon
disulfide:
.Si £i Si
^ XS^ xs^ ^
A great deal of additional investigation remains to be
carried out in the field of silicon-sulfur chemistry, as can be
seen from the following tabulation of the few silicon-sulfur
compounds which have been reported. Even these have not been
studied to any great extent, and they should be investigated
further*
■'■ ,..
,,
,. (, ;"
H
Compound
SiS2
SiS
SiOS
SiSCl2
SiSBr2
SiCl3SH
SiS(NHs|s
[ (C2He)3Si]2S
[(CH3)3Si]2S
Si(OC2H5)3SH ■
(C2H5)3SiSH
(CH3)3SiSH
(n~C3H7)3SiSH
C(n-C3H7)3Si]2S
(C2H5)2Si' Si(C2H5)2
-4-
Melting
Point °C.
Boiling
Reference
Subl.
White Heat
9
Subl. at 940
(20 mm.)
4
4
75
185
2
93
150 at
(18.3 mm,)
96 - 100
1
11
1
278-279 at
(750 mm.)
163-164
6,3
3
164
8
158
10
77-78
3
2
83-84 at
(7 mm.)
168 at
(7 mm.)
160 -* 162
10
10
3
References,
1. Blix, Ber., 36, 4218 (1903).
g, Blix and Wirbelauer, Ber., 36, 4220 (1903).
3. Chapetier, Etienne, and Kullman, Compt. rend*, 234 T 1985 (1952) v
4. Colson, Bull. Soc. chim., II, 38, 56 (1882).
5. Compt. rend. XV th Conf. Int. Union Pure and App. Chem. ,
(Amsterdam), 1949 f 127.
6. Eaborn, Nature, 165, 685 (1950).
7. Eaborn, J„ C^em. Soc., 1950, 3077.
8. Friedel and Ladenburg, Ann. Chim, Phys,, 27., 416 (1872).
9* Hempel and von Haasy, Z, ^anorg, Chem,, 23, 32 (1900),
10* Larsson and Marin, Acta Chem, Scand., 5,, 964 (1951).
11. Pierre, Ann. Chim. Phys., Ill, 24, 286 (1848),
12« Rankin and Rivington, Proc, Chem, Soc, £4, 131 (1908).
13, Sidgwick, The Chemical Elements and Their Compounds, Vol. I,
p. XXXI, Oxford University Press, London (1950).
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ANALYTICAL PROCEDURES USING
ACETIC ACID AS A SOLVENT
Donald H. Vilkins October 7, 1S53
INTRODUCTION
Acid-base relationships in nonaqueous solvents were first defin- •
ed by Franklin in 1905 (l) and later extended by Germann in 1935 (2).
Although the literature on nonaqueous solvents dates back nearly fif-
ty years, the praotical use of these solvents in analytical proce-
dures is covered only in the more recent publications. Acid-base
titrations in nonaqueous solvents are particularly useful. They may
be performed with simple techniques and without elaborate equipment.
The procedures are accurate and rapid. Many organic compounds which
are too weakly acidic or basic act as relatively strong acids and
bases in appropriate nonaqueous solvents. Acetic acid, an acid sol-
vent, is useful for titrating many organic bases which are either
only slightly soluble in water or too weakly basic to give sharp end
ooints. Most aromatic, aliphatic, and heterocyclic amines may be tit-
rated as bases in acetic acid.
DISCUSSION
Kolthoff and Williams (3) studied the order of acidity of sev-
eral acids in acetic acid. They found: HClO^HBr) H^SOy > HC1 > HN07.
Of these acids, perchloric is the most suitable titrant. It is the
strongest acid, has fewer insoluble salts, and has only one replace-
able hydrogen ion. Hydrochloric acid has a high escape velocity; a
1.0 N solution of hydrochloric acid in acetic acid fumes about the
same as 16 N solution in water, Sulfuric acid does not always act
as a monobasic in its reactions.
Many salts which do not change the reaction of water undergo sol-
volysis in acetic acid (4). The acidity of ootassium salts of the
following inorganic acids decreases in the order C10V'I'">Br~>cr>N0^.
-,...■ . t ■ ,.'
...... ;/
.;
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• •■•■■ ■■.-]• ;•''.•.""!' ' . -■ ■ ■ . . ■
........ * ~.
.... :•.-;■•'.-• •• ■' • : . ' ■ ■• *
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-6-
Sa.lts having a common ion show decreasing acidity in the order Mg*4" >
Ca"*>Sf*>Ba^>( Ag*) > Li' >Na+ > NH,} ) K%Rb+.
■Vater interferes with the end point of the titrations of bases
in acetic acid and conseauently should be removed. The water intro-
duced in making uo the titrant by dissolving 70 to 72 per cent Der~
chloric acid in glacial acetic acid is removed by adding the calcu-
lated amount of noetic anhydride. The coefficient of cubic expansion
is large enough for acetic acid so that an accurate temperature con-
trol is necessary. The temoerature of the titrant should be the same
for the standardization and the titration of the unknown; otherwise a
correction must ;rbe made for volume changes'. For precise research work
weight burets are recommended.
Hth volume burets the accuracy of titrations in acetic acid is of
the same order as aqueou titrations (0,2%).
DETERMINATION OF THE END POINT
A. Potentiometric titrations
Oonant and Hall (5) suggested the use of a saturated chloranil
electrode (tetrachloroquinone and its hydroquinone). Its advantages
are chemical inertness and low solubility. Theoretically the depend-
ability of the chloranil electrode depends on the solution being at
eouilbrium with both the solid ouinone and the hydroouinone. The
solution to be titrated is saturated by adding an excess of the solids
and stirring for ten minutes at 25°, the working tempera.ture. If a
volume of- titrating solution is now added an unsaturation occurs.
This is not important if the volume added is small Compared to the
total volume of solution. This is to be expected unless there is a
large difference in the rate of solution of two solids.
The reference electrode was a saturated calomel electrode which
was connected to a beaker of saturated potassium chloride by a. siphon
filled with saturated potassium chloride. The chloranil electrode
..-' ■•.
..
. . -i.
-7*-
was connected to the bsaker of saturated potassium chloride by a
siphon filled with a supersaturated solution of lithium chloride in
acetic acid which was prevented from crystallizing by dissolving a
small amount of gelatin in the acetic acid, The amounts of gelatin
and lithium chloride may be varied over wide limits without appreci-
able changes in the performance of the bridge. For special purposes
the siphon from the reference electrode was stoppered with glass and
dipped directly into the solution to be titrated.
The e.m.f. was measured with a potentiometer and a quadrant
electrometer.
Fritz (6) used a glass electrode as an indicator electrode and
a silver wire with a thin coating of silver chloride as a reference
electrode. A Beckman Model G pH meter was used to determine the pot-
ential. This system eliminates the salt bridge since the silver
electrode may be immersed in the solution to be titrated.
The titration curves are obtained by a plot of e.m.f. vs. ml. of
reagent added. The curves are comparable in shape to those obtained
in aqueous solutions. The end point may be taken directly from the
curve but is more accurately determined by plotting a e/^ V vs. ml. of
reagent added.
3. Visual titrations
Nadeau and Branchen (7) found that crystal violet, ec -naphthol-
benzein and benzoylauramine are suitable indicators for the titration
of bases in acetic acid. The *cid nnd alkaline colors of the indic-
ators are giv^n in table I.
These indicators, however, do not give simple color changes. The
color is dependent on the ionic strength of the solution as well as
the pH . Accurate results may be obtained by first titrating pot-
entiometrically with the indicator present to determine the proper
color at the end point. Subsequent determinations are made by tit-
i ..'.
'■ ;
...
: ■• '
. .... _.,..,.,
•' .■ •
. :A-r::
.. ■...- . '
.) *
• _ 4 . *
<• -c
...-.•.■ I, ' ; ■■ '
•+;-;,
.... .r -. ■ ,., . .->
-8*-
rating to the same color using a standard for comparison. The ind-
icators are 6#1$ solutions in acetic acid or chlorobenzene.
TABLE I
Indicator , Acid color Alkaline color
crystal violet green blue
tt-naphthol-benzein green yellow
benzoyl aur amine yellowish blue-green
To prepare a. 0.1 N perchloric acid solution in acetic acid mix
8.5 ml. of 73-t perchloric acid with 200 to 300 ml. of glacial acetic
acid and add 20 ml. of acetic anhydride. Dilute to 1 liter with
glacial acetic acid and allow to stand overnight to permit complete
reaction of the acetic with the water present*
The perchloric acid is standardized against potassium acid phth-
alate which has been used as primary standard acid in aqueous sol-
utions and is now finding use as a primary standard base in acetic
acid. About 0.5 gms, of potassium acid phthalate is accurately
weighed and added to 60 ml. of glacial acetic acid. The mixture is
refluxed gently to dissolve the salt. After cooling methyl violet
(O.lfc in chlorobenzene ) is added and the solution titrated with per-
chloric acid. The precipitate of potassium perchloratr- does not
interfere with the end point.
To prepare 0.1 N perchloric acid in dioxane dissolve 8.5 ml. of
77$ perchloric acid in 1 liter of dioxane. Standardize with pot-
assium acid phthalate as described above.
A 0,1 N solution of sodium acetate may be prepared by dissolving
8.2 gmsf of anhydrous sodium acetate in glacial acetic acid and dil-
uting to 1 liter with acetic acid. The solution is standardized
against a standard perchloric acid solution.
Amino acids (7,8)
Acids which have an acid ionization constant of 10 or less in
...
', ' t,
,.i ..,.;- V
.
- ■■* '
•
•:•■ .:
: ..
■.■■* • •. .
■*• • +
... • .. <
i*0 i. .' . rr? j"iii
-9-
aqueous Solution are neutral in acetic acid.- Consequently amino
acids may be titrated as bases.- The sample is dissolved in exactly
50 ml 4 of 0.1 N perchloric acid. Two drops of methyl violet indic-
ator are added and the excess acid back-titrated with 0,1 N sodium
acetate*
Amines (8-15)
A eample containing two to four eauivalents is dissolved in 50
ml. of acetic acid* Add two drops of methyl violet and titrate with
0.1 N perchloric acid to the first disaooearanoe of the violet color.
The end point may be determined ootentiometrically in the presence
of highly colored compounds which do not give a sharp end point.
Micro quantities of amines may be titrated with 0*001 N perchloric
acid. Most aliphatic and aromatic amines can be titrated unless they
are heavily substituted with halogens, nitro, aldehyde or other nega-
tive groups.
■Vagner, Brown, and peters (14) have proposed a procedure for diff-
erentiating between primary, secondary and tertiary amines. Ter-
tiary amines are determined first by treating the sample with 20 ml.
of acetic anhydride and 2 ml, of acetic acid and allowing the sol-*
ution to stand for three hours. P rimary and secondary amines are
converted to almost neutral acetylation products. Thirty ml, of
acetic acid is added and the solution titrated ootentiometrically
with perchloric acid* primary amines react with salicylaldehyde to
form Schiff bases which are weaker than the primary amine. By tit-
rating aliquots of the sample it is possible to determine all three
amines in a mixture.
Salts (8,9,15,16)
Salts of weak carboxylic acid reaot with a strong acid according
to the following eouation*
NaA + H+— * Na**HA
1 -. •
-10-
If the acid HA is sufficiently weak the salt may be titrated Quanti-
tatively. This procedure has been applied to the salts of a few
weak acids in aqueous solution. In an acid solvent, such as acetic
acid, the number of salts which may be titrated is considerably
larger. The alkali salts of most carboxylic acids and some inorganic
acids can be titrated in acetic acid.
Pefer and Wollish (15) have introduced a method for titrating
halide acid salts of organic bases. Mercuric acetate binds the halo-
gen anions without effecting the titration of the base portion of the
salt with perchloric acid. The reaction is believed to proceed as
follows :
Hgli
Orga.nic base
HI
HBr
HOI,
excess
HgAcx
acetate of
organic base
HgBr^
HgClA
HOlOu
(stdj
^ perchlorate of
'..organic base * HAc
voltage change
The weighed sample is dissolved in 8 ml. acetic acid. If the
salt is an acid halide of an organic base, ^dd 10 ml. of mercuric
acetate (6 gmB/lOO cc. HAc), titrate with perchloric acid using the
potentiometric method to determine the end point.
CONCLUSION
The use of nonaqueous solvents such as acetic acid eliminates
many tedious analytical procedures. The applications in industrial
work pre extensive ; however, the theoretical background for many of
these procedures has not been developed as much as might be desired.
There is a need for considerable research to develop suitable indic-
ators and electrode systems. There are over five hundred acid-base
indicators available but only a few can be used in acidic solvents
such as acetic acid. Indicators with different transition ranges
should be available for the entire transition range.
'
■■■'..■■■
r ..;
; -,*a t
-11-
BI3LI0GRAPHY
1. Franklin, J. Am. Chem. Soc. 27, 820 (1905)
2. Germann and Timpany, Ibid., 47_, 2275 (1925)
3. Kolthoff and .Villman, Ibid., £6 1007 (1934)
4. Ibid., 56 1014 (1934)
5. Conant and Hall, J. Am. Chem. Soc, 49, 3047, 3062 (1927)
6. Fritz, Anal. Chem. , 22_ 1029 (1950)
7. Nadeau and Branchen, J. Am. Chem. Soc J57 1363 (1935)
83 lH^iZ'Jl' ?' ^cid~B*se Titrations in Nonaqueous Solvents, G. F.
Smith Chemical Co. (1952) '
10. Wilson, J. Soc. Chem. Ind. (London), 67 237 (1948)
11. 'Vittman, Angew. Chem. , A60, 330 (1948)
12. Kahane, Bull. soc. chim. France, 18 92 (1951 )
13. Keen and Fritz, Anal. Chem., 24, 564 (1952)
L4, Wagner, Brown, and Peters, J. Am. Chem. Soc, 69 2609 (1947)
L5, Pifer and Wollish, Anal. Chem. 24 300 (1952)
L6. Pifer and dollish, J. Am. Pharm. Assoc, 40 609 (1951)
, >v . ; . . » •
-12-'
THE SOLVENT PHOSPHORYL CHLORIDE, POCl3
S. J. Gill October 14, 1952
INTRODUCTION
The interest in non-protonic solvents has been stimulated
within recent years both from a practical and a theoretical
point of view. The properties of a particular solvent have a
marked influence upon various compounds wi th regard to sol-
ubility, ionization, reactions, and the possibilities of
unique syntheses, Thus non-protonic solvents find a special
place by providing liquid media which are devoid of hydrogen*
Theoretical explanation of such systems has extended and
verified acid-base concepts. Experimental work with non-
protonic solvents necessitates a more accurate description of
solubility phenomena, suggests some of the physical properties
involved, and provides a possible means for clarification of
liquid state reactions. Phosphoryl chloride suggests practical
as well as theoretical utility as a non-rprotonic solvent.
PROPERTIES
Some of the physical properties of phosphoryl chloride
are listed in the following table:
Table 1
Molecular weight
Melting point, °C
Boiling point, °C
Critical temperature, °C
Density, g/cc
Molec. volume, cc/mole
Specific conductivity, 25°C
Dielectric constant
Cryoscopic constant
Ebullioscopic constant
Surface tension, dyne/cm
Heat of vaporization kcal/Mol
Heat of fusion kcal/Mol
Troutons constant cal/degree
Heat of formation (liquid)
kcal/Mol
Kinematic viscosity,
centistokes
153,39
1.25
105*8
329 (2)
1.645 (25°C) (1)
101.85 (105. 8°C) (1)
1.7 x 10-sohm icm"i
(3) 1.63 x 10""6 (i)
13.9 (22°0) (4)
76.8 (5)
54.7 (6)
31.9 (18°C) (3)
8.06 (6)
3.07 (5)
21.3
146
0.669
(7)
(1)
A less extensive list of properties of other non-protonic
solvent materials is given in Table 2 (8).
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-13-
Phosphoryl chloride resembles water in liquid range and
low self-ionization* Its dielectric constant is greater than
that for S02# The relatively large molecular volume is worthy
of note* The dipole moment of the P-0 bond has been calculated
as 3,5 (9). Molecular weight determinations in benzene give
values from 153 '4 to 166, thus showing slight if any assoc-
iation (10).
PREPARATION OF WATER-FREE P0CL3
Commercial P0C13 is readily available and costs from 13
to 15 cents a pound (11), The methods or production are:
a) heating a mixture of P2Q5 and PC15, or b) passing chlorine
over calcium phosphate and charcoal at 750°C (12) .
Purification for solvent use requires the elimination of
all traces of water. Water reacts with phosphoryl chloride
to give phosphoric acid and hydrogen chloride, thereby raising
the conductivity of the solvent. Walden (3) treated the
material with Ps05 and distilled the dry P0C13 from the result-
ing phosphoric acid* However this does not eliminate the
hydrogen chloride. To do so Cady and Taf t (13) added metallic
sodium and distilled the resulting mixture at atmospheric
pressure. This technique •ccasionally results in explosion.
In order to minimize the possibility of explosion
Gutmann (1) employed the following procedure! the commp rical
product was first distilled using a 30 cm. column ground glass
apparatus; sodium was added to the distillate and carefully
refluxed; after cooling, the material was distilled twice
under high vacuum. Explosion does not occur under these
conditions.
CONDUCTIVITY AND SOLUBILITY STUDIES
The specific conductivity of phosphoryl chloride passes
through a maximum at 88°C . Over the same temperature range
both the density and viscosity decrease monotonically , A
possible explanation (l) of this phenomenon is given by
considering the following equation:
P0C13 » P0C12+ + CI"
or 2P0C13 = P0C12+ P0C14"
These ions, more accurately thought of as in a solvated con-
dition, provide the dominant species of conductivity. The
solvated ion stability decreases with increasing temperature
especially since P0C13 has a large molecular volume. Thus it
might be expected that a revprsal of this equilibrium takes
place with the formation of P0C13 at a higher temperature.
This behaviour is not restricted to this particular solvent
alone, but has been noted for BrF3 (14) as well as other
non-pro tonic materials.'
-14-
Information suggesting the above ionic species is given
from electrolysis studies of Cady ancl Taft (13) where Fe 3
was reduced to Fe 2 in PCC13 , The formation of a highly
reactive material at the cathode has been interpreted by the
following reactions (l):
P0C12 + e~ -» (P0C12) %
2(PCC12) -> P0C1 + P0G13
Clusius and Haimerl (15) have shown that HC135 exchanges
chlorine atoms with P0C13 more slowly than wi th AsCl3 and PC13,
but more rep idly than with SiCl4 and S2C13, This implies
ionic species, A large number of compounds dissolve in P0C13 #
The following illustrative list is taken from V, Gutmann (16) .
1. Soluble
A* Cryosoopicly determined, dissociation occurs:
PC15, PBr6, AuCl3, BiCl3, IC13 , SC14 , BiBr3 , Bil3 ,
PtCl4.
B, Gryosoopic and conductance, monomolecular species:
SiCl4, SiBr4, SnBr4, N205 , 0S04
G. Compound formation: A1C13 BBr3 , SbGl5, SnCl4,
TeCl4, TiGl4 BC13.
D. Tend to associate: SnCl4 •P0C13, G1207 .
E. Colorless solutions, no knowledge of species:
ASC13, AsBr3, Snl4, Hgl8, FeCl2, Br2, KC103 .
F. Colored solutions: Cl2, Br2, I2, Nal, RbICl2,
(Ch3)NI,
2. Moderate or slightly soluble:
KIO3, KC104, KI04, FeCl3, Hg(CN)2, HAUC14
3» Electrolytes, dissociates slightly:
LiCl, NaCl, KCK, KCNC
4, Insoluble:
KNO3, K2C204, T1G1, AgCl, HgGl, CUC12, CaCl2, SrCl2,
BaCl2, ZnGl2, CdCl2, CdBr2, Cdl2, MnCl2, VG13 ,
K3Fe(CN)6, P205, CrC3, I205.
Most of these compounds have been classified only from a
qualitative point of view. Gutmann has established the
following numerical data (16) :
• ; ' > •
,K:
-15-
P0C13 at 2Q°C.
Solubility
g&L—
LiCl
O.Ob
NaCl
0.31
KC1
0.60
MUC1
0.46
RbGl
0.87
CsCl
1.26 .
(CH3)4
NCI 2.00
KF
0.40
KBr
0.51
KI
1.71
KCN
0.73
KCNO
0.80
KCNS
0.76
Specii
>!<
3 Cond .
Equivalent Cond.
at sat.
cone.
V=
=1000 1/Mol
6.6
X
10~6
4.0
3.0
X
10~5
6.4
3.4
X
10~.s
6.7
3.6
X
10~5
6.9
8.3
X
10~s
14.6
1.1
X
10-4
16.0
5.5
X
10~4
37.6
2.6
X
10"*5
6.4
4.3
X
10~5
14.5
1.2
X
10~4
23.1
3.3
X
10~5
7 0
3.1
X
10-5
9.0
2.9
X
10-5
6.6
Gutman/icalls attention to thp observations that among the
halides of the above list ionization increases, as is evidenced
by the conductance values wi th increasing size of thp cation
or anion. Such is not the case with KCN, KCNO, and KCNS,
which show virtually no change in conductance or solubility.
Veil definled instances of inorganic compound formation with
POCI3 are:
a. P0C13«BC13, rhombic, &FS9B of dissociation equals
-4.6 kcal. No BF3 compound. (17)
b. SnCl4.2PCCl3 or Cl4Sn( :0:PC13 ) 3 (18)
c. 2ZrCl4«P0Cl3 (19)
d. 3ZrCl4»2P0Cl3 and 3Hf Cl4 *2P0C13 (20)
e. A1SC16«2P0C13 (21)
1.
2.
3.
References
V, Gutmann, Monatshefte Fur Chemie, 83, 164 (1952),
Institute- for general chemistry, Vienna.
W. Ramsay and J. Shields, Jour. Chem. Soc . , 63,, 1108 (1893),
P, Walden, Z. physik. Chem., 43., 445 (1903),
4. H. Schundt, J. phys. Ch^m., 5, 515 (1901) .
5..P, Walden, Z. anorg. Allg. Chem., 68, 307 (1910).
6. P. Walden, Z. physik. Chem., 65., 271 (1909).
7. J. Thomsen, Ber. dtsch. chem., 16, 2619 (1883).
8. Spandau and V. Gutmann, Angp wand to Chemie, 64,, 94 (1952).
9. C. p„ Smyth, G. Lewis, A. J.Grossman, and F. B. Jennings,
Journal Am. Chem. Sec, 62,, 1219-23 (1940).
A. R. Morgan and S. T. Bowden, Trans. Faraday Soc, 36,
394-7 (1940) . ' ~*
Chem. and Eng. News, 3_0, 3993 (1952).
Rozhdestvenskii, Trans. State Inst. Applied Chem., (USSR),
£0., 47 (1934) . '
10.
11.
12.
••
-16-
13. H. Pr Cady and R. Taft, Journ. phys. Chem., 29,, 1057 (1925)
14. A, A. Banks, H, J. Em^leus, and A. A, Woolf , J. Chpm. Soc.,
1949.. 2861.
15. K. Clusius and H. Haimerl, Z. physik. Chem., 51B, 347-51
tl942) .
16. V. Gutmann, Monatshefte Fur Chemie 83, 279 (1952).
17. A. B. Burg and M. K. Ross, J. Am. Chem. Soc, 65, 1637-8
(1943).
18. F. B. Garner and Sugden, J. Ghem. Soc., 1929, 1298-302.
19. a. E. vanArkel and J. H. deBoer, Z. anorg. allgem. Chem.,
141, 289-96 (1924) .
20. D, M. Green and J. J. Katz, J. Am. Chem. Soc, 71,, 3843-4
(1949),
21. W. T. Dye, Jr., J.. Am. Chem. Soc, 70, 2595-6 (1948).
. \
)
-17-
M5TH0DS FOR PREPARATION OP PURE SILICON
Alex Beresniewicz
October 28, 1952
Introduction
During World War II high purity silicon was in great
demand because of its use as a semiconducting material in
crystal rectifiers of radar sets. Impurities, especially
non-metallic, have a very marked effect on some physical pro-
perties of silicon, mainly on its specific resistance, z>. (1)^(2),
Measurements of specific resistance provide, therefore, a very
sensitive and accurate method for determining the purity of a
given sample. According to Pearson and Bardeen (2) pure silicon
should have a resistance of 106 ohms cm, at room temperature,
but this value has not yet been attained.
Review of, the General Methods, for. Preparation of Si^iqoi;
I, Reduction of
furnace •
Si03 with C or CaCa in the electrio
2, Reduction of Si02 with Al in a thermite type reaction,
3, Si0a + 2Mg->Si + 2MgO.
4, Electrolysis of a solution of S102 in a molten mixture
of sodium and potassium fluorides. The alkali metai set free by
the current reduces Si02 to Si,
5, SiCl* + Zn
vap,
900°C.
*
ZnCl2 + si
6, Reduction of K8StF* with Zn or Al.
Becket (3) is usually credited with the first successful
attempt to prepare pure silicon, He treated the crude eleotric-
furnace product with an HP solution at room temperature, thus
extracting most of the Pe impurities, Schaff (4) purified the
99.8$ Si by subjecting it to a heat treatment in 3 nonoxidizing
atmosphere, The DuPont Company manufactured pure silicon on a
technical scale by recusing SiCl^ with Zn vapor, Their product
was found to be free of metallic impurities, as determined
spectroscopioally (5),
ircprpyej^ jjfe.tfoqfte, £°£ Preparation, of Pu^e SUlfiPE
Wartenbejrg tried to prepare silicon of possibly highest
purity by the following methods; (6)
. \
* y * * . • ••«
> * ■
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.- ".'.5
i
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-18-
t Crystallization Al
A, Reduction {from molten Ag ^Purification
11 IZlXm: °f withA1 lraetals: znhy™lnT
2 6 p -,_ . Vin vacuum
IB, electrolysis ^*
*— of fused
K2SiP6 + NaCl
II. Thermal dissociation of Sil4
III. Reduction of SiCl4 with Zn vapor
^ si* lron-free KsSlF6 was prepared by precipitaticn by
^oS? of KG1 to a solution of H2SiF6, which was acidified
with HC1, and decanting several times with water.
anno,?7 heatl?g pure K2SiF6 with Al and NaCl for two hours at
yuu c. a regulus of Si was obtained. Samples of it were melted
in quartz crucibles with Al, Ag and Zn respectively. These
metals have relatively low melting points and do not form
silicides.
Table 1 records the solubilities of Si in Al, Ag and Zn
and vice versa. ' &
Table I
Metal
Solubility of
Si at 900°C
Solubility of the metal
in silicon
Al
30$ by weight
0.15$ by wt.
0.14$ by vol
Ag
5$
0.5 - 0.6$
0.15$ by vol
Zn
1.6$
*
0.5 + 0.05$
Al
0.16$ by vol
//artenberg next tried to remove the metallic impurities
by melting the respective silicon samples in a vacuum oven.
After SiAg had been fused in a sintered corundum boat, its
analysis showed the presence of 0.5$ Ag and of 0.9$ Al.
Apparently the A1203 of the container was reduced by molten
silicon. At first this might appear unusual, because the
heat of formation of Al203 is larger than that of 1.5 Si02.
It can be explained, however, by making use of the approximate
form of the Nernst's heat theorem.
Attempts were made to melt silicon in such a manner that
it would not be in contact with any wall material. The result-
ing product was covered with a thin bluish film, probably of
Si3N4 or SiC.
• . \
t [
•
• -: . ' . ' ■ ii • " ":'
IT
*
■<
.U.
( j
-19-
During electrolysis of molten K2SiF6 the cathodically
separated alkali metal reduces the melt to silicon. NaCl
is usually added in order to depress the dissociation of
K2SiF6. Silicon prepared by this method is soluble even in
l^HF, Special precautions have to be taken therefore in extract-
ing it from the electrolytic residue-. Analysis of the product
showed the presence of 0*1$ Na, 1%G and of some iron from the
graphite crucible*
II. Efforts to prepare pure silicon by thermal decomposi-
tion of Sil4 were unsuccessful* At 1200°C this compound
decomposes; silicon crystallizes on a carbon wire, but the
iodine recombines only slowly with Si,
III. Zinc vapor reduces SiCl4 but not the SiOa of the
quartz container.
900°C
2Zn(vap.) + SiCl4
^ Si + 2.2nCl2; 4H = -100 Cal.
Using this method Wartenberg prepared a product which had a
specific resistance of 11000 ohms cm* The product was a fine,
wool-like material composed of transparent crystals. (6)
Table 2
Method of Preparation
Impurities
% by Vol.
T>[JI to* ]
1, Ideally pure Si
None
106
2. SiCl4 + Ln (Wartenberg)
?
1.1 x 10*
3. SiCl4 + Zn (Du Pont)
<0.001^ c
102
4. Sams as (3) , only melted
before
<0.001# c
50
5. Crystallized from Al or Ag
0.16$
2 x 103
3 x 103
6. Electr. of K2SiF6
0.25^ /l/flt
103
When Si is fused in quartz containers under He atmosphere
its specific resistance drops considerably. SiO could be
formed, but there is no characteristic absorption band of
it (2.73^) after the melt solidifies. SiO is thermody-
namically unstable at 1000°C (7) • Wartenberg proposes
following reaction: 2SiO — ► Si + SiOs;
the
"SiO -+ Si + Si02; to take place (6).
The fact that silicon after being melted becomes a much
better conductor suggests that it might be contaminated by
oxygen. It could come from hot quartz walls (6) , and it is
known to make silicon an N-conductor. (l)
It seems that the best method for preparation #f pure
silicon would be to melt a fairly large amount of Si powder
in an evacuated quartz container by means of a lens (6).
That this is theoretically possible was shown by Stock (8)
who melted silicon using a lens 40 cm* in diamster.
'
■
, •
'•
? • - S '; 5".
. ■
■ " ;.-■■'■■ ■' . V--
?"'*!'.'? ' 1 )..
: ' '»■
ta .
.1
.' • ' ■ ' '
1 ■ 1
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» •
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. ■• I
,i ' .'
-20-
References
1# hQ0rS{ and Whltner» Crystal Rectifiers, McGraw-Hill , Inc.
2, Pearson and Bardeen, Phys. Revc 75, 865 (1949)
3, Becket, U.S. Pat. 1,386,2*7, Aug. 2, 1921.
4, Schaff , U.S. Pat. 2,-502,582, June 25, 1946.
5, Lyon, Olson and Lewis, J. ^lectrochem. Soc., 9J3, 359 (1949).
6, Wartenberg, Z. anorg. u. all gem. Chem, 2j65., 186 (1951).
?* ?1950?P and H°rnle' Z# anorg. u. allgera. Chem., £6£, 261
8. Stock, Berichte, J&, 2863 (1909).
-21-
• IMlDODISUtflN^.MIDE
G, E. Johnston November 4, 1952
I, General Introduction
The availability of ammonia has focused attention and
aroused interest in the nitrogen system of compounds. This point
of view utilizes ammonia as a parent substance and solvent, as
compared to the oxygen system of compounds which places water in
that role. By the replacement of the oxygen atoms in many oxygen-
containing compounds with their nitrogen analogs, many interest-
ing compounds can be visualized. For example, by substituting
nitrogen for oxygen in sulfuric and sulfujious acids, a series of
related compounds called aquo ammono sulfuric and sulfurous acids
can theoretically be postulated!
•-, S02NHA
^ Mi2S02NH2 ^2
HOSC2OH -> NH2S02OH *7 ^ NH(S02NH2)
^» NH(S02OH)2 -*N(S02OH)3
-^ SONH
-.NH2SOHH2
HOSOOH -+ NH2S00H ^ ^* NH(S0NH2)
^ NH(SOGK)2 -^N(SOOH)3
3
Of the two similar series of compounds, more is known concerning
the aquo-ammono sulfuric acids than is known about the aquo-
ammono sulfurous acids. Most of the imides and amides of sulfuric
acid have been recognized as stable substances and can be
obtained readily by reacting either sulfur trioxide or sulfuryl
chloride with ammonia* On the other hand, less is known about
the reactions of either sulfur dioxide or thionyl chloride to give
the corresponding imid,es and amides of sulfurous acidf
II. Historical Developments
There is some confusion in the earlier literature regarding
the products of reaction between sulfur dioxide and ammonia.
In 1826, Dobereiner (l) described the product as a brownish-
yellow vapor which quickly condenses to a bright brown solid mass
which he called anhydrous ammonium sulfite, which is converted
into the hydrate with a trace of water.
. : t
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■ : . * :; ■ •
-22-
In a series of three papers, Rose (2,3,4) concluded that the
same product is obtained regardless of the ratio of sulfur
dioxide to ammonia. Rose found t^at the compound obtained contain-
ed equimolar quantities of sulfur dioxide md ammonia and was
either orange-red and viscous or a red crystalline material. He
named this material ammonium sulfite or anhydrous ammonium bi-
sulf itee
It was not until almost sixty years after the work of Rose
that Schumann (5) undertook his investigations of this reaction.
Schumann assumed that the products reported earlier were decom-
position products and not, in themselves, direct products of the
reaction. Working at temperatures around. 0°C., he succeeded in
isolating two products. The first, S02»NH3, was obtained using
an excess of sulfur dioxide; said, the second, S02»2NH3, using an
excess of ammonia. The former compound is a yellow powder which
he found to be amidosulf inic acid, IvHgSOOH, while the latter is
a hard, red material which he termed the ammonium salt of amido-
sulfinic acid. The amidosulf inic acid is very hygroscopic and is
also decomposed by moisture, into a white powder of variable com-
position. Ammonium amidosulf inate was also obtained as a very
viscous material •
In the same year, Divers ana Ogawa (6) found that sulfur
dioxide reacts with excess ammonia in an ethereal solution at
low temperatures to give, not a red material as the product
Schumann had obtained, but, instead a white powder. This compound
was also found to have the molecular composition S02»2NH3» It was
found to be very deliquescent; on exposure to air, it decomposed,
evolving ammonia to form an orange colored substance. Divers and
Ogawa conoluded that ammonium amidosulf inate is the first product
of the reaction, which subsequently underwent decomposition to
give diammonium imidosulf inate (6,7,8) and ammonia.
S02 + NH3 -* NH2SOONH4 -» NH(SOONH4)2 + NH,
Ephraim and Piotrowski (9), a few years later, confirmed
Schumann1 s results as to the product obtained from excess sulfur
dioxide and ammonia, but disagreed with Divers and Cgawa over
diammonium imidosulf inate as the final product of the reaction
between sulfur dioxide and excess ammonia at room temperatures.
They isolated a red material of molecular composition 2S02»4NH3
which they called tri ammonium imidosulf inate , NH4N(SOONH4) 2» They
succeeded, also, in preparing the trisilver salt, AgN(SGOAg)2.
These workers also studied the reaction between thionyl chloride
and ammonia, assuming that imidodisulf inamide , NH(SONH2)2| should
be obtained. Thionyl chloride reacts with excess liquid ammonia
to give an intensely red-colored solution; and the stable, red
residue, after the evaporation of the solution, was added to an
ammoniacal silver nitrate . solution to give the same trisilver salt
as obtained from triammonium imidosulf inate. Ephraim and
Piotrowski assumed that the imidodisulf inamide formed initially
is hydrolyzed to diammonium imidosulf inate;
NH(SONH2)2 + 2H20 -> NH (SOONH.J 2.
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-23-
They also assumed that the red color of the liquid ammonia solu-
tion is due to the conversion of imidodisulf inamide into an
ammonium salt, NH4N(S0NH2) 2,
Divers and Ogawa discovered two other compounds as products
of the decomposition of ammonium amidosulf inate. The first was
a compound 4NH3«5S02 or N^HisSsO^ (7)» They could not assign
a structure to it, but Sphraim and Piotrowski (9) assumed it to
be a chain-like molecule:
HOSO-MSO-NHSO-NHSO-NHSOGH »3H 20.
The second compound (8) was assumed to be sulf amide.
Since substituted sulfinimides can be prepared "(10) by
reacting amines with thionyl chloride:
SCC12 + RNH2-»S0NR + 2HC1,
Schenk (11) in 1942 attempted and was successful in the prepara-
tion of sulfinimide:
S0C12 + NH3 -» SONH + gHCl,
This substance was found to be unstable, however, and polymerized
at ~60°C. to a chain polymer. Schenk regarded the imidodisulf in-
amide, which -Sphraim and Piotrowski thought they had obtained, as
an ammonolysis product of sulfinimide.
Jander, Knoll, and Immig (12), utilized the reaction between
liquid sulfur dioxide and ammonia to obtain the same yellow
compound S02«NH3 previously obtained by the gas-phase reactionj
but they assumed it to be thionyl ammonium sulfite, [ (NH3 ) 2S0]S03«
Recently, Goehring and Kploumenos (13) have examined the pro-
perties of S02»NH3. The compound prepared at low temperature gave
a quantitative yield of ammonium and bisulfite ion, while the
compounds formed at higher temperature gave in addition trlthlon-*
ate ion, sulfate ion, and thiosulfate ion. Decomposition of
S02»Mi3 occurs at 80oC., yielding amidosulf onate , imidodisulf onate
N4S4, sulfur, and N2S4«
III, Imidodisulf inamide
In spite of the results of Sphraim and Piotrowski; Goering,
Kaloumenos, and Nessner (14) attempted the preparation of imido-
disulf inamide . They were successful in preparing this compound
in tv/o ways: first by reacting ammonia with sulfur dioxide,
and second, by reacting ammonia with thionyl chloride, the ammonia
used in excess in each instance.
The first reaction was carried out by bringing together the
dry gases at a temperature of -15 to ~2CoC. The orange-red
S0S»21JH3 melts at CoC. to a dark violet liquid; the melt was
extracted with nitrobenzene. The nitrobenzene extract was heated to
80^I0G°C. until all cf the ammonia had been driven off and' subsequar.
:•;; -;■.■"■: \ * •
-24-
cooling produced yellow crystals of imidodisulf inamide. From
40 g, of S03»2NH3, 0.3-0.75 g. of NH(S0ilH2)2 were obtained.
Desiring to improve the yield, the second reaction was
carried out by dropping thionyl chloride onto liquid ammonia at
-80°. After evaporation of the ammonia,, a yellowish red substance
was obtained; this product when subjected to the same treatment
as above, yielded imidodisulf inanide. Prom 25 g. of S0C12,
10-15 g. of NH(S0NH2)2 was obtained.
Imidodisulf inamide and its aqueous solution are stable. The
aqueous solution is surprisingly stable towards oxidation. It is
reduced by potassium iodide in 100^ formic acid to give sulfur
and sulfate ion, A trisilver salt, AgN(SONHAg) 2, was prepared
from the amide.' This salt possesses light yellow color and is
stable for several hours*
IV. Bibliography
1. Dobereiner, Schweigger1 s J,, 47, 120 (182$.
2. Rose, Pogg. Ann. 33, 235 (18347.
-3. Rose, ibidp, £&, 415 (1837).
4. Rose, ibid,, 61, 397 (1844).
5. Schumann, Z. Anorg. Chem., 23, 43 (1900).
6. Divers and Ogawa, J. Chem. Soc., 77, 327 (1900) .
7. Divers and Ogawa, ibid,, 79, 1099 (1901) .
8. Divers and Ogawa, ibid., 81, 504 (190?,).
9. xvphraim and Piotrowski, Ber,, 44., 379 (1911),
10. Michaelis and Storbeck, Ann. , 274, 187 (1893).
11. Schenk, Ber., 75., 94 (194'2) .
12. Jander, Knoll, and Immig, £• Anorg. u. allgem. Chem,, 252,
229 (1937) .
13. Goehring and Kaloumenos, ibid,, 265, , 137 (1950).
14. Goehring, Kaloumenos, and Messner, ibid., £64, 48 (1951).
<il* -'. ,
(\:
y \ ■ ' .:. /
■': '
•25-
LIESEGiiWG RINGS
Richard P. Heitmiller November 11, 1952
Introduction: '
When a drop of concentrated silver nitrate solution is
placed on a film of gelatin containing potassium dichromate, the
silver chromate is not precipitated continuously but in the
form of a number of concentric rings; r^iythmatic precipitates of
this kind are called Liesegang rings. These banded structures
have been observed to be of two types: those in which the rings
are separated by clear spaces, and those in which the precipitate
is separated by a band of peptized sol, i.e. alternate bands of
different colors occur in an otherwise continuous deposit.
Factors ^ffectinp: formation:3
)4,5)6>7 >8>9 jlO>ll}12)13|H>15
Much work has been done on this subject, and the ring
formation has been observed with many precipitates and with gels
of different materials. However, the nature of the gel is not
of fundamental importance since banded structures have been
obtained by precipitation of many salts in aqueous solution in
the absence of any gel-forming substance, provided there is a
restriction to movement by convection etc*, as in a capillary
tube. We can state, therefore, that although the presence of
the gel is not essential, its presence appears to facilitate
rhythmic instead of continuous precipitation, and its influence
is to some extent specific. It has been found, for example,
that beautiful rings of selenium metal are produced in gelatin
when selenous acid is reduced either by ferrous sulfate or
stannous chloride, but this tendency to form fine banded struc-
tures decreases when agar-agar is used, and no bands at all occur
in a starch gel#16 The nature of the Liesegang rings can be
altered in the same manner. Silver tungstate gives sharp banded
structures of type I in agar-agar, while in starch-gel, structures
of the type II are obtained.15 fl s
Light also plays an important role in the formation of
Liesegang rings: in general more rings are formed if the rings
are allowed to develop in the light, than if they are allowed to
develop in the dark. In the case of gold Liesegang rings, the
formation will only take place in the presence of light. However,
under conditions where silver tungstate forms sharp bands of
type I in the light, under similar conditions, but in the dark,
the number as well as the thickness of the bands are Increased
and the rings obtained are of a spiral nature.16
> ; i. s • »
"''< :- ♦ -.' <■
) .'■ ..
:a.j3
■".'. ;
., : • :• " : .
-26-
^xplanatlon: The Diffusion Theory
Prom a survey of the literature it has been observed that,
all other conditions remaining constant, (concentration of gel
and dissolved electrolyte, temperature, light, etc.) the number
of rings decrease with a decrease in concentration of the
diffusing electrolyte, and that the distance between two con-
secutive rings usually increases as the diffusion proceeds
downward. A mathematical expression has been derived for treat-
ing the ring formation as a diffusion phenomenon,17 For the
precipitation of silver chromate it has been postulated, that as
silver ions diffuse into the chromate contained in the gel,
reaction occurs, but the silver chromate remains for a measurable
time in the colloidal state and a visible precipitate does not
form. The depletion of the chromate ion concentration causes
chromate ions to diffuse rapidly into the reaction zone. Pre-
cipitation of the silver chromate might then be induced by the
accumulation of the ions of potassium nitrate, and at this point
some potassium chromate could be absorbed causing a further
depletion of chromate ion. Since the rate of diffusion is most
rapid where the concentration gradient is greatest there will be
a region in the vicinity of the precipitation zone where the con-
centration is low. It will then be necessary for silver ions to
diffuse some distance into the gel before reaching a zone in
which chromate ion concentration is sufficient for colloidal
silver chromate to be formed, at which point the cycle is
repeated until all of the silver chromate ion is exhausted*
In a recent paper on this subject it has been stated that
nickel and zinc cyanide; nickel, zinc and cobalt f errocyanide ;
barium and mercurous vanadate; cerous oxalate and carbonate
rings in gelatine, and thorium hydroxide, thallium oxide, and
uranium carbonate in agar-agar belie the diffusion theory since
in each case the distance between consecutive precipitation zones
becomes less and less with the dilution of the diffusing
electrolyte. The author does not present an alternate explana-
tion,1 s
Banded effects are frequently found in nature, and it is
possible that rhythaatic precipitation may be important in this
connection. Liesegang developed the view that the banded agates
were formed in a manner analogous to rhythmatic precipitation
in gelatin, but his theory is not universally accepted. It is
important to remember that periodic structures ara not all
necessarily due to the same cause.
Bibliography.;,
.1, R, £, Liesegang, Naturwiss., 11,, 353 (1896),
2. R, £« Liesegang, Phot, xirchiv. , 221 (1896),
3, R. j2. Liesegang, Z. Physik. Cham,, 88, 1 (1914).
.4, F. Hausmann, &, ^norg, Chem., 4JD, 110 (1904),
5, E. Hatschek, Zhur, Khim. Prom., 10, 77 (1912).
6. N. G. Chatterji and N, ft. Dhar, Kolloid 2i. , .3_1, 15 (1922).
.7. N. R. Dhar and A, C. Chatterji, J, Phys. Chem., 28,, 41 (1924).
.-.,■! ' * .-.-■,
1 .
. (
"> 'A
, . ....
«t '• »
-27-
8.. A- 0, Chatter jl and N. H. Phar ,. Kolloid Zfc, 3_Z, 2, 89 (1925).
9.. T..Orlowski, lbid»f 3,9, f 48 (1926).
10. M.. S.. Dunin and F. M.. Sohomjakin, ibid«r 39, 50 (1926).
11.. C. ,M.; Kuznanko, Ukraine Cham. J., 3, 231 (1928)*.
12. L. N.. Mukharjee and A. C. Jhatterji, Kollold ■■.., 50, 147 (1930)..
13.. S.. Veil, Compt.. rand. T 191. 611 (1930).
14.. S. Vail, ibid.., 234 T 1879 (1952).
15.- A,. C.. Chatterjl and M. C#. Hostagi, J. Ind. Chain. Soc, £8,
283 (1951) ..
16.. M. C. Has tori f ibid.., £&, 206 (1952).
17.. G. VJagen, J.. Coll. Sci., 5,, 85 (1950)..
i '
t ■ • «
- 28 -
FORCE CONSTANTS IN POLYATOMIC MOLECULES
Donn D. Darsow November 4, 1952
I. Definition and Calculation of Force Constants
A stretching force constant is a molecular constant associated with
the chemical bond between two atoms. Numerically it is the force in dynes
that would be acting to restore a bond distance if the two atoms were some-
how momentarily displaced one cm from their equilibrium distance in the
ever-vibrating molecule, assuming that the same proportionality of restoring
force to amount of bond displacement prevailed at such fictitious gigantic
distortions (4) p. 84.
Stretching force constants (also called bond or valence force constants
(4) or valence force potential constants (10)) have magnitudes in range of
0.01 x 10 to 25 x 105 dynes/cm (4). Bending force constants (4) are
associated with the forces tending to restore distortions of angles between
three bonded atoms. ^hey have different ranges of magnitudes from and are
negligible compared to stretching force constants in linear triatomic molecules,
For illustrative purposes I will consider only linear triatomic
molecules A-B-C (eg. HON, C1CN, BrCN, ICN, N20, C02, CS2, OCS) (l), (10). In
such molecules the atoms are continually vibrating along their common axis.
^he two pairs of atoms A-B and B-C each have a fundamental frequency of vibra-
tion which is dependent on their atomic masses and the forces between the
two atoms. Each vibrating motion approximates simple harmonic motion
(departing somewhat for the very small amplitudes), has a fundamental fre-
quency, and gives characteristic fundamental and overtone bands in the infra
red spectrum of the molecule (7).
The stretching force constant of the bond is defined by the following
equation (7), (ll) pp. 148, 501:
co = fundamental vibrational frequency
/,% 1 "ab f , " stretching force constant in
( 1 ) 00 * s— — - , ab / &
27Tj J_ 4. _1_ dynes/cm
m , m, : masses of the 2 atoms, in
atomic mass units.
Certain corrections for anharmonicity of the vibrations, if known, can be
applied when greatest accuracy is desired.
Force constants are calculated the most precisely from the vibrational
frequencies determined spectroscopically. Such determinations of frequencies
are not always possible however. An alternative method is to calculate the
force constant value from the interatomic distance (the bond length) using
empirical relations, such as Gordy's.
There has long been observed a definite regularity in the variation of
force constants and bond lengths in all series of similar molecules that have
been so studied, (7) pp. 453ff. One notes that in all cases the force con-
stant increases as bond length decreases, (l)(2 )(3)(7)(10). As yet no
theoretical derivation between force constants and bond length has been
advanced. But several empirical relations between force constants and bond
[.
_
1
f ,
?,
1
ab
1
1
- + —
1
m
a mb
L
,, . .,
.;
•-,..., ;.• •■ : ■•• '. ••. .'.■-.
. ■
i :■■:■■•
: v;.L,
. \t
sij ?.?.: '!" '■ ,■
: , %. vi ;.
u, ;.' ;::.:.', ij Lyii
'.;.•.■■■■•■
:■.'■■•■•'■'
,- \ C k#
i ... ., ! . ... ,. -
*-;>■; :-,y
■ , , ' .,'.. . : ■ ■ '■■■>■'■■■■
,";-■.■ :■:■ , ■ .»j \ ''' l * - r ■• !:- r ■■ ■•' :
-29-
lengths have been developed (2)(7) by comparing force constants and bond
lengths for many known molecules.
In 1946 Gordy (2) presented such a relation which has been quite suc-
cessful in many cases. This rule is formulated as follows:
\'h
<2' fab ' " \t |
3 f = stretching force constant for
"4* bond A-B in units of
+ d lO^ dynes/cm
N , = bond order
ab
o
R = equilibrium bond length in A
X ,X : electronegativities of atoms
a b A and B
c,d: constants whose value depends
on the rows in the periodio
table of the atoms A and B.
Table given by Gordy (2).
This equation gives force constant values within an average of two percent of
the spectroscopically determined force constants for all molecules to which
it has been applied (2).
The concept of bond order, II, is intimately related with bond lengths,
with force constants and with the nature of chemical bonds. The bond order
is a relative measure of the electron density between bonded atoms (2)(l2)(l3),
Otherwise stated, it is the effective number of single (those involving only
two electrons) bonds existing between two atoms. This applies to both extreme
ionic and extreme covalent bonds and also to those intermediate, where the
partial ionic and partial covalent character (hauling (6), pp. 171ff) are con-
sidered additive (2), Bonds in molecules whose atoms have their usual
valencies satisfied have essentially pure bond types. Typical pure single,
double, and triple bonds are those in HC1, CS2> and N2. These bonds have bond
orders of 1.00, 2.00, and 3.00, respectively. (See extensive table of bond
orders of Gordy (2)).
An empirical relation between bond order, K, and bond length, R, has
been developed by Gordy (3)
IT , = bond order for A-B
ab
o
/,\ TT c' ., R , = bond length of A-B in A
(3) IT . = + d' ab b
ab R 2
ab c^d': constants that have particular
values for particular pairs of
bonded atoms. Table given by
Gordy (3).
This equation yields bond orders from the otherwise determined bond lengths.
Equation (3) is frequently used to calculate the bond order value to be
inserted into equation (2) when using the latter to calculate force constants.
II. Significance of Force Constants
Force constants are admittedly a theoretical concept of pure chemistry
.
t ".- ■ J
;.'
-30-
which are mainly useful to the investigator who is interested in the forces
between atoms and the nature of chemical bonding. Force constant values
must be based ultimately on other fundamental quantities, such as bond
lengths or vibrational frequencies. They do not give us any additional facts.
They do present this experimental knowledge in new forms that give increased
internal consistency to the current meager but growing understanding of the
chemical bond.
Perhaps the most important use of force constants is aiding in the
examination of spectra. Force constants that have been reliably determined
for a particular bond can sometimes be utilized to calculate or identify the
vibration frequencies of that bond.
III. Estimation of Force Constants of Tri atomic Molecules by a Comparison
Method
There are instances where errors substantially larger than 2 o/o result
from calculating a force constant value only from the bond length* electro-
negativities, and Gordy' s equations (2; and (3) (see reference (2), Table I).
It is desirable to have a more accurate method for use when vibrational fre-
quency data are not known for a molecule, The two recent papers of Thomas,
(l)(lO), present the following method which purports to do this.
By solving certain secular equations involving the force constants of
a molecule, the force constants may be expressed in terms of a parameter, p,
which may have a large range of possible values (9)(l0). 1'hus a large num-
ber of possible sets of the force constants for the molecule are obtained.
These are most conveniently represented by graphing the p vs. f values.
If the one correct p value could be obtained, the correct value of all the
force constants would be indicated. Thomas obtains this correct p value by
calculating the force constant of one of the bonds in the molecule from the
knowledge of the ratio of bond lengths of that bond and the length of the
same bond in a similar "reference molecule", the force constant of which bond
is accurately known* The ratio of the bond orders is then calculated and
used with the ratio of bond lengths in Gordy* s equation (2) to calculate the
force constant of the bond in the molecule being investigated. The signifi-
cant feature of this method is that a comparison xvith the reference molecule
force constant value is employed.
In (10) Thomas applies this method to calculate force constant values
for C1CN, BrCN, ICN and OCS. HCN was the reference molecule used for the
cyanogen halides and CS2 was used for OCS. Vibration frequencies which have
been since determined for OCS have given the precise value of the force con-
stant for Ocs, Thomas' value checks within two percent with it.
TV, Extension to Other Polyatomic Molecules
Thomas (l) then makes an interesting extension of his method to the non-
triatomic molecules
N-C-0 *-C-S hA-HS-N3 H-C-C-B
a. a a I
H
which have the angular shapes indicated, the three heavy atoms being collinear
in each case, ^e states that "the motions of the H atoms are not coupled to
any great extent with those of the rest of the molecule and to a good approx-
imation the CH„ and H — N groups may be considered as rigid groups."
.
I,
■ •■
■ r. ■■'■•■• \ ' y
- :
'••
■~~- -"> '''■' ~ "
■ : :
•
; ■.-■
-Si-
He makes the interesting but otherwise unsubstantiated implication
that the vibrational (and force constant) characteristics of these molecules
should be very similar to their isoelectronic counterparts, respectively,
0-C — 0, 0-C — S, and 0~ Ns~ ^3* Accordingly he uses these as the
reference molecules for calculating values for their force constants, using
the f-p comparison method as in (l0). ue uses aCN as the reference molecule
for CH„CN. Strangely, Thomas (l) tabulates vibration frequencies for each of
these molecules but does not mention whether the force constant values have
been calculated from these, or, if so, whether they are in agreement with
the values resulting from his method.
REFERENCES
1. Thomas, W. J. Orville, J, Chem. Soc, 1952, 2383-89.
2. Gordy, W., J. Chem. Phys. . 14, 305 (1946)
3. Gordy, vf., J. Chem. Phys. 15, 305 (1947)
4. Linett, J. W., Quart. Revs., 1, 73-90 (1947)
Excellent Review entitled "Force Constants"
5. huckel, 1.., "Structural Chemistry of Inorganic Compounds," Vol. I, pp. 359-63,
(1950) Elsevier
Excellent overall but non mathematical exprsition of force constants.
6. Pauling, Linus, "The Nature of the Chemical Bond," p. 171, (1938) Cornell U.
7. Herzberg, G., "Spectra of Diatomic Molecules," 2nd Ed. 1950, (which is
Vol. I of "Molecular Spectra and Molecular Structure"), Van Nostrand Co.
8. Wu, C. K., "Vibrational Spectra and Structure of Molecules", 1946, Edxvard
Bros., Ann Arbor, Mich.
9. Torkington, P., J. Chem. Phys., 17, 357 (1949)
10. Thomas, W. J. 0., £, Chem. Phys . , 19, 1162-66 (1951 )
11. Glasstone, Samuel, "Textbook of physical Chemistry," 2nd Ed., 1946, Van
Nostrand,
12. Coulson, C. A., Proc. Roy. Soc. (London) A169, 413 (1939)
13. Longuet-Higgins, H. C. and Wheland, G. W. "Annual Review of Physical
Chemistry", ("Theories of Valence"), Vol. 1, 1950, p. 135."
. " ; !
i?t <•- UO
:•
: '■
~*o-
•3
METATHESIS. IN LIQUID .ARSENIC TRICHLORIDE1
Harold J. Matsuguma November 11, 1952
Introduction:
^bout fifty years ago Walden (2) recognized liquid arsenic
trichloride as an ionizing, non-aqueous solvent. He found that
binary salts and quar ternary ammonium salts are soluble in this
solvent. The recent finding that water-free arsenic trifluoride
is a solvent in which metathesis is possible has revived interest
in Walden' s observations.
Properties of Liquid arsenic Trichloride :
arsenic trichloride has many of the desirable properties of
an ionizing, non-aqueous solvent. Physical properties of the
pure solvent are summarized in Table 1. according to Raman
spectra the molecules of this compound possess a pyramidal
structure.
Table 1
Molecular Weight 181,28 g/g-mole
Melting Pt. -18°C.
Boiling Pt 130. 2°C. (2,3)
Specific Conductivity (0°) 1.4-1.6 x 10""?
-/v ~1x cm"1
Density of liquid at melting pt 2,163
Dielectric constant 12.8 (20°) (4,5)
12.35 (21°) (4,5)
Specific refractivity 0,2732
Molar volume at the melting pt 83.8
Ebullioscopic constant 6,48-7.25 (6)
Heat of formation of the liquid 71.4 Kcal/g (7)
Parachor 212.0 (8)
Heat of vaporization 69,7 cal/mole
Critical temperature 356°C.
Paramagnetic susceptibility 7,39 x 10 6
Dipole moment in benzene 2,15 (9)
Dipole moment in dioxane 3,11 (9)
Decomposition voltage in nitrobenzene 0,78 v (10,11,12)
General:
according to Walden (2), and confirmed by Gutmann (1),
the self-ionization of arsenic trichloride may be expressed by
the equation: 2^sCl3 ^ ..) AsGla + 4*sCl4""*. Therefore, sub-
stances capable of forming *bC1s or ^sCl.t~ ions in arsenic
-33-
trichloride would represent acids or bases, respectively. It
was hoped that neutralization reactions between such compounds
in liquid arsenic trichloride would lead to the formation of
hitherto unknown types of complex chlorides.
arsenic trichloride dissolves many different substances
whose solutions are much better conductors then the pure solvent
itself. The solubility relationships of various compounds and
elements in arsenic trichloride are summarized in Table II.
Reactions between the solvent and several compounds dissolved in
it are given in Table III,. Some reactions between arsenic
trichloride and organic compounds are summarized in Table IV,
Figure 1 depicts the considerable increase in the conductivity
of a solution of tetramethylammonium chloride in arsenic tri-
chloride as a function of concentration.
Table II
Solubility Relationships in arsenic Trichloride
Soluble
Less Soluble
Slightly soluble
Insoluble
(CH3)4N^SC14
[(CH3)4N]SnCl6
K2C03
KMn04
(CH3)4NSnCl5
(CH3)4NTiCJl5
K2S04
Cr203
(JH3)4NVC15
Fe3Fe(CN)6
KN03
Nb205
(CH3)4NSbCl6
[ Li CI
AgOl
Ta205
KI
NaCl
%G12
Mo03
HC1
Mi4Cl
BaCl2
W03
Rbl
KCI
BaBr2
Ti02
I CI
RbCl
VC13
CuO
cia
CsCl
v2o5
Cu
Br2
NbCl5
CrCl3
A6
I«
TaCl5
Fe203
^u.
S
p2o5
FeS04
Mg
P
KCH
Mn02
Si
*ilCl3
CuCl
Ti
TiCl4
CuCl2
Nb
SnCl4
Ta
(CN)2
Cr
C0C12
Mo
i N0G1
Pt
PC1S
Fe
SbCl3
■
V
SbCl5
W
FeCl3
Hgl2
Col 2
-sl3
-34-
Table III
Some Reactions in arsenic Trichloride
Starting Product
Reaction Product
Observations
(CH3KNCI
(CH3)4NaSC14
C2H5I
^sl3
(13)
HgO
HgCl2
With heating
BBr3
BCI3 and ^sBr3
(14)
NH3 (liq-.)
AsCl3*4NH3
(15)
N204
xis205, HG1 and N0C1
PH3
-sP3 and HG1
(16)
PI3
PG13 and ^sl3
(17)
Pb02
PbCl2
With heating
and shaking
asH3
arsenic and HG1
H23
^•s2S3 and HC1
(18)
Cr203
Heterogenous powder
KI03
Not investigated
HI
asl3 and HG1
Snl4
-si 3 and SnCl4
(19)
Table IV
Reactions of Organic Compounds
Starting Product
Reaction Product
(C2H5)NH3C1
[(G2H5)NH3]2.sCl5 (20)
(C2H5)2NH2C1
(C2H5)2NH2/iSCl4
(CH3)NH3C1
(CH3NH3)3^s2Cl9 "
(CH3)2NH2C1
(CH3)2NH2,iSCl4 "
(CH3)3NHC1
[ (CHaJaNHJa^SaCln "
Guanidine
[H2;iG(iMH)NH3]3A8aCl9 "
Pyridine
C6H5NH^SC14 "
Qui no line
(C9H7lMH)2.,s3Cl11 »
Quinoline
C9H7NH01^sCl3«2H20 (21)
-35-
Fi poire 1
25 T
Specific 20
Conductivity
x 10* 15
s>*
10 -
/
/
.01
.02
.03
.04
.05
Moles/liter
Specific conductivity of solutions of (CH3) ^N^sCl.!
in liquid arsenic trichloride at 20° C.
Gutmann prepared arsenic trichloride by the direct chlor-
ination of the metal. The initial product was yellow due to
the presence of excess chlorine. Excess chlorine was removed
by bubbling a stream of dry nitrogen through the liquid and by
repeated distillations under reduced pressure. The final product
was obtained as a colorless liquid whose specific conductivity
varied between 3 and 6 x 10~v ^t-~ix cm' i *
Isolation of Bases, from, Liquid ^rsenic Trichloride :
Compounds of the type Me ^sCl4 were considered bases in
liquid arsenic trichloride, but very few such compounds have
been reported up to this time. The dimethyl ammonium and pyridine
salts are known and two salts, Rb3-iS2Cl6 and CssAsgClc, , had been
prepared by Wheeler (22). The alkali metal chlorides were found
to be only slightly soluble in liquid arsenic trichloride.
However, the residue upon evaporation of a solution of potassium
chloride in arsenic trichloride was found to correspond to the
formula Kn.sCl4 upon analysis. When this compound was prepared
in an ignition tube at 250 o there was no increase in the yield,
showing that only the portion of the chloride which did dissolve
would form the compound.
A solution of tetramethylammonium chloride in arsenic tri-
chloride was evaporated under reduced pressure until all of the
solvent had distilled off. The residue was a pure white, non-
hygroscopic powder which was stable to temperatures up to 175°
and which had the formula (CH3) 4N;xbC14- For neutralization
reactions tetramethylammonium chloride was added to solutions of
solvo-acids in arsenic trichloride rather than adding the base
itself. The end result was the same in both cases.
-36-
Isolation of ^.cids in Liquid, arsenic Trichloride :
The compounds Ti
in liquid arsenic tri
solution.. VC14 was f
theless, a weak acid
to dissolve in arseni
to the formula -xsSbol
solution. Since that
the compound was assi
Cl.j, and SnCl4 were found to be very soluble
chloride and to exhibit acid character in
ound to be much less soluble, but never-
in the arsenic trichloride. Sb0l5 was found
c trichloride and a compound corresponding
s was isolated from the residue of such a
solution exhibited acid characteristics,
gned the formula .is Jl2SbCl6 .,
Metathesis and the Formation of Pentachloro Complexes of Te tra-
valent Titanium. Tin and Vanadium:
Figure 2 depicts the course of the conduc time trie titration
when tetrame thylammonium chloride is added to a solution of
SbCl5 in arsenic trichloride. When the mole ratio of 1:1 was
reached there was a sharp break in the curve and a subsequent
steady, marked rise in the conductivity.
30
Fig. 2
Specific
conductivity
x 104
20
,
/
10
1
I
/
" "
0
f^^^0
1.0 2.0 3.0
Mole Ratio (Base:~cid)
In the case of the neutralization of a solution of SnCl4 in
trichloride the reaction was found to proceed in two
distinct steps. (See Figure 3) The maximum point (1:1) repre-
arsenic
tine l sieps. voee figure 0/ me maximum poini, u:i; repre
sen ted the formation of the pentachloros tannate (IV) salt. The
addition of more base caused the formation of the less soluble
hexachloro complex and the subsequent decrease in the conductiv-
ity. Onoo the latter complex had been completely formed,
conductivity again increased steadily due to the presence in the
solution of the excess base ions* Gutmann actually isolated
the two compounds tetrame thylammonium pentachloros tannate (IV) and
bis( tetrame thy lammonium) hexachloros tannate (IV) by evaporating
solutions of SnCl4 and te trame thylammonium chloride in arsenic
trichloride. These compounds are light yellow, hygrscopic
solids.
-37-
30
Fig. 3
Specific
conductivity
x 10*
1
1
1 /
20
<
/
10
/
y\
*v
/
•
/S
'"V
/
0
/
V,
/
1.0
2.0
3.0
4.0
Mole Ratio (Base :^cid)
The behavior of TiCl4 toward neutralization with tetramethyl-
ammonium chloride in a solution of arsenic trichloride was
found to be exactly analogous to that of Sn^l<t • The curve shown
in figure 4 shows two distinct breaks. The first is caused by
the formation of the soluble pen tachloroti tana te (IV) complex and
the second, by the formation of the more soluble hexachloro-
titanate(IV) complex. The final sharp increase was caused by
the addition of excess base. In this case too gutmann was able
to isolate both chloro complexes. Both are light yellow,
hygroscopic solids.
Fig. 4
40
Specific
conductivity
x 10*
1
!
30
r*
1
■' \
20
s
i
10
r
■'■>'■
.--•
<<-
0
.^-- "*"*
, .
i
i
0.5 1.0 1.5 2.0 2.5 3.0
Mole Ratio (Base:^cld)
3.5
-38-
With the addition of te trame thylammonium chloride to a
solution of VC14 there was found only one break in the conductiv-
ity titration curve (Figure 5) • The increase in conductivity
was caused by the formation of the soluble pentachlorovanadate
(IV) compound. The absence of any other break in the curve
showed that no hexachlorovanadate (IV) compound was formed. The
compound tetrame thylammonium pentachlorovanadate (IV ) prepared
by Gutmann was a dark reddish-violet powder, readily soluble in
water to give a bluish-green solution. This compound could not
be obtained completely free of arsenic. Since VC14 is a *eak
acid in arsenic trichloride, the presence of arsenic in the
product was due to solvolysis.
Fig. 5
30
Specific
conductivity
x 104
/
20
/
10
. /
0
/
t,
• , ... . , i
1.0 2.0 3.0
Mole Ratio (Base:,.cid)
Summary:
Water-free arsenic trichloride is an ionizing, non-aqueous
solvent in which metathetical reactions are possible. Such
reactions produce complex chlorides of the elements titanium,
tin, antimony and vanadium. The compound potassium tetrachloro-
ar8©nite(lll ) was also prepared by reaction in liquid arsenic
trichloride. TiCl4 , SnCl4 , VC14 and Sb^l5 were found to be
acids in arsenic trichloride. - solution of tetrame thylammonium
chloride in liquid arsenic trichloride provided a strongly
basic solution.
i-
-39-
Bjbllopcraphy:
1. Gutmann, V,; z. anorg. u. allgem. Chem. 266 . 331 (19511.
2. Walden, P.: Ibid. 25, 210, 214 (1900).
3. Biltz, W. and Meinecke, £.: Ibid, 13.1, 16 (1923).
4. McCusker, P. -. and Curran, B, C#: J. ^m. Chem. Soc . 64,
615 (1942) .
5. Schlundt, H.: J. Phys . ^hem, 5,, 512 (1901).
6. De Kolossowsky, N.: J. Chin. Physique 2£, 356 (1926).
7. Thomsen, J.: Ber. 16, 39 (1883).
8. Henly, W# J. R. and Sugden, S.: J. Clhem. Soc. London 1929 .
1060.
9. Smith, J. W.: Proc. ^oy. Soc, London ..136, 262 (1932).
10. Ratten, H. E. and M0tt, W. R.: J. Phys. Chem. 8, 153 (1904).
11. Finkelstein, W.: 2, physlk. Chem. 115if 303 (1925).
12. Ualden, P.: Ibid, .43., 42.0-1, 436 (1903).
13. Ko'hnlein, B. : Liebigs ^rm. Chem. 225, 176 (1885).
14. Tarible, J.: Corapt. rend. 132, 206 (1901) .
15. Besson, .,.; Ibid, 110, 1258(1890) .
16. Jancvsky, J. V.: Ber., 8,, 1636 (1875).
17. Karantassis, T.: Corapt. rend. 182 . 1391 (1926).
18. Biltz, W. and Keunecke, £.: Z, anorg. u. allgem. Chem., 147,
173, 187 (1925).
19. Karantassis, T.: Corapt. rend. 182,, 699 (1926).
20. Petzold, W„: Z. anorg. u. allgem. Chem. 214, 355 (1933).
21. Dehn, W. M#: J. ..m. Chem. S0c. 4j3, 375 (1926).
22. Wheeler, H, L#: Z, anorg. u. allgem. Chem. 4., 452 (1893).
23. Me 11 or, J. W., 0/, Comprehensive Treatise of Inorganic and
Theoretical Chemistry", Longmans, Green and Co., New York,
1929. Vol. £, pp.- 239-245.
1 '" "■ ,
-40-
THiS RHENIDE OXIDATION ST..TE
Robert L, Rebertus November 18, 195?
I . Chemical Studies.
In 1937, Lundell and Knowles (l) found that if a cold
solution containing potassium perrhenate and sulfuric acid is
passed through a Jones reductor, the rhenium(VII) is reduced to
rhenium(-I), Eight equivalents of oxidizing agent were required
to oxidize one equivalent of the rhenide to perrhenate* These
investigators point out that rhenium has lower oxidation states
analogous to the halogens, and they speculated that if a solution
containing rhenium (-1) and sulfuric acid is warmed, hyporhenous
acid (HReC) and sulfur dioxide should be produced ♦
The reducing effects of lead, cadmium, bismuth, and zinc
amalgams on potassium perrhenate in sulfuric acid and hydrochloric
acid media were studied by 0. and F. Tomioek (2) « Bismuth amal-
gam in 80$ sulfuric acid yielded blue solutions in which Re (VII)
was reduced by 2.2 valency units. In 40-50$ sulfuric acid
solution, brown solutions of Re (IV) were obtained. Lead amalgams
in 10$ hydrochloric acid and cadmium amalgams in 20-25$ sulfuric
acid gave a solution containing Re (IV). Only saturated zinc
amalgam in 10$ hydrochloric acid reduced the Re {VII) to Re(-I)»
Lingane (3) has confirmed the reports of Lundell and Knowles
(l) and of Tomicek (2)* He prepared rhenide solutions by reducing
potassium perrhenate in a Jones reductor. The reduced solutions
were run directly from the reductor into an excess of standard
eerie sulfate solution under an atmosphere of nitrogen, and the
excess eerie ion was back titrated with a standard ferrous
ammonium sulfate solution, using Ferroin as an indicator. Table
I shows the results obtained,
Tx.BLE I
Oxidation State of Reduced Rhenium Solutions
KRe04 Volume reduced 0,01 N Equiv. Ce(IV)
Millimolar solution, cc, Ce(S04)2f cc, per mol total
Re
3.26 8.3
5.28 7.6
12.1 8.0
23.0 7.8
0.105
37.6
.209
33.3
.418
36.4
.836
35.2
-41-
The analogy between rhenide and h allele ions suggested by
Lunclell and Knowles (1) was extended by Rulfs and living (4),
who reasoned that thallous rhenide, like the thallous halides,
should be relatively insoluble,
concentration was prepared. This
thallous chloride at 25° and then
nitrogen for ninety minutes at 0°
with thallous chloride, separated
the crystals were not analyzed,
rhenide. The solubility of the
. rhenide solution of known
solution was saturated with
was allowed to stand under
White crystals, isomorphous
from the solution, although
they were presumed, to be thallous
thallous rhenide estimated from
the known rhenide concentration in the presence of excess thallous
ion was approximated to be at least 39 mg, per 100 ml, of 2.4 N
hydrochloric acid at 0°, The solubilities of the thallous halides
in mg. per 100 ml. water at 0° are: T1C1, 161; TlBr, 22; Til, 2.
Maun and Davidson (5) attempted to prepare more concentrated
solutions of Re(-I) than those prepared by Lundell and Knowles (l)
(10*~4 F) and Lingane (3) (8 x 10 4 F) by reducing hydrochloric
acid solutions of perrhenate in a Jones reductor. They were
unable to reduce quantitatively He (VII) to He(-I) if the initial
Re (VII) concentration was greater than 0,001 F. Table II shows
the results of some of their experiments.
TiBLz II
apparent average Oxidation Numbers
of Reduced Rhenium Solutions
HRe04
HC1
F
F
0.00097
4
0.00097
4
0.00097
4
0.00242
6
0.00242
6
0.00483
4:
0.00483
4
0.00504
4
0.00970
4
Time of Reduction
minutes with
Ce(S04)2
average Oxidation No,
wi th wi th
KI03 I a
3-5
3-5
3-6
2-3
2-3
3-5
3-5
3-5
3-5
0.1
-1,4
-0.9
-0.3
0.5
2.4
•0.2, 0.1
0.6, 0.7
2.9
-0,4,0.4
0.8, 1.0
0.6,0.8
0.7,0.8
The average oxidation numbers listed, were claculated en the basis
of the initial formality of the perrhenic acid, It is believed
that the incomplete reductions resulted from the precipitation of
rhenium in the reductor. The resulting solutions would contain
pure Re(-I) but at a lower concentration than the initial Re (VII)
concentration. The solutions obtained from the reductor were
colorless which f act supports the explanation of the incomplete
reductions •
-42-
The possibility of concentrating Re(-I) by volatilization
or by evaporating the solutions was also studied by Maun and
Davidson (5). No volatility of rie(-I) resulted, and a brown
solid residue, which resulted from volatilizing the hydrochloric
acid and water from the rhenide solution, was shown by analysis
with Ce(IV) to contain rhenium in an intermediate oxidation
state •
Hignly colored intermediate oxidation states were formed
when rhenide was treated with perrhenate, copper sulfate, or
oxygen (5). From spectropho tome trie studies of the reactions of
Re(-I) with perrhenate and oxygen, and from analyses on the
reaction product of Re(-I) with copper sulfate or oxygen, the
first oxidation product was concluded to Re(l), which is followed
by the formation of Re (IV), It is to be noted that the compound
Re pO has been prepared (6) and that s solution possibly containing
Re(I) from the reduction of Re2Cl6 has been described (7).
Recently, Kleinberg (8) and his coworkers succeeded in
isolating a potassium rhenide-potassium hydroxide mixture.
Potassium perrhenate (0%003 M) in e thylened.iamine-water solution
(9,8 g. water per 100 ml* solution) was reduced by potassium
metal (4 g. per 100 ml. solution) in a closed system under
nitrogen, aqueous solutions of the solid formed were prepared and
titrated with potassium dichromate. Under the conditions des-
cribed 55% of the potassium perrhenate reacted to form potassium
rhenide.
When an aqueous solution of the solid mixture of potassium
hydroxide and potassium rhenide is treated with thallous ion, a
white precipitate, presumed to be thallous rhenide forms. This
rapidly decomposes to thallium metal and perrhenate ion,
II., JlLxSUTROCHaMIC.iL STUuI^S
The oxidation states to which perrhenate ion is reduced under
various conditions can be determined from polarographic waves by
means of the Ilkovic equation,
A& = 607 n US' m* ts
c
where id is the diffusion current (microamperes) , D is the
diffusion coefficient (cm,3 sec.. 1) , c is the concentration
(millimoles per liter) of the reducible substance, m is the rate
of flow of mercury from the dropping electrode (mg. sec. x) ,
t is the drop time (sec.), and n is the number of electrons
involved in the reduction.
Lingane (9) found that perrhenate in 2-4 N hydrochloric acid
solution reduces to Re(IV), In neutral unbuffered solutions con-
taining 2 N, potassium chloride a double wave is produced,. The
first part of this wave (E^ = 1,41 v, vjg,, S.C.tC.) is due to the
reduction to the rhenide ion. The data in Table III verify the
eight-electron reduction. From the Ilkovic equation the theo-
retical value of id/c for n = 8, corresponding to the reduction of
perrhenate to rhenide ion, is 40.3 ffiicroamperes/millimole/liter
for the particular capillary used,.
-43-
T.,BL^ III
Diffusion Current of Perrhenate Ion as a Function of Concentration
in 2 I Potassium Chloride.
Re04~
id
id/c
millimolar
Microanvcc
jres
Microampere s/millimol/li ter
0.0192
0.97
50.5
0.0487
2 #38
48.8
0.0703
3.39
48.3
0.192
7.85
40.9
0,311
12.5
40.2 (Theor. (n=8)=40
Eulfs and living (10) also found that the first reduction
wave of perrhenate in 4 N hydrochloric acid is a three -electron
reduction. To confirm that the reduction product is (RelV) , they
prepared potassium hexachlororhenate (IV) and found that it
reduced at a potential in good agreems.V$ with the double reduction
wave of the perrhenate ion. The eight-electron reduction at
-1.41 volts in neutral solution reported by Lingane (9) was
confirmed.
Geyer (11) has measured the half -wave potentials of the steps
in the reduction of perrhenate in neutral, alkaline, and acid
solutions. In alkaline, neutral, and very weak acid solution,
he believes that the reaction proceeds in the steps
Re (VII) = Re(0) - Re(-I)
In high acid concentration, he finds three waves corresponding
to the reduction steps
Re (VII) - Re (IV) - Re(0) - Re(-I)
The waves obtained from the reduction of perrhenate ion are very
poorly defined, and Geyer has chosen to interpret a slight
inflection in the wave as the reduction to Re(0). However, his
results are in fair agreement with the results of previous
investigators.
Lingane (3) investigated the oxidation of the rhenide ion
at the dropping mercury electrode. The oxidation process le
described by the five steps listed in Table IV.
T.3L£; IV
Stepwise Oxidation of Rhenide Ion at the Dropping
Mercury Electrode
Step
Reaction
E-g., volts v£. saturated
calomel electrode
a
He -h
Re++-
Re+++
Re 5 -
Re++ + 3e~
» Re+++ + e"
► Re ' 5 + 2e
Re+7 + 2e~
Mercury dissolution
-0.54
-0,34
-0.07
-44-
That He(II) is the first oxidation product is not substantiated
by the chemical evidence given above .
Lindane (3) suggests that since rhenium (0) has the ground
state configuration,
5ss5p65d56ss
one might expect either of the following configurations for the
rhenide ion:
A, 5s2 5p6 5d6 6s2 B# 5s25p66s26p6
Since the rhenide ion has halide-like character, Lingane suggests
that configuration B is the more likely. This configuration
requires the promotion of all 5d electrons and makes for instabil-
ity. (Pauling (IS) proposes that the structure of Re(-I) is like
platinum(II) with four coordinated water molecules at the corners
of a square . )
The oxidation of the rhenide ion has also been studied
polarographically by Rulfs and Slving (13) • although these
investigators obtained five steps in the anodic wave, they believe
the first oxidation product to be Re(I). They further postulate
that the oxidation of Re(-I) to Re (I) is thermodynamically
reversible. On the basis of this postulation, Rulfs and Elving
present evidence for the ability of the rhenide to form complexes^
In hydrofluoric acid medium the half -wave potential for the first
step is -0*426 volts, but more negative values are obtained in
the presence of chloride, bromide, iodide, hydrochloric acid-
pyridine , hydrochloric acid-thiocyanate, and e thylene diamine ,
indicating that the reductions proceed with greater difficulty.
R^PiJR^NCES
1. G.ji.F. Lundell and H.B. Knowles, J. Research Nat. Bur. Stand-
ards 18, 629 (1937),
2. 0. Tomicek and F. Tomicek, Collection Czechoslov. Chem. Commun.,
11, 636 (1939) •
3. J.J. Lingane, J. M, Chem. Soc, 64, 2182-90 (1942).
4. C.L. Rulfs and P.J. living, ibid.. , 72., 3304 (1950).
5. jS.K. Maun and N. Davidson, ibid.. 72, 3509-14 (1950).
6. R.C. Young and J.W. Irvine, ibid,. , 59, 2648 (1937).
7. W, Noddack and I. Noddack , Z. anorg. allgern. Chem., 215,, 129
(1933) ,
8. S« Grlswold, J. Kleinberg and J.B. Bravo, Science 115, 375-6
(1952).
9. J.J. Lingane, J. ~m. Chem. Soc. 64, 1001-7 tl94g) .
10. C.L. Rulfs and P.J. living, ibid., 73., 3284-86 (1951).
11. R. Geyer, 2. anorg. Chem. 263, 47-62 (1950).
12. L. Pauling, Chem. dng. News, 25, 2970 (1947).
13. C.L. Rulfs and P.J. living, J. -m. C^iem. Soc., 73_, 3286-92
(1951) .
t *
-45-
H-LOGiN CUTIONS
L# K. Diamond November 18, 1952
Introduction:
The electronegativity of the halogens ie continually
stressed, and it is pointed out that the chlorine atom, having
the greatest affinity for electrons, can displace bromine and
iodine from solutions of their salts. The electropositive
character of the halogens, however, is frequently overlooked.
The electroposi tivi ty of the halogens increases with increasing
atomic number, and the risultant positive character can be
observed either in the polarization or polar izability of halogen
compounds in which a positive charge becomes centered on the
halogen atom, or in the ability of the halogen to form positive
ions. With chlorine, electropositive character amounts to little
more than measurable polarity in covalent bonds. With iodine,
however, electropositive character is so increased that positive
iodine ions can exist in simple or complex compounds. Studies
of the chemical properties of astatine by tracer techniques
indicate that the trend would probably extend to this element
(1) . The subject of positive halogens has recently been
reviewed by Reeve (2) ,
Trivalent Iodine:
Partington and Bahl (3) observed that ozone will convert
iodine into a compound I4O9 , or I(I03)3f This same compound
has been prepared by the treatment of iodic acid with phosphoric
acid (4). Ficter and Stern (5) prepared I(C2H302)3 by oxidizing
iodine with fuming nitric acid in the presence of acetic anhy-
dride. From the respective halogenated acetic acids, I(GH2C1C02)
I(CHC12C02)3, I(CC13C02)3, and I(CH2BrC02)3«I(l03)3 are obtained
whereas in the presence of phosphoric acid the normal phosphate
IP04 is formed. These compounds are unstable and hydrolyze in
accordance with the equations of the type:
5IP04 + 9H20-»I2 + 3HI03 + 5H3P04
When a saturated solution of iodine triacetate in acetic acid
is electrolyzed, iodine is quantitatively liberated at the
cathode. The cation I+++ reacts with aromatic nuclei in the
following manner;
I+++ + RH -> RI++ + H+
I+++ + 2RH -> R2I+ + £H+
Univalent Halogens.:
The treatment of a solution of iodine in a non-aqueous
solvent such as chloroform with silver perchlorate or acetate
-46-
forms the unstable salts IC104 and IC2H302 (6). When silver
nitrate is added to an alcoholic solution of iodine, IN03 is
formed.
Kikindai and coworkers (7) have recently indicated the
possibility of the fixation of the cation Br and I bv the
cationic exchanger ^mberlite IR100H , Using a solution of iodine
or bromine in absolute alcohol, Kikindai was able to prepare
solutions of compounds of Br and I • «■ typical procedure was
to pass a solution of (X.5 g. of iodine in 100 ml. of absolute
alcohol through the -*mberlite IR100H »
H + Res~ + I s -> I^Res~ + HI
The formation of HI caused a measurable decline in the pH of the'
solution. The amount of positive iodine in the resin was deter-
mined by passing a saturated solution of KI through the exchange
column and titrating the liberated iodine with thiosulfate*
I+Re s" + KI -> K+Re s~ + 1 2
approximately 0,02-0,03 grams of I can be fixed upon the resin
by this method, and solutions of I2S04 and IN03 were obtained
by passing a solution of H2S04 and HN03 - in absolute alcohol
through the exchanger containing the positive I+ ion. With this
procedure, the salts of positive iodine and bromine can be pre-
pared from the corresponding acid if the solutions are completely
anhydrous and the activity of the H ion is great enough to
effect the change,
_+
I may be stabilized by coordination with pyridine, and the
resulting compounds may be isolated in the crystalline state •
-udrieth and Birr (8) postulated the formation of I (py) from
conductivity measurements of solutions of iodine in pyridine,
and Kleinberg (9) has presented spec tropho tome trie evidence for
the existence of I(py) in solutions of iodine in pyridine.
Carlsohn (10,11,12) prepared a series of salts of the hypo-
thetical bases I(py.)0H and I(py)20H. These salts are prepared
by treating the silver or mercury salt of the necessary acid
with the calculated amount of iodine and a slight excess of
pyridine in a solvent such as chloroform, according to the
equation:
~g-^n + py + 1 2 — ► I(py)-in + ^gl
Crystalline compounds prepared include l(py)N03 I(py)2N03,
I(py)aC104, I (py) sC2H3C2, and a series of l(py)+ salts of some
thirty organic acids prepared by Kleinberg ©t al (13) • The
chemical reactions of the salts indicate the halogen is positive
and univalent. When dissolved in sodium hydroxide containing
potassium iodide and then acidified, they all liberate free
iodine.
r i ,"i X jii
. y !
-47-
The reactions of I (py) 2N03 are sirnmarlzed below. The other
salts behave similarly.
1, The iodi-pyridine salts hyr'rolyze slowly in water.
Kpy)2NC3 + H0H-*I(py)0H + pyHN03
5l(py)0H -> 21 2 + pyHI03 + 2H 20 + 4py
2. Upon treatment with sodium hydroxide the bases
I(py)OH and I(py)2OH are liberated* These immediately
revert to their respective anhydrides, which can
be isolated.
2l(py)20H-*I(py)-0-I(py) + HOH
3, The salts of iodine react immediately with phenol
to produce the iodophenols.
C6H5OH + 3l(py)2N03 -» C6H2I3OH + 3HN03 + 6py
4. The salts of the halogens are fixed by addition to
a double bond,
. ^2 >. y GH2 x
CH2 CH2 CH2 CHON02
I M + l(py)2N03 -+ I I + 2py
GH2 CH CH2 CHI
V GH 2 y v CH 2
In the case of an iodine salt with bivalent anions,
as in the case of iodine succinate reacting with
cyclohexene, the following product is obtained (14).
^GH2 / GH2 N
CH2 CHGO2CH2GH2CO2GH CH2
, I II
GH2 CHI ICH CH2
^CHs^ ^CHsy
5. The salts of iodine in a chloroform solution will
dissolve the noble metals.
3l(py)3N03 + ^u=-*au+++ + 3N03"~ + 3/2 I2 + 6py.
This confirms Finklestsin' s conclusion that the I+
ion should be placed with the noble metal ions in
the electromotive series (14) ,
6, Electrolysis of I (py) 2N03 in chloroform or methanol
produces iodine at the cathode.
• {
-48-
Methyl substituted pyridines such as P-plcoline, 2-6-
lutidine, 2-4-lutidine , and 2-4-6-collidine yield similar 1
compounds. By similar procedures the bromine compounds Br(py)2
N03 and Br(py)2C104 and the chlorine compound Cl(py)2N03 have
been prepared. (15,16)
Evidence for the existence of a uniposltiva bromine cation
is advanced from studies of the brominating strength of solutions
of aqueous bromine, alkaline hypobromite, and acid HOBr (17,18).
The relative brominating power observed experimentally is acid
HOBr>) bromine > alkaline 08r~~. The enhanced brominating strength
of acid KOBr results from the formation of a positive bromine
cation, , ,
HOBr + H -* Br + H -0 and
HOBr + H+ -> (H2OBr)*
attempts to prove the existence of cationic bromine in
acidified hypobromous acid solutions by electrolysis have not
been completely successful (18). The existence of Cl+, probably
in the form of (H20C1) + has been recently deduced from studies
of the chlorinating power of acid solutions of hypochlorous
acid (20) . The hydrated form of Cl+ is not formed at as low
acidities as (H2OBr) + , since H0C1 is a weaker base than HOBr by
a factor of IO .
With fluorine electropositive character is reduced to a
minimum,* however, Cady (21,22) has prepared some so-called '
"hypof luorites" , FN03 and FC104. These are gaseous compounds
in which the fluorine is linked, by a covalent bond to oxygen.
Fluorine in the "hypof luorites" has a valence of +1, but there
is little likelihood that any ionic forms of the acid or its
derivatives will ever be isolated because of the covalent nature
of the molecular bonds.
Bibliography;.
1. G. Johnson, R. Leininger, and jd). Segre, J. Chem. Phys., 17,
1 (1949).
2, K. Reeve, Revs. Pure nppl, Chem. , 2,, 108 (1952).
3* J. Partington and R. Bahl, J. Chem. Soc, 1258 (1935).
4, F. Fichter and H. Kappeler, Z, anoi g. Chem., 91., 134 (1915).
5. F. Fichter and S. Stern, Helv. Chim. -.eta, 11, 1256 (1928).
b. L. Birkenbach and J. Goubeau, Ber. , 65., 395 (1932).
7. T. Kikindai, Bull soc, chim., 18, 799 (1951); Compt. rend.
23,2. 1110, (1951); 232, 1840 (1951).
8. L.F. Audrieth and E, Birr, J. am, Ghem. Soc., 55., 668 (1933).
9. R. Zingaro, G. Vender Werf and J. Kleinberg, J. ^m. Chem.
Soc. 73, 88 (1951) .
10. H. Carlsohn, Uber eine neue Klass von Verbindungen des positiv
einwertigen Iods. Verlag J, tfirzel, Leipzig (1932).
11. H. Carlsohn, -ngew. Chem., 46, 747 (1933).
12. H. Carlsohn, Ber., 68B, 2209 (1935),
13. R. Zingaro, J. Goodrich, J. Kleinberg, and C. Vander Werf
J, rim, Ghem. Soc. 71,, 575 (1949).
-49-
14. V. Finkelstein, Z. physik. Chem., Ig4, 285 (1926).
15. M. Uschakow and W. Tohistow, Bar., 68B. 824 (1935).
16. M. Uschakow and W. Tchistow, Bull. soc. chim, , 3_, 2142 (1936)
17. W. Wilson and P. Soper, J. Chem. &ca., 3376 (1949).
18. D, Derbyshire and W. Waters, J. Chem. Soc, 564 (1950).
19. P. Korosy and G. Szekely, Nature, 168,, 77 (1951).
20. u. Derbyshire and W. Waters, J. Gnem. Soo.. 73 (1951).
21. G. Cady, J. ^m. Chem, Soc, 56, 2635 (1934).
22. Heport of Symposium: Unfamiliar Valence States of Familiar
Elements, Chem. ang. News, 28, 1478 (1950).
-50-
REACTI01TS OF THE ITITROSYL 101!
M. Km Snyder November 25, 1952
A. Introduction
In 1909, Hantzsoh suggested the existence of the nltrosyl
Ion, N0+, when he obtained anomalously low values for the cryo-
scoptcally determined molecular weight of nltrosyl sulfuric acid
(12). Later, the existence of the nltrosyl ion was proved by
An^s and Leckie (7) who studied the Raman spectra of nltrosyl
sulfuric acid. Recently, Addison and coworkers (5) and Partington
and '.Jhynes (18) investigated dinitrogen tetroxide. They concluded
that in certain instances dinitrogen tetroxide reacts as though
it were nltrosyl nitrate,
Addison recognized that dinitrogen tetroxide can dissociate
In three ways: n
M204 ^ N02 + N0a
iT2o4 ^=^ no3+ + nos~
n3o4 ^=^ no+ + nc3~"
The first of these dissociations is the familiar thermal reaction.
The second represents the apparent way in which dinitrogen tetrox-
ide reacts with covalent compounds, such as organic amines. The
third shows the type of dissociation involved when dinitrogen
tetroxide reacts with ionic compounds, such as amine hydrochlorides.
Actually, Angus, Jones, and Phillips (6) have shown that the liquid
dinitrogen tetroxide does not undergo self-ionlzation. Rather,
they picture the molecule in a polarized state, such that under
the proper environmental conditions it reacts as though it were
ionized.
For the sake of convenience, a solvent system may be developed
using dinitrogen tetroxide as the solvent and considering it to be
nltrosyl nitrate. In this way the reactions of nltrosyl compounds
may be classified in a familiar manner.
E. Neutralization
In the nltrosyl nitrate system, the nltrosyl ion is the
acidic species and the nitrate is the basic species.. '.Jhen a
nltrosyl halide reacts with a nitrate, nltrosyl nitrate is pro-
duced, along with a metallic halide.
H(II03)X + xN0Cl->IIClx + xH0N03
This reaction has been run for the chlorides of zinc (4), silver
(5), thallium (17) and lead (13).
•
-51-
C. Solvolysis
Addison, Conduit, and Thompson (3) have shoim that amine
hydrochlorides undergo normal solvolysis with nitrosyl nitrate to
give the amine nitrates and nitrosyl chloride, according to the
e cuation
R4ITC1 + N0?J03-*R4I!N03 + NCC1
The reaction ta!:es place *rhen R is methyl, ethyl, methyl and
hydrogen, or ethyl and hydrogen. This is a food method for the
preparation of the amine nitrates.
Certain metal nitrates are difficult to prepare in the
anhydrous state. If the chloride undergoes solvolysis with
nitrosyl nitrate, this affords a relatively easy method for the
preparation of the anhydrous nitrates. Addison and coworkers (3)
have worked this out for anhydrous zinc nitrate:
ZnCl2 + :iO:T03 F11^ Zn(lT03)a J^I^Il (H022n(NO3)
+
N0C1
-"Va/"*heat v'™ a/4
As the equation indicates, the end product is actually the complex
dinitrosyl tetranitrato zincate (II ) , which decomposes to the
nitrate upon heating. A similar reaction with titanium tetra-
chloride ^as not very successful due to undefined side-reactions
(18).
The most commonly applied hydrolysis reaction takes place in
the Glover to^er of the lead chamber process for the production of
sulfuric acid (30). The reactions take place as follows:
2302 + :T0 + :T02 + "20 + 02 -* SNOHSO4
2NQHS04 + H30 -> NO + NO 2 + 2H3S04
The first step of the hydrolysis of the nitrosyl sulfuric acid
probably produces nitrous acid, i.e. NO* OH which, in the presence
of acid, decomposes into nitric oxide and nitrogen dioxide.
The acid salts of diethyl amine undergo solvolysis to give
diethyl nitrosamine (o).
PaHaJpfHNOa + IT0II04 -♦ PaK J|-N0 + 2HN03
PsH^jgH-HCl + IT00H -> p8HjN-N0 + HOI + H20
(C3Hb)1H*HN03 + I-TOCl -* (C2H^!-N0 + HC1 + KN03
D. Solvation and Desolvation
Addison and Conduit (l,2) have shown that the diethyl nitros-
amine will solvate the nitrosyl ion.
NOI!03 + (CaHjg-NO -* [N0p2(C2H.s)N-N0]N03
(NO)2Zn(N03)4 + fcsH5^N0->'[N0-2(C2H^-N0]2Zn(i!03)4
2
Although they realized that they were able to prepare nitrosyl
compounds, the early investigators did not know the exact nature
of these compounds. As a result, they reported then as addition
compounds:
MxXy + rNOX->KxXy«rNOX
At the present, such compounds are usually written as coordination
compounds:
IIxXy + rNOX
(uo)riixx
r^x^r + y
Many of the complex halides of metals have been prepared in this
manner, using the appropriate nitrosyl halide. Table I lists these
compounds with the metal salt from which the3r were derived.
Table I
ilXAy
Old Notation
A1C13
Al 01 3 •11001
AsF5
AsFgTTOF
CuCl
Cu 01 '17001
FeCl 3
Fe CI 3 -HO CI
GaCl3
GaCl3'N0Cl
HgCl3
HgCl2 •1T0C1
InCl3
InCla-IIOCl
MnCl3
MnCla'NOOl
?F5
PF5«*T0F
PdCla
PdCl2'2:T0Cl
ptca.4
?tCl4'2TT0Cl
SbCl5
SbCl5»lT0Cl
SbF5
SbFs»XI0F
SnCl4
3nCl4*2I!0Cl
TiCl4
TiCl4*2N0Cl
TlOls
TlOla-NOCl
ZnCl2
ZnCl2-NOCl
New Notation
Investigators
N0A1C14
11,17,18
N0AsF6
20
NOCuCls
8,9,17
N0FeCl4
17,27
N0GaCl4
17
NGHgGl.
17
C(N0)2Hg2Cls]
I!0InCl4
17
NOHnCla
8,17
[(N0)2Nn2Cl6]
IJOPPe
14
:io)2paci4
(N0)2PtCle,
18
18,27
NOSbOle
22
NOSbFfl
8
(NO)aSn01a
17,27
(N0)aTiCl6
17
N0T1C14
17,27
N0ZnCl3
17,27
[(N0)2Zn2Cl6]
Some of the modern notations in Table I indicate odd co-
ordination numbers. The alternative structures in brackets are
suggested on the basis of a de solvation reaction studied by
Addison and Lewis (4). Their work may be illustrated by the
following scheme :
NO
S/"
NO a*"
3
o3
NOCl^. 0°T 1 hr.
2N0C1
}
NO
NO
01 ,01
Zn
CI' , N03
jo"
m ^
cr ^ or 01
or
vci/
sci
-53-
In the light of there reactions, it seems likely that the zinc,
mercurj^, and manganese compounds in Table I are actually the
dimers.
E. Other Iletathesls Reactions
In certain instances, dinltrogen trioxide reacts like a
nitrosyl compound. From inspection, it may be called either
nitrosyl nitrite or nitrosyl oxide. Either formulation will pre-
dict the products .of the reactions with selenic acid (15), tetra-
fluoroboric acid (9,^8), and fluorosulfonic acid (13).
'Joolf (29) carried out some reactions in liquid bromine tri-
fluoride which undergoes self-ionization:
23rF3 -> BrF3+ + BrF3*F~
Typical reactions take place with metal compounds and nitrosyl
chloride in bromine trifluoride:
MxXy + r2rF3-* xlIFr + x(BrF2)xX
x
N0C1 + 3rF3 -^ NOP + BrF2Cl
ZNOF + HFr -> NOI-IFr
x
T + Z
Table II shows a few of the complex compounds which Woolf has
prepared in this way.
Table II
Reactant Product
SnCl4 (NO)2SnF6
As203 NOAsF6
SbP03 ITOSbFe
b2o3 :tobf4
?Br5 NOPFe
G-e02 (NO)aGeFe
SnF4 (NO)2SnF6
Si02 (N0)5iFe
(ijO)2S207 1!0S03F
In addition it was found that it is possible to replace the
coordinated chlorides in a complex compound with fluoride by
treating the complex with bromine trifluoride. So if a metal such
as gold is dissolved in nitrosyl chloride to f'iv0 nitrosyl tetra-
chloro-aurate (ill ), treatment with bromine trifluoride produces
nitrosyl tetrafluoro aurate(lll).
In Germany, Seel and coworkers have prepared nitrosyl com-
pounds in liquid sulfur dioxide, a non-solvolytic solvent for these
reactions (22). By combination of tetramethyl ammonium nitro-
prussate and nitrosyl hexachloroantimonate (V) , they have obtained
a mixture of nitrosyl tetramethyl ammonium nitroprussate and
-54-
dinitrosyl nitroprussate (2S). Using the same starting nitrosyl
compound with tetramethyl ammonium hexafluophosphrte, they obtain-
ed the nitrosyl hexafluophosphate without having to handle nitrosyl
fluoride or "bromine trifluoride (24).
F. Oxidation-Reduction
Active metals will react with water to produce hydrogen and
a base. In a similar fashion zinc (4), the alkali metals (6),
mercury (5), and copper (5) will react with nitrosyl nitrate to
produce the metal nitrate and nitric oxide. If an acid, such as
nitrosyl chloride, is used, in addition to these metals, iron (5),
tin (5), thallium (5), indium (5), gallium (5), gold (18), and
platinum (18) will react.
The intensity of the color of nitrosyl halides increases as
the size of the halide increases. In the same way, the degree of
dissociation of the halides increases with increasing size of the
halide. In fret, nitrosyl iodide has never been prepared. Seel
and coworkers (23) have investigated the reaction of nitrosyl
compounds with iodides. In all cases the liberation of free iodine
was observed:
• 2N0X + 2X1 -* NO + ila + 2KX
ICI could be any convenient iodide. They found, also, that the
presence of a very slight number of nitrosyl ions could be detect-
ed in this way. Even ethyl nitrite, which didn't react itself,
did react when dissolved in acidic solvents such as sulfur dioxide,
sulfur trioxide, boron trichloride, or hydrogen chloride. This
reaction may serve as a test for the nitrosyl group.
Similarly, Seel and coworkers (25) studied the reactions of
nitrosyl compounds with azides. Instead of the very unstable
nitrosyl azide, they obtained nitrogen and nitrous oxide. Like-
wise, this reaction can 3erve as a test for the nitrosyl group.
Seel (2l) found that the salt like nitrosyl compounds reacted
violently with water to produce nitric oxide and nitrogen dioxide.
This reaction proceeded according to thc ecuation:
NO+X~ + HOK -> NO*CK-> N0N03 -> NO + N0;
+ +
HX HQH
However, it was found that the coordinated nitrosyl group in
nitrosyl nitroprussate did not react in this way.
G-. Summary
Assuming the existence of the nitrosyl ion and the ability of
dinitrogen tetroxide, dinitrogen trioxide, and nitrous acid to
dissociate into nitrosyl ions, the typical reactions of a solvent
system allow systematization of the preparation of nitrosyl salts,
complex nitrosyl compounds, and certain anhydrous inorganic
nitrates. Oxidation-reduction reactions serve as a means to detect
the presence of the nitrosyl group in its ionic form as well as
its covalent form.
-55-
Bibllography
1. Addison, C. 0. and Conduit, C. P., J. Cher.. Soc. 1952, 1390.
2. Addison, C. C. and Conduit, C. P., J. Chen. Soc. 19_52, 1399.
3. Addison, C. C, Conduit, C. P., and Thompson, R., J. Chem.
Soc. 1951 , 1298.
4. Addison, C. C. end Le-ris, J., J. Chem. Soc. 1951, 2843.
5. Addison, C. C. and Thompson, R., J. Chem. Soc. 1949, S211 .
6. Angus, T.T. R., Jones, R. Ti., and Phillies, G-. 0., Nature 164,
433 (1949).
7. Angus. W. R. and Leckie, A.H.. Trans. Faraday Soc. 3JL, 958
(1935) > Nature 134., 572 (1934); Proc. Roy. Soc. (London) A149T
387 (1935).
8. Asmussen, R.Tf., Z. anorg. u. allgem. Chem. 243, 127 (1939).
9. Balz, G-., and llailander, E., Z. anorg. u. allgem. Chem. 217,
161 (1934).
10. Cuttica V., Tarchi, A., and Alinari, P., Gazz. chim . ital. 53,
189 (1923),
11. Gall, H. and Ilengdehl, H., Chem. Ber. 60E, 86 (1927).
12. Hantzsch, A., Z. physi!:. Chem. 65, 41 (lOOC).
13. Lange, W., Chem. Per. 60S, 962 U.927).
14. Lange, h'., Chem. Per. 6JLB, 799 ( 1928).
15. Iteyer, J., and "fagner, W., J. Am. Chem. Soc. 4£, 1032 (1922).
16. Iloeller, G.T., J. Chem. Educ, 23, 441 (1946).
17. Partington, J.R. and ".Ihynes, A.L. , J. Chem. Soc. 1948, 1952.
18. Partington, J.R. and T.7fcynes, A.L., J. Chem. Soc. 1949 , 3135.
19. Rheinboldt, H. and 7as serf utor, R., Chem. Per. 60B, 732 (1927).
20. Ruff, 0., Z. anorg. u. allgem, Chem. 58, 325 (1908).
21. Seel, F., Z. anorg. u. allgem. Chem. 5H, 325 (1908).
22. Seel, F. and Bauer, H., z/ Naturforsch. 2Jd, 397 (1947).
23. Seel, F., Bocz , A .P., and Ilogradi , J., Z. anorg. u. allgem.
Chem. £64, 298 (1951 ).
24. Seel, F. and Crossl, P., Z. anorg. ul allgem.. Chem. 263, 253
(1950).
25. Seel, F. and Nogradi , J., Z. anorg. u. allgem. Chem. 264, 311
(1951).
26. Seel, F. and '.falassin, P.H., Z. anorg. u. allgem. Chem. 261.
85 (1950).
27. Sudborough, J.J., J. Chem. Soc, 5_9, 655, 662 (1891).
28. Wilke-DTfurt, E. and Balz, C-., Z. anorg. u. allgem. Chem. 159 f
197 (1926).
29. Woolf, A. A,, J. Chem. Soc. 1950, 1055.
30. Yost, D.H. and Russell, H., "Systematic Inorganic Chemistry,"
Prentice-Hall, Inc., ?T.Y., 1943, p. 49.
/ <-\ \
' -r. ■ *
\V •
V > /x s .W < <fc^ s *
-56-
THE OCODBRSITCS OF MAXIMUM OXIDATION STATES
A1IONG- THE FLUOROCOltPLEXES OF THE FIRST TRANSITION SERIES
D. H. Busch November 25, 1952
Introduction
In his book on unfamiliar oxidation strtes, ICleinberg (l)
states that the potentialities of fluorine as an oxidizing agent
for the study of higher oxidation states have scarcely been tested.
The implications of this statement are emphasized by the studies
under consideration here* The higher of the normal oxidation states
of an element commonly occur in combination with fluorine. (2)
This is illustrated in Table 1 which lists the binary fluorides
of the elements of the first transition series. (3)
ScF
Cr?2
hnra
FeF2
CoF3
TiF3
VF3
Cr?3
MnFa
FeF3
CoF3
TiF4
VF4
CrP4
VFB
Table 1
ITIPa CuF2 ZnF;
Binary Fluorides of the First Transition Series.
The tetra and pentafluorides of vanadium and the tetrafluoride of
chromium are quite significant in pointing out this relationship.
The highest valencies so far obtained for many of these
elements are associated with the formation of oxy-anions. In these
compounds the dual effects of the presence of oxide ions, which
are difficult to oxidize, and the formation of complex ions are
both operative. It is then to be expected that fluorocomplexes
can be prepared in -which high valence states are associated with
the central metal atoms. The highest valence states known in
fluorocomplexes prior to the work to be discussed here are summar-
ized in Fig. 1. The maximum valence states which have been observed
for these elements in oxj'complexes are shown in Fig. 2.
-57-
7
6
5
4
© ®
3
® ©© ® ®
2
@ @ ®
1
-
Sc Ti V Cr In Fe Co I'i Cu Zn
Fig. 1
The maximum valencies found in fluoro-
complexes of the transition elements as of
1948. The circled number gives the refer-
ence as listed in the bibliography.
7
6
5
4
3
2
1
o
o
Ti V Gr Hn Fe Co Hi Cu Zn
PI i
9
The maximum valencies found in oxy-
comple::es of the transition elements.
-58-
Preparation of Fluoroconplexe s Containing Central Atoms of High
Valence
Xleram and Hugs (2,14) prepared fluorocomplexes of the ele-
ments vanadium through copper by fluorination of mixtures of
potassium chloride and the chloride of the appropriate metal. In
each case the molar ratio of potassium chloride to transition
metal chloride was varied until a homogenous product was obtained.
Each product was then characterized by analysis, weight change
during fluorination, and by x-ray and chemical means.
lied potassium hexafluoronichelate (IV) (IC21T1F6) prepared in
this manner was found to be isomorphous with potassium hexaf luoro-
silicate (2). It is hydrolyzed by water, and reduced by hydrogen
according to the equation
KsNiFe + Hs -> K2NlF4 + 2HF
TJith copper (2) a pale green complex of the composition
K3CuF6 was obtained. This compound is also decomposed by water.
Hydrogen reduction produces copper (i) fluoride.
In the case of cobalt, (2) bright blue potassium heptafluoro-
cobaltate (IV) (l£3CSoF7) was prepared. This substance is isomor-
phous with K3ZrF*. It is fairly stable toward reduction by
hydrogen; however, at 450° it is slowly converted to potassium
hexaf luorocobaltate (ill ) •
IC3CoF7 + -|H3 -> K3CoF6 + EF
Potassium heptaf luorocobaltate (IV) reacts with water releasing
only a small part of its fluorine.
Fluorination of iron salts produced only a trivalent hexa-
fluorof errate (ill ) complex(2) which was found to be identical with
the product obtained by a wet method.
The only product obtainable by the method of Klemm and Huss
with manganese (14) was potassium hexaf luoromanganate (IV) (K2IlnFQ)*
This product was identical with that obtained by Teinland and
Laurenstein (5).
Similar experiments with chromium (14) were complicated by
the formation and volatilization of chromium (v) fluoride. With
the molar ratio of 2 ICC1 to 1 CrCl3, the amount of chromium lost
by volatilization was a minimum, and yellow potassium hexaf luoro-
chronate(lV) (K2CrF6) was found in the solid residue. Potassium
hexaf luorochromate (IV) is isomorphous with potassium hexafluoro-
manganate (IV)« '.Then heated in the presence of hydrogen, the
chromium disproportionate r -ith the result th«t volatile chromium
Kv) fluoride sublimes away, leaving behind green potassium
nexaf luorochromate (ill ) .
2K2GrFe-r*CrFfi +::,>?« +
nr
ihe same reaction occurs in an atmosphere of nitrogen although it
does not proceed to completion at 300o.
-59-
Difficulties similar to those associated with the chromium
compound were encountered with vanadium (14). At reaction temp-
err tures in excess of 100° vanadium (V) fluoride sublimes away.
The best results were obtained by subjecting the green trivalent
complex Iv2VF5 to fluorination. The nearly colorless product of
the fluorination reaction had the composition ICsVFe# The potassium
pentaf luorovanadate (ill ) was prepared (6) by the addition of
potassium fluoride to r solution of vanadium (ill ) fluoride in
anhydrous hydrogen fluoride.
(15) succeeded in preparing the potassium,
f the hexafluorovanadate (v) ion by the
Erne leu s and Gutmann
barium and silver salts of the hexaf luorovanadate (V) ion bj
reaction of vanadium (ill ) chloride and potassium chloride in
anhydrous bromine trifluoride. Potassium hexafluorovanadate (V)
(KVF6) was also prepared by the reaction of potassium fluoride
with vanadium (V) fluoride in a sealed tube. These salts fume in
air liberating hydrogen fluoride and vanadium (V) oxide. KVF6
is decomposed in vacuo at 330°,
Two tetravalent fluorocomplexes of manganese were prepared
by Sharpe and Tfoolf . (lG) The reaction of potassium permanganate
with the solvent bromine trifluoride produced the pink salt,
potassium pentaf luoromanganate (IV) • The x-ray pattern of this
product showed that no potassium fluoride, manganese (ill ) fluoride,
or potassium hexaf luoromanganate (IV) was present. K!:nF5 reacts
with water liberating manganese (IV) oxide, hydrogen fluoride, and
potassium fluoride. Yellow potassium hexafluoromanganate (IV) was
prepared by the reaction of potassium permanganate with bromine
trifluoride in the presence of potassium chloride. This product
was identical with IC3IInFe prepared by a wet method (17).
The reaction of chromate ion with bromine
produced only mixed oxyfluoro complexes.
;rilfuoride (16)
Interpretation of the gx")pr Omental Results
The most obvious conclusion that can be drawn from the pre-
ceding review is that the higher valence states are not so well
stabilized through the formation of fluorocomplexes as might be
expected in view of the hnown oxycomplexes . The highest oxidation
states obtained in fluorocomplexes are summarized in Fig. 3.
JJI
L> •
7
6
The maximum
valences known5
for fluorocom-
plexes of the 4
elements of
the first 3
transition
series. 2
Ti V Cr I In Fe Co Mi Cu Zn
■ , t
-60-
IClemm (10) calculated the heats of reaction for the reactions
shown below.
3KF + VF3 -» K3VF6
2"F + VF4 -» K2VF6
KF + VFS — > XVF6
Kis calculations indicate that, on the basis of electrostatic
relationships, the tetravalent complex should be the most stable
of the three while ICVF6 should be relatively easy to decompose
into vanadium (V) fluoride and potassium fluoride. This is in
agreement with the experimental findings.
The failure of investigators to prepare fluoro complexes in
which chromium, manganese,, iron, cobalt, and nichel exhibit their
maximum known valences ( though admittedly an inferior type of
evidence for their inability to exist) should lead to a reexam-
ination of the reasons given for expecting such compounds to form.
The stabilizing effect on valence state which accompanies complex
formation is not the sane with fluorine as with many other ligands.
This is associated with the fact that the bonds formed between
fluorine and the transition elements are primarily of the ionic
type (10), so that the stable electronic configurations attainable
by the central atom do not involve filling of its electronic
orbitals with electrons contributed by coordinating ligands. The
stable configurations are, to the contrary, the same as those
associated with gaseous atoms, i.e., completely filled, half-filled
and empty electronic shells. (2, 10 ) Recognition of this fact
facilitates the explanation of the stability of the FeF6~~3 ion and
the stable existence of the CoF7"3 ion.
The absence of experimental proof for the existence of stable
fluorocomplexes in which chromium, manganese, and nichel attain
such stable configurations may be interpreted to mean that there
is an upper limit to the number of electrons which may be removed
from these atoms and placed on associated fluorine atoms (18) •
Two factors favor the lihlihood of attaining higher valencies
with oxy complexes than with fluorocomplexes. According to Klemm
(18), the negative two charge on the oxide ion favors a higher
positive charge on the nucleus of the central atom on the basis of
electrostatic and space considerations. The electronegativity
difference between the average value for the elements of the first
transition series and that for oxygen indicates that the corres-
ponding bonds should be about 55$ ionic (l9)» On this basis it is
reasonable to expect the formation of bonds in which electron
sharing plays an important part so that an upper limit on the
oxidation state of the transition element need not be expected on
the basis of electrostatic considerations.
-61-
Bibliop-rar?hy
1. ICleinberg: "Unfamiliar Oxidation States and their Stabil-
ization," p. 12, University of Kansas Press, Lawrence,
Kansas, . 1950.
2. Klemm and Huss: Z. anorg. allg. Chen., 25£, 221 (1949).
3. Haszeldlne and Sharper "Fluorine and its Compounds," p. 48,
I'ethuen and Co., Ltd., London, 1951.
4. Sldgwick: "The Ohemlcal Elements and their Compounds," Vol. I,
p. 441, The Oxford University Press, London, 1950.
5. Kovalewski: Z. anorg. allg. Chem., 25, 139 (1900).
6. Petersen: Ber., 21, 3257 (1888).
7. Fabris: Gaz.. 20, 582 (1890); Helmolt: Z. anorg. allg. Chem.,
3, 125 (1898).
8. tfeinland and Lauerstein: Z. anorg. allg. Chem., £0, 40 (1899).
9. Peters: Z. Phys . Chen., 26, 195, 219 (1898).
10. Pauling: "Nature of the Chemical Bond," p. 116, Cornell
University Press, Ithaca, Hew York, 1948.
11. Jones and Tasher: J. Chen. Soc . , £5, 1904 (1909).
12. Kaas: Chen. Ztg., 32, 8 (19CS).
13. Si dgwi cli : "The Chemical Elements and their Compounds," Vol. I,
p. 285, The Oxford University Prers, London, 1950.
14. Klemm and Huss: Z. anorg. allg. Chen., 262, 25 (1950).
15. Smeleus and C-utnann: J. Chem. Soc, 1949, 2979.
16. Sharpe and Uoolf: J. Chen. Soc, 1951 T 798.
17. Bellucci: Atti R. Accad. Lincei, 22 II, 579 (1913).
18. Klemm: Naturwissenschaften, 37, 175 (1950).
19. Pauling: "Mature of the Chemical Bond," p. 70, Cornell
University Press, Ithaca, Hew Yorh, 1948.
-62-
POLY- and LETAPHOSPHATLS
V. D. Aftandilian December 2, 1952
The dehydration of NaH;jP04 yields sodium metaphosphate
according to the following scheme:
NaHaPO* -A — * Na3H3P307 -^ — * NaP03
Many investigators have prepared metaphosphates by heating
mixtures of Na2HP04 or Iia4?a07 and an ammonium salt (l;. Prac-
tically all the known metanhosphates have been prepared by a
slight modification of this method (4,5,6), i.e. by changing the
ratio of the ammonium salt aid the final temperature of the
dehydration process. However, this method has not been investi-
gated systematically nor has the exact function of the ammonium
salt been elucidated.
Kantzer (2) carried out the reactions of Na2HP04 with NH4N03
using a Chevenard Thermobalance (3) to investigate the intermediate
products of dehydration. Typical graphs shoeing the relationship
between mole ratios of NH4N03/Na3HP04, rate of heating end temp-
erature are presented in Figures I, II, III#
In a series of experiments where equal weights of ammonium
nitrate and disodium hydrogen phosphate were used Kantzer (2)
found the final products of dehydration to be either Na4P307 or
NasPsOjo, depending on the final temperature and ;n the rate of
heating of the mixture. The weight of the excess NH4Jf03, decom-
posed in these experiments, was found to be proportional to the
weight of the initial Na&HP04. And for the same weight of di-
sodiumhydrogen phosphate the loss of the excess nitrate, in the
case of the formation of Ka6P3Qio and- Na4p307 was found to
correspond to the ratio of 2/3. (See table IJ
Table I
A B
Anhydrous Loss of xfs Final Product Ratio
Na2HP04 NEi^Oa A/B
gm ga
0.1517 0.0579 Na5P3010 2.6
0.0841 0.0311 n 2.6
0.0795 0.0 20 Na4P207 3.9
0.1208 0,030 n 3.9
In order to explain this phenomenon Kantzer postulates that the
first product of the reaction must be a metaphoSDhate. Depending
on the final temperature either Na4P307 or NasP30iO is formed when
<a
u
P.
6
Eh
300 °C
250
200
*- - 4.47
constant
100
^0°/hr
rate of heating
Fig, I Temperature vs. Rate of Heatim
(D
P,
300 °C
250
/ /'. fff/tnic
200
~~§ ratio^of moles
NH4 NO 3/moles
Fig. II Temperature vs. Ratio of NtUNOg/NagKPO* * *
- 265°
F1* ttt n 4.4 NH4W0, 5^V?r lOOVhr rTte^of heating
Fig. Ill Ration {g*gfr vs. Rate of Heating at const, temp. *
T# i*
« » «
>*•'
-64-
the metaphosphate reacts with sodium nitrate according to the
following equations:
(a) 2NaP03 + 2NaN03 ) Na4P307 + N20s
(b) 3NaP03 + 2NaN03 ) NasPsOio + N30e
It is interesting to note that these reactions take olace either
below or above the region where the trimetaphosphate can exist.
(See Fige. I, II and III), Therefore the metaphosphate postulated
by Kantzer cannot be the trimer. The metaphosphate, initially
formed, is very unstable and reacts with sodium nitrate. If the
temperature is kept constant N20s is said to be evolved(?) and
Na4P207 is the only product of dehydration. But, if the temper-
ature is raided to the domain of the existence of Na5P3010 the
following reaction could take olace:
6NaP03 + 6NaN03 v ^ 3Na4P307 + 3N205 = ■■ ^ 2(NaP03)3
+ 6NaK03
Trimetaphosphate could react with sodium nitrate to give Na5P30lo.
3v«aN0>|fc-+ 8NaP0a ^ NaBFaOio + N306
Kantzer (2) isolated the dehydration product of a mixture of
equal weights of Na3HP04 and wH4N03, heated to exactly 236*C. He
obtained a yellowish liauid which gave an opaque material when
chilled auickly. This was dissolved in a very small auantity of
cold water. Oily droplets were formed immediately upon the
addition of ethanol. Acetone, dioxane, and carbon tetrachloride
^ave the same result. The resulting compound gave no precipitate
with AgN03 (characteristic of the ortho phosphate anion). It was
found to be very hygroscopic and its density diminished rapidly
when exposed to moisture giving a solution which had the character-
istics of pyrophosphate. (vVhite precipitate with ethanol). A
poorly formed crystalline compound was obtained when it was dried
in dry atmosphere. X— ray powder diagram and the infra-red
absorption spectra showed that this compound is neither NaBP30io
nor Na4P207» The existence of dimeta-nhosphate, among the -products
of dehydration of disodium dihydrogen pyroohsoohate was presumably
verified by Kantzer (2) by cryoscopic measurements. The reactions
of the systems Nai'i03(rJaP03 )3 and lv:aN03-(NaP03)3 were investigated.
For the first of these, Na4P307 was found at 260°C, (NaP03)3 at
270 *C, and NaBP3010 at 450°C. The second mixture did not react
below 300°C, but between 305 - 700°C. Ma5P30lO was obtained.
No Na4P207 was detected in the second mixture. Taking into con-
sideration the above experimental facts Kantzer (2) summarizes
the reactions of Na2HP04 and NH4M03 to take place as follows:
1. Na2KP04*acn ) Na2HP04 + aa.
2. Na3KP04 + NH4NO3 ■) NaHaP04 + Na.N03 + NH3
3. 22JaK2P04 -^ Na2H2?207 + H20
4. decomposition and volatilization of excess NH4N03
5. NaaK^aO, ^ 2NaP03 + H20
6. (a) 2NaP0s + 2NaN03 7 Wa4P207 + N20B
. (b) 3NaP03 + 2I\IaN03 ^ NaaPaOio + Na0B
I .
. . - - -■■■„',■-■
• -• < ■ .;
• ,"
-65-
Steps 4 and 5 are reversed if the rate of heating is increased.
Boulle (8) investigated the dehydration of NaK3?04 and found
that (NaPC3)a and an insoluble compound, called Kaddrell's Salt,
are the dehydration products. Ke also reported that as the temp-
erature of dehydration is increased the content of kaddrell* s Salt
is also increased. Investigations carried out on Kaddrell's Salt
showed that it is composed of two crystalline forms, called E and D
by Boulle (8). He summarized his investigations by the following
scheme;
m v r>n v m u t> n \ / Soluble \ trim eta
NaHgPO* } Na3h3P307 \ i tZ~+ r> zio^op * <
' ~ Unsoluble fmeta I t8d0C) trime^
\metB. D 550° T trimeta
Partridge, Everett, Kicks and Smith (9) also reported two insoluble
forms II and III in Maddrell's Salt and summarized their work as;
HrfO.XlI gO-WCj, NaPOaII 475-S00°C; HaPOg j 686^ fuglon
Kpntzer (2) studied the infra-red absorption spectra of the various
forms of meta-, pyro-, and triphosphate, and found that certain
bands are common to all the<3e compounds, indicating a close
similarity between the structures of these compounds*
Kantzer suggests two possible structures for dimetaphosphate.
^v _ 0 - P<- o x 1
N,
(I) (II)
It is obvious that the second structure could give rise to cis-
trans isomerism. The close analogy between the infra-red absorp-
tion spectra and the A^-ray diffraction data of dimetaphosohate
and KaaHaP307 suggests (I) as a more favorable structure for
(NaPOs)a, This structure also explains the extreme reactivity
of this compound since one of the phosphorus atoms lacks a pair of
electrons and acts as a Lewis acid. It adds water and many donor
groups, Kantzer (2) found that dimetaphosphate decolorizes Fe(SCN)r
but he was not able to isolate the resulting compound.
The author also suggests structure (II ) as an explanation
for the two isomeric forms, B and D, of Kaddrell's Salt. However,
he gives no experimental evidence to support this suggestion.
No complexes have been reported for trimetaphosphate.
The tetrametaphosr>hate can be prepared by allowing hexagonal
phosphorus(V) oxide to react with sodium carbonate 10-hydrate (10).
The tetramer is capable of forming complexes with heavy metals
(11 K Kantzer (2) prepared Fe [Fe(P03 )4C13] and Na3£Fe(P03 )«Cl£f
by adding a solution of xeCla to a solution of tetrametaohosphate.
-'. -; i $
... i -.
•i i
• V
i i
r ....
■*V-Jl
..• . i ; » -J
A •:;■•;■-■- j
\ , , . t
: , c h'^s-H
. / ' '• t. ; . ? tr •■-. ,•■:
-66-
The structure suggested for this complex is shown in figure I,
0
0->p
CI
!
!
-t-0-
1
i
' n
r —
/
/
0
1
---.0
V
/ 1
1 .
/
o — ; —
i
Cl
I
-P
V
0
Figure I
Graham's salt, sometimes called hexametaphosphate, is a
linear molecule made up of F04~ tetrahedo- linked together by
common oxygen atoms. This salt forms, complexes. Kantzer (2)
prepared the lead complex and by gravimetric analyses established
its formula as iJa2fPb(P03)4j «1-|- Ka0. Ke postulated that the lead
(II) ion is capable of breaking the linear chain of this poly-
phosphate into eigfc£-m ember ed cyclic structures. No structure is
suggested by Kantzer for this compound.
Two complexes were prepared from triphosphate and FeCl3
according to the following reactions;
NaePdOio 4- FeCl3 ^ Na3CFe(P30lo )Cl} + 2NaCl
NasPaOio + 2FeCl3 — -}. Fe !Fe(P3010 )Clj + 5NaCl
When a solution of Fe(i\i03)3 is added to a solution of triphosphate
? yellowish— white precipitate is formed. Kantzer identified this
compound as FeB(P30iO)3.
Bibliography
1. Knorre: Z. anorg. Chem., 24, 369 (1900).
2. Kantzer: Ann. Chim., 12, 819 (1950 ).
3. A.D.A.k.E.L. : Chevenard Thermobalance; 4, Passage Louis
Philippe, Parix(Xl), France.
4. Travers and Chu: C.R. , 196, 2169 (1934).
5. Bonneman: Ann. Chim., l|7~395 (1941).
6. Rechid: Thesis, Paris, France, 1933.
7. Rechid: ibid.
8. Boulle: Thesis, Paris, France 1941.
9. Lecomte, Boulle and Domine-Berges: Bull. Soc, Chim., £3, 764
(1948).
10. Topley: Quart. Revs., j3» 345 (1949).
11. Kantzer: C*B., 2£0» 661 (1945).
J2.
1
.
■ f
»■
r
•i * ' ' '
■ . ,.' r*
-67-
PRODUCTION OF SILICON CHLORIDES BY ELECTRICAL
DISCHARGE AND HiaH TEMPERATURI TECHNIQUES
W. E. Cooley December 9, 1952
X. Nomenclature
Compounds of the general formula SinClgn+2 are known as
silanes," and prefixes denote the number of silicon atoms present.
(1,3)
SiK4 silane
3i2H6 disilane
3i3H8 trisilane, etc.
Cyclic silicon compounds of the general formula (SiH2)n ace
known as cyclosilanes. The number of silicon atoms is denoted as
before, always with cyclo- as a first prefix, (l)
SiH3
/ \ Cyclotripilane
SiH2 SiK2
SiH2 SiH2
Cyclotetrasilane
SiHa- SiK
2 uj-i*2
Halogen derivatives of the silanes are named by prefixing to
the name of the parent compound the nature and number of halogen
atom** present. (1,6)
SiH2Cl2 dichloro silane
5i2Cl6 hexachlorodilsane
3i3Cle octachlorotri«ilane
oiClg— 3iCl2
octa.chlorocy clotetrapilane
SiCl2 — Si CI
II. Early '.work
The first production of completely chlorinated silane* by
electrical discharge was accomplished by Besson and Fournier in
1908. (2) These investigators subjected a mixture of hydrogen
and trichlorosilene va^or to a discharge and obtained an oily
liauid oroduct. From the liouid were distilled SiCl4j Si2Cl6,
and 3i3Cle. Si4Cli0 and 3i5Cl12 were uncertainly reported as
viscous liauid^ and Si6Gli4 aoreared as a white solid, subliming
-68-
at 200 °C. in a vacuum. After continued heating in vacuum, a
reddish-yellow solid remained. It was thought to he a mixture of
higher chlorides, hut no further identification was attempted.
When tetrachlorosilane was used instead of trichlorosilane, a.
solid residue again remained after removal of the lower chloro-
sllanes.
III. More Recent Studies
silicon
ur
In 1937 Schwarz and his associates at the University of
Kc^nigsberg published the first of a series of reports on silj.
chlorides of high molecular weights. Schwarz and Keckbach (14)
used a so-called "hot-cold" tube for their preparations. A mixti
of tetrachlorosilane vapor and hydrogen gas was admitted into a
tube previously filled with an ar^on atmosphere and fitted with an
electrically heated ouartz rod. The rod was heated to 1000° -
1100°C. as the gas mixture Ilowed over it. Condensation and
collection of products was accomplished in a flask attached to the
hot portion of the. tube, the flask being cooled by a liauid air
bath. A product mixture was obtained whose more volatile com-
ponents consisted of SiCl4, SiHCl3, Si2Cl6, and HC1. Fractional
distillation under high vacuum yielded a very viscous oil boiling
at 215°-220°C. The results of analyses and cryoscopic molecular
weight determinations- led to calculation of the formula Sll0Cl23.
The compound is hydrclyzed vigorously by water, as may be expected
for all members of the SinClon+g series, according to Schumb. (8)
Schwarz and Thiel (16), employing the "hot-cold" tube, sub-
stituted a hydrogen atmosphere for the argon atmosphere formerly
used. The compound 3ii0Cl20h2 was separated as a viscous oil.
This compound and related chlorinated silanes undergo alkaline
hydrolysis in accordance with the following general eauations:
(1) SinCl2n+2 + (4n+2) KOK > nK23iQ3 + (2n+2)KCl + (n-l)Ha
+ (n+2)H20
(2) 3inCl2nE2 + 4n KOH } nKaSIOs + 2n KOl + (n+l)H2 + nH20.
Schwarz (10) used the compound Si10Cl2oH2 to prepare an
amorphous polymerized silicon chloride of empirical formula
(SlCl)x. A flask, containing benzene solution of Si10Cl20H2 in a
carbon dioxide atmosphere, was attached to a vacuum pump. After
all the benzene had been drawn off, the contents were heated to
30G°C. The volatile decomposition products, collected in receiver
cooled by liauid air, consisted of SiCi4, SiHCl3, Si2Cl6, Si3Cl8,
SI3CI7H, Si4Cl10, and in some case^ hydrogen. A yellow residue
remained behind. Hydrolysis of the yellow substance with dilute
KOH permitted measurement of the evolved hydrogen. Chlorine was
determined titrimetrically, and silicon was determined gravimetric-
ally as Si02. Several analyses gave good agreement with a Si:Cl:H
ratio of 1:1:3. Such a ratio is in harmony with the general
equations for hydrolysis, given above, in that one H2 molecule is
-69-
evolved for each Si-Si bond. The molecular weight of (SiCl)x
could not be determined, because of its insolubility in organic
solvent?.
Schwarz and Gregor (11 ) pointed out that the insoluble pro-
duct (SiCl)x, could not be a mixture of silicon and members of the
series SlnClgn+2, since benzene or ether would dissolve out the
homologs of tetrachlorosilane. No loss of weight was experienced
by the yellow material on treatment with solvents* no residue
remained behind on evaporation of the solvent. Hydrolysis of a
mixture of the above type would leave elemental silicon; actually
a pure white, homogeneous hydrolysis nroduct was formed with care-
ful hydrolysis. Furthermore, the yellow material remained un-
changed at 400 °C, and one atmosphere pressure. A conjugated
structure,
CI CI CI CI CI
-Si = Si - Si = Si - Si=,
was tentatively proposed by the authors; Rochow (7) and Sidgwlck
(17) have, however, objected to a structure consisting of chains
of doubly-bonded silicon atoms.
IV* Present Studies
Work in this field is now being carried out by Hertwig and
Wiberg (4,5) and Schwarz and Kd'ster. (l£,13) Hertwig and Wibergt
working at the University of Munich, have used the glow discharge
apparatus of Schwarz and Pietsch. (15) An evacuated tube system
was fitted with two aluminum electrodes arranged to produce a glow
discharge of 10,000 volts. Hydrogen gas and SiCl4 (or S1HC13)
were admitted at low pressure. It was found that under prober
conditions 90-100^ of the chloride could be converted to stable
materials ranging in composition from SICIq^ to SiClp g. Each
of the several compositions wao found to be renroduciBie, and
Quantitative recovery of all products was shown by agreement of
stoichiometric calculations with analyses of the non-volatile
products. The materials with lowest chlorine content consisted of
solid powder «s" those of higher chlorine content were resinous and
oily, and increasingly soluble in benzene and ether. The individ-
ual reproducible compositions were apparently often mixtures in
themselves, for chlorine— richer components could be dissolved by
benzene or ether treatment of chlorine-poorer substances. The
following data are based upon the results of cryoscopic deter-
minations.
Empirical Formula Molecular Weight Proposed Formula
SiC1l 10 2555 3i38Cl42
S1C:L1*27 1610 Si2SCl28
SiCll*47 1526 SinClae
SiCl^eo 848 Si10Cli6
S1C12#00 891 Si?Cl18
- -■ •
•
:
f -i r
-70-
Schwarz and K6"ster have verified the findings of Hertwig and
Wiberg by isolating similar materials through the thermal decomp-
osition of Si10Cl20H3 in a stream of argon. By varying the
decomposition temperature from 260°C. to 320^0., it was possible
to produce chlorides of composition SiCl-j_ qq to SiCl-^ -^q.
Hertwig and Wiberg propose for the resulting chlorides a
structure composed of sections from the diamond lattice of silicon,
with the chlorine atoms occupying the positions necessary to
saturate each silicon atom.
Si<jCl18| m.w. 891 Si10Cl16, m.w. 848
Hydrolysis with KOH is reported to yield a auantity of hydrogen
in agreement with the number of Si-Si bonds in these pronosed
structures.
Schwarz and KdJster argue that the amorphous nature of the
products precludes any ordered structure such as a diamond lattice,
A blcyclic structure is proposed for Si10Cl1Q, for example:
Cl-Si^GLOL^Si'Cl G&S1-C1
ci-sta ?l /S,i ci 0.-51-31
^i^Cl^Si-"'
ci di
Further condensations, to give products of even higher molecular
weight, occur in a random three-dimensional manner, resulting in
irregularly interlaced "bundles"* As the temperature is increased
above 290°C, Sii0Cl18 decomposes to substances of higher Si'.Cl
ratios. Above 800°C. a residue of silicon is obtained which is
shown by X-ray studies to be amorphous.
V. References
1. American Chemical Society Official Reports, Chem. Eng. News,
30, 4513 (1952).
2. Besson and jfournier: Comnt. rend., 148, 839 (1908); 149 34
(1909).
3. Crane: Chem. Eng. News, 24, 1233 (1946).
4. Hertwig: Z. Naturforsch. , 6b, 337 (1951 ).
5. Hertwig and Wiberg: Z. Naturforsch, 6b, 336 (1951 ).
6. Patterson: J, Am. Chem. Soc, 55, 3905 (1933).
-71-
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
Rochow
2nd ed
Schumb
Schumb
Schwar
Schwar
Schwar
Schwar
Schwar
Schwar
Schwar
Sldgwi
p. 551
MAn Introduction to the Chemistry of the Silicones, n
, V* 77, John Wiley and Sons, Inc., New York, 1951.
Chem. Revs., 31, 587 (1942).
and Gamble: Inorg. $yn. , I, 42 (1939).
z
z
z
Z
z
z
z
ntr •
Angew. ohem.,
51, 328 (1938).
and Gregor: Z. anorg. allgem. Chem., 241, 395 (1939).
and KdJster: Z. Naturforsch. , 7b, 57 (1953).
and Kd'ster; Z. anorg. allgem. ""Chem. , 270, 2 (1952).
and Meckbach: Z. anorg. allgem. Chem., 232. 241 (1937).
and Pietsch: Z. enorg. allgem. Chem., 232, 249 (1937).
and Thiel: Z. anorg. allgem. Chem., 235, 247 (19,38).
flm-i
he Chemical Llements and their Compounds
The Oxford University Press, London, 1950.
Vol. I
-72-
FLUORINE CONTAINING GXYHALIDLS OF SULFUR
E. H. Grahn December 16, 1952
Introduction
In his work on fluorine containing oxyhalldes of sulfur,
Jonas (5) has described the preparation and properties of thionyl
chlorofluoride, sulfuryl bromofluoride, and ^thionyl tetrafluoride".
Several compounds containing sulfur directly linked to fluorine
and oxygen are known (7).
Melting Boiling
Name Formula Point Point
Thionyl fluoride
Sulfuryl fluoride
Sulfuryl chlorofluoride
Thionyl chlorofluoride
Sulfuryl bromofluoride
"Thionyl tetrafluoride"
SOF3
SO gif g
S02C1F
SO GIF
SO a Br F
SO F4
-110 °c.
-43.8°C
-120
-52
-124.7
7.1
-139.5
12.3
- 86
"40
-107
-48.5
Thionyl fluoride
Thionyl fluoride, S0Fs, was first prepared by Meslsns (6) in
1896 by reacting thionyl chloride with zinc fluoride. A better
method, however, is that of Booth and Mericola (3), who obtained
the compound by reacting thionyl chloride with antimony trlfluorlde
using antimony oentachlorice a" a catalyst. Thionyl fluoride is a
relatively stable compound which hydrolyzes only slowly in water
and does not attack gla«s below 400°C. It is rapidly attacked by
dilute alkali solutions.
Sulfuryl fluoride
Sulfuryl fluoride, S02F3, was first described in 1901 by
Moissan (8), who prepared it by the reaction of fluorine with
sulfur dioxide in the presence of a platinum catalyst. It can be
obtained in fair yields by decomposing barium fluo sulfonate (9).
Sulfuryl fluoride is a very stable compound, not being decomposed
by hot water. It does not attack glass even at high temperatures,
and molten sodium has no effect upon it.
Sulfuryl chlorofluoride
Sulfuryl chlorofluoride, 303C1F, was carefully characterized
by Booth and Herrmann in 1936 (2). It was prepared in good yields
from sulfuryl chloride and antimony trifluoride in the presence
of antimony pentachloride as a catalyst. This compound is inter-
mediate in stability and in reactivity between the chloride and
the fluoride, It hydrolyzes slowly in water and very rapidly in
dilute sodium hydroxide solutions. It does not attack dry glass,
mercury, or the common metals at room temperature.
*»■ -■' --—.
. • 1 Or ,-i
. ■■ .
'. ... u'fV
,- '
■73-
Thlonyl chlorofluorlde
Thionyl chlorofluorlde, S0C1F, was prepared by Booth and
Mericola (3) In 1940 by fluorinating thionyl chloride with antimony
trifluoride using antimony r^entachloride as a catalyst. Actually
this is the same method used for the preparation of thionyl
fluoride. On the fractional distillation of the reaction mixture,
Booth and Mericole obtained approximately fifty percent thionyl
fluoride, twenty percent thionyl chlorofluorlde, five percent sulfur
dioxide, and twenty-five oercent unreacted thionyl chloride, Jonas
(5) prepared it by reacting thionyl chloride (0,5 mol) with iodine
pentafluoride (0,2 Mol) and then fractionally distilling the pro-
duct. The following six fractions were obtained.
Fraction
Temperature Range
Products
1
2
3
4
5
6
-70 nC. to -50 °C,
-50 to -36
-36 tO -22
-22 tO 10
10 to 18
Rest
SiF4 and HCl
KCl and S0F2
S0Fs (12 cc. )
S0F8 and S0C1F
S0C1F (10 cc.)
SO CI 2
PHysical proper
ties
of
thionyl chlorofluorlde
Vapor density = 103.2 g./22.4 1. (calc. 102.5 g./22,4 1.)
The vaioor pressure follows the equation, log p = 7,83 - 1409
T
The density may be ex^reoced by the eaua.tion, D=l. 576-0 ,00224t
Boiling point = 12.3°C.
Jonas was unable to determine a melting point for the compound.
However, Booth recorded the melting -noint as -139°C,
Jonas hypothesized that ttiionyl chlorofluorlde is actually a
mixture of two stereoisomer^, hence the abnormally low melting
point as compared with thionyl chloride and thionyl fluoride. The
following structure* for thionyl chlorofluorlde were suggested.
This is the only known inorganic thionyl comoound with two differ-
ent subotituents. However, thionyl bromochloride was thought to
have been prepared by Bespon (l) from thionyl chloride and
hydrogen bromide, but it was later proven by Hayes and Partington
(4) to be an eauimolecular mixture of thionyl chloride and thionyl
bromide.
-74-
Thionyl chloro fluoride at 0°C. is a colorless volatile liquid.
It reacts only very slightly with dry glass or quartz at room
temperature. Under the catalytic effect of moisture, it forms
silicon tetrafluoride, sulfur dioxide, and hydrogen chloride. It
is hydrolyzed very quickly by water giving as products the fluoride
ion, chloride ion, and sulfur dioxide or sulfite ion.
Sulfuryl bromo fluoride
Jonas used two different methods for the preparation of this
compound; the reaction of trichloromethanesulfuryl chloride with
bromine trifluoride and secondly, that of reacting a mixture of
bromine trifluoride, bromine, and sulfur dioxide. The fir«t method
gave a poor yield, since it appeared that the chief reaction was
the formation of sulfuryl fluoride. The second method gave prac-
tically a quantitative yield.
Physical properties of sulfuryl bromofluorlde
Vapor density = 162.0, 162.6 g./22.4 1. (calc. 16? g./22.4 1.
The vapor pressure follows the eauation, log p = 8.03-lf 10
T
The density may be expressed by the equation, D = 2.75 -
0.00298t
The density of solid S02BrF at the temperature of liquid air
is 3.16 g./cc.
Melting point = -86°+0.5cC.
Boiling point at atmospheric pressure = 40.0°C.
Sulfuryl bromofluorlde is a colorless choking gas. At room
temperature it resets with dry glas^ but not with auartz. With
water in a sealed tube there is a vigorous reaction witr. the
liberation of hydrogen bromide, hydrogen fluoride, and sulfuric
acid. With a very small ouantity of water, free bromine is
liberated. The vapor of sulfuryl bromofluorlde is fairly etpbl e.
It Just begins to change color at a temperature of 320°C. and at
340 °C. is completely decomposed into sulfur dioxide, sulfuryl
fluoride, and bromine. There is also a reaction with quartz, to
form silicon tetrafluoride.
There was an attempt to prepare sulfuryl iodofluoride in a
similar manner. Apparently the sulfur dioxide reacts with iodine
pentafluoride alone, but not with a mixture of ioc'ine pentaf luoride
and iodine, kore clarification of this reaction is needed.
"Thionyl tetrafluoride"
This compound was prepared by passing a mixture of thionyl
fluoride (18 g. ) and fluorine over a Platinum gauze at s temp-
erature of 150°C. The product was distilled through a auartz
fractionating column, and five fractions were collected.
t >:■£
-75-
Fraction Temperature Range Pro duct s
1 -90 °C. to -64 °C. SiF4 (small amount)
2 -64 to -52 SF6 U cc.)
3 -52 to -49 SF6, S03F2, and S0F4
4 -49 to -48 SO4 (10 cc.)
5 Rest SOi-a (small amount)
Physical properties of "thlonyl t etrafluorlde"
Vapor density = 122.5 g./22.4 1. (calc. 123.9 g./22.4 1.)
The vapor pressure follows the equation, log p = 7.76 - 1092
T
The density may be expressed by the equation, D = 1.653
- 0.00360t
Melting point = -107°+ 0.5rC.
Boiling point (760 mm. Hg) = -48.5°C.
At room temperature "thlonyl t etrafluorlde" is a colorless
pungent gas. It reacts violently with water and alkaline solu-
tions, liberating sulfuryl fluoride and hydrogen fluoride. The
sulfuryl fluoride slowly undergoes hydrolysis to give the
fluoride ion and the sulfate ion. It reacts slowly at room temp-
erature with mercury to give thlonyl fluoride and mercurous
fluoride.
Bibliography
1. Besson, A., Compt. rend., 122, 320 (1896).
2. Booth K. S. and Herrmann, C. V. J. Am. Chem. Soc., 58 63
(1936).
3. Booth, H. S. and Mericola, *. C., ibid., 62,. 640 (1940).
4. Hayec, K. A. and Partington, J. R. , J. Chem. Soc, 1926. 2599,
5. Jonac, K., Z. anorg. u. allgem,^ Chem., 265, 273 (1951).
6. Medians, M« Bull. aoc. Chim. France, 15, 391 (1896).
7. Miller, H. C. and frrall, F. J., Ind . Eng. Chem., 42, 2224
(1950).
8. Moissan, K. and Libeau, P., Compt. rend,, 132. 374 (1901 ).
9. Traube, W., Hoerenz, J. and Wunderlich, F. , Ber., 52B. 1272
(1919).
-76-
PREPARATIGN AND PROPERTIES OF URANYL CARBONATES
Richard A. Rowe December 16, 1952
Introduction
Discovery that minerals with the type formula X * (U02) (C03 )3 ■
n H20 exist in nature (3) (where X = Na2Ca, CaMg, Mg2, and
n = 6,12,18 respecitvely) prompted the French workers M. B ache let,
S« Cheylan, M# Douis, and J.C. Goulette (1,2) to study the forma-
tion of the alkali and alkaline earth uranyl carbonates* C,
Guillemin (4) had prepared the magnesium compound and determined
its physical constants. J. a. Kedvall (5) had previously pre-
pared and described the monouranyltri carbonates of silver,
mercury(I), mercury(II), copper (il), lead, cadmium, zinc, iron(II)
cobalt and nickel.
Uranyl Carbonates of Soc ium
The only uranyl carbonate of sodium to have been prepared
previously was Na4 (U02) (C03) 3, called sodium monuranyltricarbonate
Bachelet and coworkers subsequently prepared Nae(U02)2(C03) 5 and
Na6(U03) 3 (U02) 2(C03) 3. To simplify the nomenclature, these com-
pounds will be designated as (A),(B), and (C) respectively.
By adding sodium carbonate to a solution of a uranyl salt,
insoluble uranyl carbonate separates out,
(1) U02(N03)2 + Na2C03 -> U02C03 + 2NaN03
(2) U02C03 + 2Na2C03 -»Na4U02(C03)3
The presence of sodium nitrate in the solution interfers with the
separation of the uranyl carbonate. This disadvantage can be
overcome by adding uranium (VI) oxide instead of the nitrate, but
the formation of base somewhat limits the reaction.
(3) Na2C03 + U03 + H20->U02C03 + 2Na0H
(4) UOsC03 + nNa2C03 -»U02C03»n Na2C03
The free NaOH formed in accordance with equation (3) will react
with U03 to form sodium iriSE^ge. ^A> ^"»*^c «t'^"»^e.
(5) 2NaCH + 2U03 -> Na2U207 + H 20
The formation of (A) can therefore be represented as follows:
(6) 3Na2C03 + 3U03 -^Na4U02(C03)3 + Na2U207
The preparation of uranyl carbonates is best accomplished by
interaction of uranium(VI) oxide with sodium bicarbonate since
only uranyl carbonates remain in solution. By varying the
quantities of the reagents it is possible to obtain not only
compound (A) and also (E) and (C) »
-77-
(7) U03 + 4NpHC03 -* ,'tfa4.V02(C03)3 + C02 + H20
(8) 2U03 + 6N9HC03 -> )° Na6 (U03) 2(C03 )5 + C02 + 3H20
(9) 5U03 + 6NaHG03 -> tG)Na6 (U03) 3 (U02>3(C03> 5 * G03
+ 3H20
The formation of compounds (B) and (C) is not to be considered
as simple as is shown in equations (8) and (9), but rather as
involving the stepwise addition of uranium ( VI )cxide and carbon
dioxide to compound (a)#
(10) 3Na4U02(C03)3 + U03 + C03 -» 2Nae(U03) 3(G03) s
(11) Na6(U02)2(C03>5 + 3U03 -> Nae(U03)3 (U0a)a(C03)5
Compounds (a) , (B) and (C) have been prepared by mixing"
stoichiometric quantities of uranium (VI) oxide and sodium bi-
carbonate. At room temperature the rate of the reaction is very
slow since the uranium(VI) oxide goes into solution slowly. The
quickest method for obtaining compounds (A) and (B) is to
dissolve, in the case of (a) a 1:4 mole ratio of U03/NaHC03 in a
liter of water, and in the case of (B) a 1:3 mole ratio. After
continuous agitation for two days the undissolved uranium (VI)"
oxide is filtered off. Within two days the filtrate will de-
posit a greenish precipitate. Precipitation will be accelerated
by adding alcohol or acetone. Compound (C) can be obtained by
starting with a. solution of compound (B) and treating it with
excess uranium (VI) oxide. Carbon dioxide is bubbled through the
solution during the process. Compounds (A) and (B) dissolve
slowly in water to give stable solutions at 20° which contain
204 g./l,, and 190 g./l. respectively. Dilute sulfuric acid
and (A) react with e%rolution of C02 gas.
(12) 3H2304 + Na4U02(C03)3 -> U02S04 + 2Na2S04
+ 3C02 + H20
Both (A) and (B) decompose to form insoluble uranyl compounds
when the pH is brought into the range of 4 to 9.
Ammonia incompletely precipitates the uranium in the form of
ammonium uranate • Addition of sodium hydroxide results in the
precipitation of sodium uranate. Addition of sodium carbonate
or bicarbonate to compound (B ) brings about a conversion into
compound (A).
(13) Na6(U03)3(C03)s + 2NaHC03 -> 2Na4U02(C03 )3
+ C02 + H20
Complete precipitation of uranyl phosphate U02(HP04) #4H20 takes
place in the presence of phosphate ion. When oxygen is passed
through the solution of a uranyl carbonate uranium peroxide
(UC4»2H20) is presumably formed. Thermal decomposition begins
for each of the three compounds at the following temperatures:
U) 4200 f (B) 410°, fC) 400°. These compounds can be regarded
as the double salts - (A) 2Na2C03 •U03C03 , (B) 3Na2C03«2U02C03,
and (C) 3Na2(C03) •2U02C03»3U03. Uranyl carbonate decomposes
between 400-430° as follows:
(14) U02C03 -*C03 + U03
WS'G
... < ,
-78-
Precipitation occurs when solutions of these compounds are heated,
Dissolution of the resulting products is slow. It is therefore
assumed the solution is accompanied by hydration and that
hydra ted state is unstable and loses water when heated
dehydrating solvent is added. If a coordination
assumed for the uranium VI,
written for
authors) •
(A) Na*
this
or when a
a coordination number of 6 is
the following structures can be
compounds (A) and (B) in solution (according to the
0
(C03)=U«-0
H20
(B) Npe
Tiy,0
H20-U = (C03)2
t
CO,
Co
;co3)2=u-o
h2o
The dehydration of the hydra ted products as shown in the above
structures leads to the formula for the precipitated products* ~
The X-ray patterns of compounds (A) and (3) do not present qual-
itative differences.
Uranyl Carbonates of the Alkaline Earths2
Two methods have been employed to prepare the alkaline earth
uranyl carbonates. Uranlum(VI) oxide may be allowed to react with
the alkaline earth bicarbonate s. However, only the alkaline
earth bi carbonates of magnesium, calcium, and strontium are
soluble enough to make this method feasible. The alternative
second method involves double decomposition between soluble salts
of the alkaline earth metals and txhe corresponding sodium uranyl
carbonate. Neither of these methods has given satisfactory
results in the formation of uranyl carbonates of beryllium.
Uranyl Carbonates of Magnesium - Compound Mg(A) - Mg2U02(C03)3
was prepared by reacting the stoichiometric amounts of U03 and
MgH2(C03)2 in water kept saturated by a stream of C02 gas.
(15) U03 + 2M#a(C03)2->MgU) + C02 + H20
MgC&) is a lemon yellow solid which is strongly hydra ted, probably
18 molecules of water, and is very efflorescent in air. It
dissolves into its water of crystallization at 60°. If a U03 and
a solution of Hg(A) are agitated with C02 gas, en equilibrium
mixture of compounds Mg(B) and ilg(C) is formed. After 48 hours
of deposition, a product was collected which had a U/Mg ratio
equal to 2.4.
Compound Mg(B) was obtained by reacting the proper amounts
of Mg(A) with U02CC3#
(16) 3%2(U02)(C03)3 + U02C03 ->2Mg3(U0a)a(C03)
-79-
Uranyl Carbonates of Calcium - Two methods of preparation
can be used; The purest products are obtained by allowing the
bicarbonate to react with uranium( VI) oxide . Compound Ca(A) can
be obtained in a very pure state. It precipitates as a deca-
hydrate in the form of greenish-yellow crystals, i^hich are stable
in air. It loses 8 molecules of H 20 at 90°, and it becomes
completely anhydrous at 180°. Under ultraviolet light it emits
a greenish luminescence. Compound (B) was prepared by allowing
stoichiometric amounts of U03 and Ca(A) to react. It can also
be made by the interaction of calcium nitrate with Na(B) as
follows:
(17) 3Ca(N03)a + Na6(U02)2(C03)s -*Ca(B) + 6NaN03
If an excess of UC3 is allowed to react with Ca(B) a precipitate
is obtained which corresponds to neither compound Ca(B) nor
Ca(C). As In the case of the magnesium salts, it is likely that
an equilibrium mixture of these compounds exists in solution.
The solubility of calcium uranyl carbonates increases with an
increase in temperature. The solution becomes turbid at 60°, but
the precipitate can be dissolved by passing carbon dioxide
through the solution.
Uranyl Carbonates of Strontium - The best method for pre—
parting Sr(A) is by the interaction of an excess of strontium
chloride with Na(A).
(18) 2SrCl2 + Na4(U02) (C03)3 -* Sr2(U02) (C03) 3 + 4NaCl
The precipitate contains 9 molecules of water and exists in the
form of small crystals which exhibit a green luminescence under
ultraviolet light. Compound (B) cannot be prepared by either
of the two general methods due to the extreme insolubility of
Sr(A) which precipitates before Sr(B) can be formed.
Uranyl Carbonates of Barium - Interaction of Na(A) with
barium chloride yields a greenish yellow precipitate corresponding
to the 6 hydrate of 3a (a) . The compound Ba(B)»4H20 can be
obtained by a similar method. This compound is not fluorescent
like Ba(A) . Both compounds are insoluble and decompose in
boiling water.
Bibliography
1. M. Bachelet, E« Che y Ian, M. Douis, and J. C. Goulette:
Bull. soc. chinu, France (5) 18, 55-60 (1952).
2» H. Bachelet," iS» Cheylan, M. .uouis, and J. C. Goulette: ibid.
(5) 18, 565-9 (1952) .
3. J* M, Axelrod, P. S. Grimalde, C. Hilton, K. J. Murata:
Amer. Minerologist 36, 1-22 (1951).
4. G. Branche, J. Chevet, C. Guillemin: Bull. soc. France
Mj.nerology 74, 458 (1951).
5. J. a. Heclvall: Z, anorg. u. allgem. chem. 146r 225-29 (1925)
6. M. E# Burker: Journ. Pharm. Chem. (4) 27,347 (1878).
o
-80-
TH E NATURE OP IODINE SOLUTIONS
Ervin Col ton January 6, 1953
Introduction
It is well known that iodine solutions are either violet or
brown, the color depending upon the solvent (l) • Such solvents
as carbon disulfide, carbon tetrachloride, and saturated hydro-
carbons form violet solutions wl th iodine, while brown solutions
result with alcohols, ethers, ketones and organic nitrogen bases.
Solvents yielding violet solutions are called "non-active" sol-
vents, while those giving brown solutions are called "active"
solvents. The case of iodine in benzene is somewhat anaroolous,
for the color of the resulting solution is red-violet which
changes to violet when the solution is heated.
Precise molecular weight measurements have shown that iodine
is present in the diatomic state in all solvents investigated
thusfar (2). Hildebrand (3) has shown that the violet solutions
form a series of "regular solutions", i.e., solutions which are
formed from their components with the same change in entropy as
occurs in the formation of an ideal solution of the same concen-
tration. In these cases, the solubility of iodine may be repre-
sented by the family of curves
log N2 = 2.264 - [876.5 + k(l-N2)2]xl/T
where N2 is the mole fraction of iodine in a solution saturated
at the absolute temperature T, and k is a parameter, approximate-
ly independent of temperature for each solvent, whose value
depends upon the energy of vaporization of the solvent. Approx-
imate values of k for solutions of iodine in various solvents
are shown below:
Solvent CS2 CHC13 TiCl* CC14 C6H16 SiCl*
k 190 300 310 390 450 490
No such family of solubility curves can be usecL for iodine in
brown solutions, thus indicating the probable formation of new
molecular species.
Spec tropho tome trie investigations show an absorption maximum
in the region 515-525 mu for iodine in "non-active" solvents.
-81-
Solvent Hay . . m u Reference
Chloroform 520 (4)
Carbon tetrachloride 520 (4)
517 (5)
518 (6)
520 (7)
518 (8)
Carbon disulfide 520 (4)
518 (5)
518 (9)
n-Hexane 535 (10)
Iodine in brown solutions is much more reactive than the
free element in the violet solutions. Physical measurements
indicate that the brown solutions contain iodine chemically
bound to the solvent, in equilibrium with free iodine. £ach of
the compounds forming brown solutions with iodine contains an
element capable of acting as a donor in the formation of co-
ordinate covalent bonds. Hildebrand and Glascock (11) studied
the change in the freezing point of solutions of iodine in an
"inactive" solvent brought about by the addition of a small
amount of a third substance, itself either an "inactive" or
"active" solvent© In those cases where the lowering was less
than that calculated for the iodine on the basis of the dilute
solution law, thsy concluded that there must be at least partial
combination of iodine with the second solvent. Only with the
"active" aolvents was a large deviation observed.
I2 + n("active" solvent) ^ ("active" solvent)n'I2
Brown solutions of iodine in ether, alcohol, dioxane, among
others, react with the silver derivative of saccharin to yield
the theoretical quantity of silver iodide. Violet solutions
only react slightly.
J NNAg + I8 -* I ^NI + Agl
fl_-so/ li^L-soi
Iodine solutions of some completely fluorinated tertiary ali-
phatic amines as hep tacosaf luorotributyl amine, (C4F9 )3N, penta-
decafluorotriethylamine, (C2F5)3ii and heptadecaf luorodie thyl-
propylamine, (C2F5 )2N(C3F7 ) , are viole t (12). Violet solutions
are ordinarily observed for "non-active" solvents; in these
amines the presence of three large alkyl groups and the numerous
fluorine atoms appears to decrease tremendously the availability
of the electron pair on the nitrogen atom and thus prevents the
formation of a brown solution such as usually results with a
nitrogen base.
-82-
Hildebrand and Benesi (5) have investigated, spec tropho to-
me trically, the interaction of iodine with various substituted
benzenes. In the visible region of the absorption peaks of these
solutions showed moderate shifts toward shorter wave lengths.
Each of the aromatic hydrocarbon solutions had an intense absorp-
tion band in the ultraviolet region, shown to be characteristic
of a complex containing one I2 and one aromatic hydrocarbon
molecule.
v u.v, * visible
Solvent /\ max. , m u. ^\ max.. m a>
Benzene 297 500
Toluene 306 497
^-Xylene 319 497
^-Xylene 315 495
Mesitylene 333 490
Experimental Investigations
The exact nature of the binding between iodine and solvent
in the brown solutions is still not well understood. The case
of iodine in pyridine is an. example. Audrie th and Birr (13)
studied the change of conductivity of iodine in pyridine with
time and explained their observations by assuming the presence
of the following species in such solutions:
py + I2 ^=z Ipy + I" ^=z py + 21*"
2(pyl2) ^= IPy+ + I* pyl3"
A recent spec tropho tome trie investigation (14) seems to Indicate
that the broad band initially observed for a solution of iodine
in pyridine gradually develops towards a maximum at approximately
373 mx< in 19 days. This peak is attributed to the triiodide ion.
The situation with iodine in quinoline is somewhat easier
to study since the reaction between iodine and the amine is very
rapid (12,14). Immediately after mixing, an initial maximum is
observed at 370 m/* • A sharp increase in optical density, with
no shift in the position of the maximum, is subsequently noted.
This increase suggests that the maximum is due to a species
resulting from reaction between the iodine and the quinoline.
The maximum at 370 rn^ is attributed t~> the triiodide ion. Proof
of this is offered by the. marked similarity of the absorption
curves of iodine in quinoline to those of a known solution of
triiodide in quinoline which also shows a maximum at 370 m/* .
Even more conclusive proof that the raaxinum at 370 m/* is caused
by the triiodide ion is offered by the rapid loss, up to 50$, of
titra table iodine in solutions of iodine in quinoline. Upon
standing, the strong absorption and maximum characteristic of
the triiodide ion disappear; a new peak appears at 350-355 m/^. .
The latter maximum is attributed to an iodinated derivative that
could not be isolated. ■
•I
w\-.: ■
» T ■• * .
L ' <
If "Y
-83-
It is surprising to find that the relatively stable triiodide
ion also reacts with quinoline, although at a. much slower rate
than does iodine. The triiodide maximum in solutions of tetra-
n-butyl ammonium triiodide in quinoline gradually disappears. In
solutions of iodine in quinoline in which, after 50$ of the
iodine has been consumed and the remainder presumably exists as
triiodide ion, slow loss of titratable iodine still continues.
It seems probable that the reaction proceeds via free iodine
formation according to the equilibrium:
I3~ + quin £. ' r I~ + quin*I2
Freshly prepared quinoline solutions of a series of sub-
stituted benzoates, e.g., o~, m-, and p_-chlorobenzoates, o~ and
m-iodobenzoates, ra— and p_-nitrobenzoates, and £-bromobenzoate, of
monoquinoline iodine (I), also show an absorption maximum in the
region of 355 m^ • Hather rapid reaction involving removal of
unipositive iodine, followed by formation of triiodide ion, is
indicated by the initial decrease in optical density at 355 m/^
and the subsequent rise in optical density with appearance of
the maximum at 370 m \m • This latter peak disappears evsntually
and a third maximum is again observed at 355 m^ • This latter
peak is due to the formation of an amorphous, 'violet, polymeric
solid with an absorption peak at 355 mM . These observations
suggest a complex series of reactions 'in which the initial
material undergoes decomposition with formation of iodine and
polymeric products.
Bibliography
1. J. Kleinberg and A.W, Davidson, Chem. Rev., 4£, 601 (1948).
20 £, Beckman, Z. phvsik. Chem. f 5J3, 543 (1907).
3. J.H. Hildebrand, Solubility. 2nd edition, pp. 153-7. Rein-
hold Publishing Corporation, New York 11936) .
4. H. rfigollot, Compt. rend.. llg, 38 (1891).
5. H.a. Bene si and J.H. Hildebrand, J. Am. Chem. Soc., 71, 2703
(1949).
6. O.J. Walker, Trans. Faraday Soo . f 31, 1432 (1935).
7. &•■£. Gill an and H.a. Morton, Proc. Roy. Soc. (London), 124A,
604 (1929).
8. W.R. Brode, J. Am. Chem. Soc. 48, 1877 (1926). '
9. J. Groh, 2. anor,?. allgrem. Chem. f 162, 287 (1927).
10. F.H. Getman, J. Im. Chem. Soc., 50, 2883 (1928).
11. J.H. Hildebrand and B.L. Glascock, Ibid . t 31, 26 (1909).
12. E, Colton, i-Iaster's Thesis, University of Kansas, 1952.
13. L.F. Audrieth and JS.J, Birr, J. a^. Chem. Soc., 55, 668 (1933)
14. J. Kleinberg, £. Colton, J. Sattizahn and C.A. Vander^erf ,
ibid... In press*
■
-84-
SOkE REACTIONS OF OZONE
Barbara H. Weil January 6, 1952
Van Marum in 1785 observed that oxygen subjected to an elec-
tric discharge had a peculiar odor and tarnished mercury. It was.
not until 1840, however, that Scho'nbein recognized that the^e pro-
perties were characteristic of a new gas which he named "o^.one"
from the Greek ozo , I smell, (l)
Ozone is more soluble than oxygen in water and irore soluble
in glacial acetic acid and chloroform than in water. It can be
decomposed into oxygen by heat, by contact with powdered glass
and, catalytically, with silver, platinum or manganese dioxide,
lead dioxide and silver, cobalt and iron oxides. It is a powerful
oxidizing agent, liberating iodine from a solution of nota^ium
iodide. However, it has no action on potassium permanganate or
chromic acid. Sulfur dioxide is oxidized to the trioxide and
stannous chloride to stannic chloride in contact with ozone.
Ozone bleaches indigo solution and vegetable colors and con-
verts moist sulfur, phosphorus and arsenic into their highest
oxy-acid«. It liberates halogens from their hydracids. Moist
iodine is oxidized to iodic acid while dry iodine is converted
into a yellow powder I4O9 with liberation of oxygen. (2,3)
For a long time it has been known that ozone reacts with, the
hydroxides of the alkali metals to produce peculiar, highly
colored solid*. (4) The reaction was fir«=t investigated extens-
ively around 1900 by Ba eyer end. Villiger (5) who used the hydrox-
ides of sodium, potassium, and rubidium for their investigations •
They postulated the formation of "ozonates" of the formula K4O4.
At about the same time, bach (6) presented rather weak evidence
for"ozonic acid", H304, formed in a manner analogous to the forma-
tion of H203 from Ka303 and K3S04. The potassium ozonate of
Baeyer and Villiger was thus, he believed, the acid salt of ozonic
acid, KO4H. Evidence for the existence of ozonic acid ha* aince
been disproved, however. (9) kanchot and Kampschulte (7) found
that the stability of the alkali metal ozonates, or ozonides, ?s
they are now commonly named, increases with increasing atomic
weight of the metal, the cesium derivative being the most stable
and the lithium compound the least stable. They investigated the
ozonides of the alkaline earths and found that here, too, stabil/fy
increases with increasing atomic weight of the metal. The form- >
ation of all the ozonides is an exothermic reaction.
TraUbe, in a later work, postulated the oroduct to be an
addition compound of the metal hydroxide and oxygen, (K0H)202, (^\
Recently, the same reaction has been studied extensively by
Kazarnovskii and co-workers for the purpose of elucidating the
structure of the product, (10, 11) They obtains, cfter react:-. - >?
at low tenroeratures and extraction with ammoni 1 roduct con-
.
■V
,{
V » .
-85-
taining 88-93$ K03 with some K03, KOH and KOH-HaO. The mechanism
KOK + 03 = KOa + OK is excluded by their thermodynamic data, so
they postulate a possible mecha.nicm involving the intermediate
radical H03, namely, KOK <-+ 03 = K0H*03; KOH*03+03-*K03 + 02 + H0a .
The decomposition in water oroceeds according to: K03 + H30 =
K + OH" + 03 + OH, followed by 20H = K30 + l/2 03,
The free-radical nature of K03 was confirmed by these workers
by magnetic measurements, molar susceptibility = +1185 x 10~6;
hence, the magnetic moment = 1.57 Bohr magnetons, very close to
the theoretical 1.73, corresponding to one unpaired electron.
This confirms the formula K03 to the exclusion of K30s or K0H»03*
From results of+molar electric conductivity, the ccnvoound can
be formulated K 03~. The 03_ ion can be conceived as a resonance
hybrid of 0-0-0- and -0-0-0 or the three- electron bond structure-
O.mO-0 and 0*-0-^, with a third structure O-o-O, the 0-0 bond
being a hybrid between an ordinary and a three-electron bond.
Preliminary X-ray deffractions studies reveal close similarity
between the lattices of ?I03 and KN3, tetragonal, of the K3F3
type.
tost recent work on the alkali metal ozonides has been done
in this country by Whaley and Kleinberg who investigated the
sodium, potassium and cesium derivatives.1^ They agreed with
the Russians that assignment of the formulas ixia03, K03 and O03
to the colored, paramagnetic products is mo°t consistent with the
data obtained. However, chemical behavior of the product ob-
tained, from KOH, although mostly K03, would seem to indicate that
the red solid is not a single chemical species. Analysis of the
extracted cesium material even more strongly suggests the oresence
of another substance.
During the course of their investigations of the reactions
of the alkali metal hydroxides and ozone, the Germans (6-9) had
employed liauid ammonia as solvent and as a means for keeoing the
reactions cold. Manchot and Ksmoschulte (7) observed, further,
that the bright orange color is observed even with oure, dry
liauid ammonia by itself or such organic bases as methyl amine,
dimethyl amine, plperidine, toluidine or aniline, but that was the
extent of their investigations, Strecker and Thienemann, (13)
thirteen years later, carefully investigated this reaction of
ammonia and its derivatives. Experiments using nure, dry liauid
ammonia and dry ozone showed that an unstable ozonide of ammonia
is formed which on evaporation of excess ammonia is transformed
to ammonium nitrate (98fc) and nitrite (2$). Carius (14) and
Ilosvay (15) had already carried out o^onization experiments with
dilute aaueous ammonia solutions, the former obtaining MH4N03,
NH4N03 and H303 as oxidation oroducts; the latter could detect
only nitrate and nitrite. Strecker and Thienemann found that
ozone does not react nearly so completely with aaueous as with
liauid ammonia.
With hydro? ylamine, these workers found that hydro xylamine
nitrate, which is rapidly formed, is the only product. Hydrazine
hydrate yields essentially only nitrogen and water with small
auantities of hydrazine nitrate and presumably ammonium nitrate.
Hydrazoic acid is obtained in trace auantities. Tertiary amines
such as trim ethyl amine react violetnly, yielding exclusively
the N-oxide.
•
-• '
' ■
■
• .-v '
t
-86-
Papke, (16) studying relative rate? of oxidation of ammonia
in aqueous solution by ozone u^ing a*3 catalysts copper, cobalt
and nickel salts obtained hyc roxylamine a? the first and rate
determining step of the whole process.
In recent work on inorganic frEe radicals the decomposition
of ozone in aqueous solutions catalyzed by OH*" has been studied
by Weiss, (17) He showed that this can be accounted for by a
chain reaction catalyzed by OH" ion: :
03 +
0H~
j.
o3~
*
H02
■ s
o3 *
03 +
HO 3
— ».
20 2
o3
4-
OH /
fa
H03^J
OH
+
H02 +
OH —
o2
+
Ha0 "/
HO 2 +
KD
%
o2
+
1
H202 1
nu S
initiating step
chain propogation
termination
Further, the reaction between ozone and hydrogen peroxide was
explained quantitatively on this basis.
The most extensive work in the field of chain reactions of
or.one in aqueous solutions ha<? been done by Taube and Bray (18)
and by Volman (19). The interaction of ozone and hydrogen per-
oxide in acidic aaueous solutions can be presented by two overall
equations;
A. 03 + H202 * H20 + 20 3
B. 20 3 * 30 3
The mechanism involves the following steps:
(1) H202 * 03 — JU Oh + H03 + 03
(2) H02 * 03 — *<£-> OK + 203
(.3) OH + 03 — ^_> H02 + 02
(4) OH + Hs02 — £*-> H03 + H20
Reaction (l) is the chain-initiating *tep. Reactions (2) and (4)
are the chain-propogation steps for the overall reaction A and
reactions (2) and (3) those for reaction B. Whether the reaction
HO 2 + ha0a * OH + H20 + 02
occurs in the system will not only depende on the ozone/HgOg ratic
but also on the acidity, as it is likely to proceed via the 02"
ion. In acid solutions, e.g., 0.2N, it cannot occur to a measur-
able extent in this system, the pK of H0a being 2# As the reac-
tion 02~ + K20a * 02 + OK" + OK is likely to require
an activation energy of c_a. five kcal., it could be considerably
suppressed in the presence of ozone even at a pH of 2.
• ,.
-87-
A study wa" made by Taupe and Bray of inhibition of this "
reaction. All organic substrates (such as alcohols) act an in-
hibitors. Organic acids occupy a special Position. Chloride^
and bromides inhibit but not fluoride, in agreement with thermo-
dynamic data.
Some metal" are found to be effective catalysts, the greatest
effect being obtained when t^o conditions are fulfilled, (a) the
cation is oxidized by ozone by a monovalent valency change and
(b) the oxidized cation is reduced rapidly by H303. The net effect
is then a. larger stationary concentration of free radical", thus
leading to an increased rate of reaction. Colbaltous and cerous
ions fulfill these conditions and Taube and Bray have shown that
these cations induce the decomposition of ozone even in the absent
of hydrogen peroxide. They found it difficult to exolain the
effect of Cu 9 as the existence of an oxidation state Cu+3 has nou
been established beyond doubt. Uria6 expressed the opinion that
the basic interpretation of the effect of Cu+s ion might be
similar in principle to its action as promoter in the catalytic
decomposition of H303, Owing to a large e ^3/R factor, the
reaction Cu+S + HO a ■» Cu+ + H4" + 03 (or the corres-
ponding reaction with the 03~* ion) is extremely fast and could
occur in this sy&em, being followed by the oxidation of Cu+ by
ozone, accompanied in turn by the formation of OH radicals to
act as chain carriers, !,€,, Cu+(KOH) + 03 ^ 03 + Cu+S
+ OH- + OH. The lack of inhibitory effect of choride ions in
the presence of cupric ions is not surprising as it is in agree-
ment with the considerably increased efficiency of cuoric— chloro-
complexes as catalysts for the decomposition of hydrogen peroxide.
Hill (21,23) recently studied the decomposition of o*one
induced by cobaltous ion and postulated a mechanism involving
formation of the Co+30H"" ion-pair. He presented thermodynamic
data to support this proposed mechanism.
Yeatts and Taube (22) indicate that the reaction between
chloride ion and ozone in aqueous solution does not proceed
through free-radical intermediates such as OH radicals and chlor-
ine atoms, but reacts through the hypochlorite ion:
(a) 03 + CI"" ^ 03 + CIO"*
(b) CI- + CIO"* + 2H+ not rate- +
determining * 2
Szabo', (24) on the other hand, proposes a six-step mechanise
for the reaction whereby the intermediates are the free-radical «•••
CIO and C10a.
It is evident that much work remains to be done in the fieO*
of ozone reactions since most workers do not yet ->,gree on the
nature of these reactions or, indeed, on the ii& cure of the pro-
ducts obteined.
-
:3 -. ■
.. *
-88-
Blbliography
1. J. N. Friend, "a Textbook of Inorganic Chemistry," Vol. VII,
Part I, pp. 138-^55 Chas. Griffin and Co., Lim., (1924).
2. J. A. Partington, "A Textbook of Inorganic Chemistry", pn.
152-53, Sixth Ed., .KacMillan and Co., Lim., (1950).
3. W. Mane hot and W. Kampechulte, Ber., 40, 2894 (1907).
4. 'iurtz, "Dictionaire de Chimie ^ure et aprjliquee", Vol. II,
p. 721 (1868).
5. A. Baeyer and V. Villiger, Ber., 3_5, 3038 (1902).
6. A. Bach, ibid. . 35, 3424 (1902).
7. W, Manchot and W. Kamoschulte, ibid. . 40, 4984 (1907).
8. W. Traube, ibid. . 45, 2201, 3319 (191277
9. Neuman, J. Chem. Phys,, 2, 31 (1934).
10. I. A. Kazarnovskii, G. F. Nikolskii and T. A. Ableteova,
Doklady Akad. Nauk 9S3R, 64, 69 (1949).
11. Or. P. Nikol'skii, L. I. Kazarnov^ka;a, Z. A. Bagdasar'yan and
I. A. Kazarnovpkii, ibid.. 72, 713 (1950).
12* T. P. Wlraley and J. hleinberg, J. Am. Chem. Soc, 73, 79
(1951).
13. W. Strecker and h. Thienemann, Ber., 53, 2096 (1920).
14. L. Carius, Ann., 174, 49 (1874).
15. Ilo^vay, Ber., 40, 4985 (1907).
16. S. I. Paoko, Zhur, Priklad, Khim. , (J. Applied Chem.,) 23,
8 (1950); CA 44, 4319 (1950).
17. J. Weiss, Tran<=. Faradaj 3oc, 31, 668 (1935).
18. H. Taube and W. C. Bray, J. Am. ""Chem. Soc, 62, 3357 (1940 ).
19. D. h. volman, ibid., 73, 1018 (1951).
20. N. Uri, Chem. Rev,, 50, 375 (1952).
21. G. R. Hill, J. Am. Chem. Soc, 70, 1306 (1948).
22. L. R. R. Yeatts, Jr. and H. Taube, ibid., 71, 4100 (1949).
23. M. G. Alder and G. R. Hill, ibid.. 72, 1884 (1950).
24. Z. G-. Szabo, Acta Chem. et Phyc., 3, 20 (1950); CA 45, 8386
(1951).
-89-
HYDRAZlNxi: BY ELECTROLYSIS IN LIQUID AMMONIA
Robert N. Hammer January 13, 1953
Introduction:
Because of the great increase in the large-scale use of
hydrazine, considerable attention is being given to the improve-
ment of methods of synthesis. For the manufacture of hydrazine,
the low cost and availability of ammonia make it the logical '
starting material, in spite of the highly unfavorable thermo-
dynamic character of the reaction:
2 NH3 (liq) = N2K4 (liq) + H2 (g) ; &F = +41.0 kcal mole -1.
In principle, this non- spontaneous reaction may be driven to
completion by supplying low voltage electrical energy in an
electrolytic process in liquid ammonia.
Theoretical considerations:
Just as the electrolysis of many aqueous solutions results
in the formation of oxygen at the anode and hydrogen at the
cathode, so in acidic or basic liquid ammonia solutions the anode
and cathode products normally are nitrogen and hydrogen, as is
indicated by the following electrode reactions:
In acidic solution (i.e., NH4CT)
Anode : 4 NH3 = -^ Na + 3 NH4+ + 3 e (a)
+
Cathode: 3 NH4 + 3e = |H2 + 3 NH3 (b)
In basic solution ( i . e . , KNH 2 )
Anode : 3 NH2~ = * N3 + 2 NH3 + 3 e (c)
2
Cathode: 3NH3 + 3e = -H2 + 3 NHa (d)
Liquid ammonia and water have many similar properties,
but some of the fundamental dissimilarities of these two sol-
vents play an important role in electrochemical phenomena in
liquid ammonia. One of these is the unusually low value of the
standard free energy of formation of ammonia. The recent data
compiled by Jolly (l) gives a value of -2.7 kcal mole"-1- for
j\ Ff (298°K) ' whiif the corresponding value of water (2) is
-54. b4 kcal mole -1- — over twenty times greater. If the value
of /\ F„ for ammonia is corrected to -50°C (a temperature at
which many electrochemical investigations have been carried out),
it is found that ^ Ff (p23°K) = ""6.0 kcal mole . This leads
to an extremely low value for the reversible decomposition
potential of an ammonia solution in which the electrolysis pro-
ducts BTe N2 and H2:
Ed(223°K)= Z*| - ^%6uL = °-087 volt
-9 Cl-
in comparison, the reversible decomposition potential of water
is about 1.18 volts, or almost fourteen times pyeater.
If the experimental decomposition potential of liquid
ammonia were near this value, the possibility of hydrazine
formation by an electrolytic process would be exceedingly
unlikely. Fortunately, the measured decomposition potential
of solutions in liquid ammonia shows that in actuality this
property may be as large as several volts, even at low current
densities, owing to the existence of a large nitrogen overvol-
tage (3) .
Electrolysis of solutions of bases
Pleskov (4), in 1945, suggested that hydrazine as well
as nitrogen should be formed at the anode in the electrolysis
of basic liquid ammonia solutions. In 1950, Gessler and Ples-
kov published (3) the observation that hydrazine is formed at
a current efficiency of 5 to 10 per cent during the electrolysis
of a KNH2 solution. Even though all attempts by other investi-
gators to repeat this work have failed, there is reason to be-
lieve that such a reaction should be possible. As is indicated
in equation (c) , the primary anode reaction in the electrolysis
of an amide solution in liquid ammonia appears to be the dis-
charge of the NH2~ ion and the formation of nitrogen. However,
formation of hydrazine might occur through such a mechanism as
the following:
NH2~ = [NRV] + e (e)
2 [NHV] = HSN— m2 (f)
for which the overall half -re action would be
2 NH2~ = NgH* + 2 e (g)
Even the reaction of the [NH2"J radical with the solvent could
lea.d to hydrazine formation:
[NH2-] + NH3 = H2N— NH2 + 1/2 H2 (h)
Apparently no effort has been made to detect H2 in the anode
gas.
The combination of the anode half -reaction (g) with the
cathode half -reaction (d) shows that KNH2 is not consumed in
the process inasmuch as amide ions which disappear from the ano-
lyte are replaced by an equivalent amount formed at the cathode.
The overall reaction, then, represents the direct electrochemi-
cal conversion of ammonia into hydrazine and hydrogen:
2 M3 electrolysis 9 NaH4 + H2
Since hydrazine does not undergo electron reduction in liquid
ammonia (5) , it would seem unnecessary to prevent mixing of
-91-
the anolyte and catholyte.
electrolysis of solutions of acids
Another possible approach to the electrolytic synthesis
of hydrazine in liquid ammonia involves the use of a solution
of en acid such as I'IH4C1» Almost twenty years ago Howard and
Browne reported (6) that very small amounts of a reducing
agent, presumably hydrazine, are formed by the electrolysis
of Mi4Cl in liquid ammonia be tween platinum or graphite elec-
trodes. Again, efforts to repeat this work have failed. Con-
sideration of the mechanism of such a reaction suggests that the
formation of hydrazine in acidic solutions is less likely than
in basic ones. The formation of nitrogen at the anode in an
acidic solution, as is indicated in the equation:
4 NH3 = 1/2 N2 + 3 Mi4+ + 3 e (i)
may be the primary anode process or it may be the result of
secondary reactions between free halogens, or other' products
of the electrolysis, and the solvent. If the latter were the
case, the possibility of hydrazine formation would be more
promising. However, there is reason to believe that this does
not occur. The anode overvoltage of nitrogen in acid solutions
is practically uniform, regardless of the anions present in the
solution, except in the case of NH4I (3) . This seems to indicate
that the primary electrode process is the same in all cases
end is probably the direct electrochemical oxidation of the
ammonia molecule, inasrnuch as the concentration of NH2"" ions
in acidic solutions must be exceedingly small. Furthermore,
there has never been found in the products of electrolysis
any free halogens, other than iodine, or eny intermediate
products of their interaction with the solvent, such as NH2C1,
NGI3 , etc. In the case of electrolysis of wH4I solutions,
Gessler and Pleskov reported (3) that about 20 per cent of the
current is consumed in the separation of iodine. This might
be expected from a consideration of the data in Table I, where
it may be seen that the differences between the normal poten-
tials of the halogens and the potential of nitrogen separation
considerably exceeds the 1.4 volt overvoltage (3) of nitrogen
in all cases except for iodine in acid solution.
Table I . The difference of normal potentials of halogens
and the potential of nitrogen separation in
liquid ammonia (3)
Solution CI? Bra £2
%2"~~ Ex2 Acid 2*03 1-90 1*45
% — ^Xp Alkaline 3.39 3.19 2.81
-92-
Arc electrolysis.:
Still another approach to the electrochemical formation
of hydrazine in liquid ammonia involves arc electrolysis.
Experiments by Ingraham (7) with 400C volt direct current
discharges between a platinum anode and a catholic surface in
liquid ammonia have shown that hydrazine is formed in yields
up to 0.2 per cent, based on ammonia disappearance. The yield
increases rapidly with increasing current density. Since the
yields are in excess of those calculated by Faraday' s laws for
the simple discharge of the amide ion, Ingraham suggested that
hydrazine is probably produced by a chain reaction operating
in the gaseous discharge above the liquid surface.
Bibliography:
1. W. L. Jolly, Chem, Revs. 50, 351-61 (1952)
2. National 3ureau of Standards Circular, "Selected Values
of Chemical Thermodynamic Properties," U. S. Govern-
ment Printing Office, Washington, D.C., 1952
3. N. w. Gessler and V, A. Pleskov, Zhur. Fiz. Khim. 24,
445-54 (1950)
4. V. A. Pleskov, i^cta Physicochim U.R.S.S. 20, 578-87 (1945)
5. G. W. Watt, Chem. Re-re. . 48, 301 (1950)
6. D. H. Howard, Jr., and A, T". Browne, J. Am. Chem. Soc. 55 .
3211-14 (1933)
7. T. A. Ingraham, Can. J. Chem. 30, 168 (1952)
-93-
NAPHTHAZARIN COMPLEXES OF THORIUM AND h$R& iSARTH M2TAL IONS.
Melvin Tecotsky Thesis Report January 20, 1953
Solutions o£ thorium salts show no absorption in the range
2000 - 10,000 A , thus thorium cannot be identified by direct
absorption spectra measurements (3 ) . In general, there is a lack of
color reactions for thorium. Formanek (1) has shown that when the
dye alkannin is added to thorium salt solutions, one obtains a vio-
let solution showing maximum absorption in the range 6055-6084A0.
Recently, in seeking a more satisfactory method for the determi-
nation of beryllium, Underxvood and Neuman (5) developed a method
first using alkannin and then an equivilant procedure using naphtha-
zarin. Since alkannin is a substituted naphthazarin, the two com-
pounds would be expected to behave similarly. These recent in-
vestigations of the naphthazarin complexes of beryllium have indi-
cated that a study of the reactions of other met?l ions with this
reapent might prove fruitful. Inasmuch as alkannin and thorium
had been shown to give a color reaction, it seemed logical to in-
vestigate the thorium naphthazarin system from this point of view,
TH^ RdAGSNT
Naphthazarin is prepared by the Friedel-Craf ts Reaction
of hydroquinone and maleic anhydride (4). The product is recrys-
tallized from high boiling petroleum ether and is obtained as
brown needles. Naphthazarin has the structure shown in Figure A.
OH 0
dH 5
Fig. A
When dissolved in an organic solvent such as dioxane or ethanol,
naphthazarin produces a red solution. Upon the addition of various
metal ions the color of the solution changes, the color change
being from red to reddish-blue or violet. Absolute ethanol has
been used as a solvent for naphthazarin in this investigation.
Metal ions have been used as aqueous solutions of their nitrate
or chloride salts.
Upon the addition of a solution of thorium nitrate to a naphtha-
zarin solution, there is a distinct color change from red to violet
or purple, due to the formation of a thorium-naphthazarin complex.
Solutions containing as little as 1 X 10 3 mg of thorium ion per
ml. give visible evidence of complex formation. The color pro-
duced by rare earth ions with naphthazarin is a deeper red than
the naphthazarin alone. The color change produced by rare earth
ions is no where near as pronounced as the color change produced
by thorium ion.
-94-
THORIUM-NAPHTMAZaRIN complex
.ABSORPTION SPECTRA
The absorption spectrum of naphthazarin inethanol is character-
ized by twin peaics at 4875 A5 and 5166 A and an inflection
point at 5500 a. The spectrum of the thorium -naphthazarin com-
plex |s also characterized by two peakg . The first appears at
5700 A and the second appears at 6185 A, This spectrum also has
a point of inflection at 5375 A. The spectrum of the complex
is sufficiently different from that of the uncomplexed reagent to
permit a study of the complex.
To determine the number of complex species present in solu-
tion, solutions of the thorium-naphthazarin complex were prepared
in different mole ratios. Using ithe Gary Recording Spectro-
photometer, it was observed that the curves obtained for all the
solutions coincided quite well, thus indicating that only one
colored complex species was present in solution.
THi£ MOLAR RATIO OF THd COMPLEX
An attempt to elucidate the empirical formula of the thorium-
naphthazarin complex was made through the use of the method of
Continuous Variations (2,6) . Use was made of the difference in
optical density between a naphthazarin solution and a solution
of the complex of the same molarity. This difference in optical
density, known as Y, between the complexed and uncomplexed re-
agent was plotted against concentration. The molar ratio where
Y is a maximum is the molar ratio of the complex.
Similar results were obtained using the Gary Recording Spect-
rophotometer and the Beckman Spectrophotometer. When the values
obtained were plotted, the maximum appeared at 34 mole % thorium
and 66 mole % naphthazarin. This indicates a two to one complex,
two naphthazarin molecules to one thorium atom.
Further evidence for a two to one naphthazarin- thorium com-
plex was obtained by the application of the Molar Ratio Method.
It has bien reported by Yoe and Jones (7) that for a very stable
complex, a plot of optical density against molar ratio of component
B to component A, with A constant, rises from the origin as a
straight line and breaks sharply to constant optical density at
the molar ratio of the components in the complex. In every case
the curve broke at a ratio of two moles of naphthazarin to one
mole of thorium. This indicates a two to one complex and agrees
wi-th the data obtained from the method of Continuous Variations.
STABILITY OF TH£ COMPLEX
A stability study of the tnorium-naphthazarin complex over
a period of forty- three days indicated a very small ch^n^e in
optical density over this period of time. The complex a/ooears
to be stable enough in solution to work with for normal intervals
of time without appreciable decomposition. When the complex was
exposed to the direct rays of the sun, a much more r^pid decom-
position was found to take place., The spectrum of the complex
was shifted to shorter wave lengths, and there was a considerable
decrease in optical, density.
-95-
LAKE LIKE CHARACTERISTICS OF TIIE COIIPLEX
Upon the addition of ammonium hydroxide to a violet tborium-
naphthazarin solution, containing excess thorium, a precipitate
of thorium hydroxide forms. Upon agitation the solution becomes
colorless and the precipitate becomes violet in color. If an
excess of naphtbazarin is present, J:he precipitate forms and the
resulting solution is blue- in color. Naphthazarin is blue in a
ba<-,ic solution. The thorium-naphthazarin material becomes adsorbed
on the thorium hydroxide precipitate, in a manner similar to the
adsorption of aluminon reagent on an aluminum hydroxide precipitate.
The colloidal properties of the thorium-naphtha zar in material have
been shown through dialysis and flocculation of the material by
strong electrolytes.
ADHERENCE TO BEER1 5 LA'f
Various studies of the adherence of Beer!s Law of solutions
of the thor ium-naphthazarin complex have been made. '.Then the
solutions prepared contained less than a ^rro to one ratio of
naphthazarin to thorium, no adherence to Beer's Law was found, to
exist. When, however, the ratio of naphthazarin to thorium was
two to one or greater, the solution of the complex was found to
adhere to Beer's Law in the range 10~52i/l, Thin adherence to
Beer's Law is important since it may provide s method for the
nuantitative determination of thorium.
RARE EARTH ?-IAPTTT-IAZARIN COIIPLEX
ABSORPTION SPECTRA
The absorption spectra of solutions of lanthanum, praseodym-
ium, neodymium, samarium., gadolinium, erbium, vn^ yttrium ions
combined T-rith naphthazarin -rere examined Trith the Cary Recording
Spectrophotometer. The spectra obtained were all -similar in nature.
The rare earth-naphthazarin spectra does not have as pronounced a
characteristic peak as the thorium- naphthazarin complex does. The
characteristic oeak for the rare earth-naphthazarin complex is
located between 6000 and S05C A. This is over 100 A from the
thorium-naphthazarin ■neali. This fact may Trell make it possible
to determine thorium in the presence of rare earths through the
u s e of naph th a ? ar i n .
The rare earth naphthazarin complexes exhibit the same prop-
erties ns the thorium-naphthazarin complex; such as one complex
species, a tT,o to ' one ratio of naphthazarin to rare earth ion,
adherence to Beer's Law and lake like characteristics.
An attempt has been made to evaluate the stability constants
of the complexes formed. The determination of thorium in the
r>re*.e*ice of rare earths is being attempt s-'1 at the present time.
-96-
BIBLICCrRAPHY
t
(1) Formane!;, J., S. anal, Chem. J£, 673 (1900).
(2) Job, P., Ann. Chera., (10) 1, 113 (1923)
(3) ":oeller, T., Schweitzer , G.K., and Starr, B.D., Chem. Rev.
ia, 63 [19^5.
(H) Toribara, T.Y, and Underwood, A. L. , Anal, Chen. 21, 1352
(19^9).
(5) Underwood, A.L. , and Newman, T'r.F., Anal. Chen. 21, 1345 (1949)
(6) Vosburg, *i.C., and Cooler, G-.H. , J. Am. Chen. Soc . , 6j, 437
(19*4-1).
(7) Yoe, J.H. s and Jones A. L. , Ind. Enp:. Chen., Anal. Ed. l£, 111
(19^0.
-97-
THESIS REPORT
Perry Kippur January 20, 1953
INTRODUCTION
The potential use of hydrazine as a specialty fuel and the
commercial exploitation of certain hydrazine derivatives as a result
of their biological activity has recently increased the amount of
research activity in this field. These efforts have been directed
towards the development of new procedures for both the preparation
of hydrazine and the synthesis of hydrazine derivatives that may be
employed as pharmaceuticals and agricultural chemicals.
The purpose of this investigation has been first, to develop a
practical method for the preparation of thlocarbohydrazide, a com-
pound which exhibits insecticldal, herbicidal, fungistatic and growth
regulating properties and secondly, to condense cyanamide with cer-
tain hydrazides in order to extend further the general type reaction
involving the addition of hydrazine, N-substituted hydrazines and
hydrazides across the triple bond of the cyano group.
X X NH
(l) H3NC3N + -CNHNHs -* -CNHNHCNH8 (where X* 0, S, or NH)
I. THI0CAR3CWYDRAZIDE
Historical
Thlocarbohydrazide has been orepared by the elimination of
hydrogen sulfide from hydrazinlum dlthiocarbazinate (l) and the
hydrazinolysis of S-methyldithiocarbazlnate (2), diethyl xanthate (3)
and thiophosgene . (l)(4)
The amphoteric character of thlocarbohydrazide is evidenced by
its solubility in acids and bases and the formation of metallic and
acid salts in suitable solvents, Thlocarbohydrazide enters into the
reactions typical of the hydrazide group l) reduction of silver ion,
2) solvation of compounds containing the cyano group (HOCN, H3CN) to
produce normal or substituted N-carbamyl or N-thiocerbamyl deri* •.
vatives, 3) solvolysis of acyl halides to yield N-acyl derivatives
and 4) solvolysis of carbonyl compounds to yield the corresponding
thiocarbohydrazones. Although alkylation of thlocarbohydrazide re-
sults in the formation of S-substituted derivatives, N-alkyl and
N-aryl derivatives may be prepared by the reaction of certain thlo
compounds (C3S, CSC1S) with substituted hydrazines. The 1,5-di-alkyl
and/or aryl thiocarbohydrazides may be oxidized to " thiocarbazones"
(RN=NC (S) NKNHR) which are useful in the quantitative, color ime trie
determination of metallic ions.
, . , ■ '.
» ..••■■
■ ':.' ■•; .
V . <
i .:'. :1 ' ■ '
, >,
... -^ r, 1 .,.
v
.. ■■■•>. .. - '
r> ' ' ■ - '
■ -■:■■:
V
-.;.■■ , : - •
' - - v -
-9&-
Experimental and Discussion
The solubility of thiocarbohydrazide in water, elshanol, chloro-
form, carbon tetrachloride and hydrazine hydrate has been measured
quantitatively; appreciable solubility is found only in the case
of hydrazine hydrate. Thiocarbohydrazide is non-hygroscooic up to
and including a relative humidity of 9tf£, It undergoes gradual
thermal decomposition at 110°C. The dH' of a saturated solution at
room temperature Is 6.95. From on x-ray diffraction pattern of
thiocarbohydrazide, the d-spacings have been calculated.
A practical method for the Preparation of thiocarbohydrazide by
the direct reaction of carbon disulfide and excess aqueous hydrazine
has been developed in which a yield of about 60^ is obtained. The
t>rocedure requires the use of approximately a 5:1 mole ratio of
hydrazine: carbon disulfide. For the quantities of starting materials
employed (0.2 mole OS,) in this investigation, it wes found that a
reflux time of from 1/2 to 1 l/2 hours is desirable. Acidification
of the reaction mixture results in the decomposition of hydrazinium
dithiocarbazinate which is formed by the initial combination of
hydrazine and carbon disulfide. The use of dilute aqueous solutions
results in a reduction of the yield of thiocarbohydrazide. The
employment of an alcoholic solvent offers no advantage over the use
of an aqueous Solution, Additional quantities of thiocarbohydrazide
(as much as 15^) are obtained by refluxing the mother liquor,
subsequent to removal of the product from the initial reaction
mixture. If the reactlon:ls carried out on a larger scale (using
molar quantities of CS2 rather than 0.2 mole), an Increase In the
yield of several percent is obtained.
This new procedure Possesses several advantages over previously
reported methods: l) better overall yields (60#) are obtained, 2)
commercially available raw materials are employed (other methods
require the Preparation of special starting materials) and ?) the
starting materials are not as difficult to handle as in the prepara-
tion of thiocarbohydrazide by the hydrazinolysis of thlophosgene,
(l)(4) The new procedure is believed to be adaptable to a semi-
continuous and/or continuous operation.
A mechanism for the elimination of hydrogen sulfide from
hydrazinium dithiocarbazinate to give thiocarbohydrazide has been
proposed. This mechanism involves the formation of an Intermediate
dithiocarbazinate ion which could then react with hydrazine prior to
the elimination of hydrogen sulfide. Credence Is lent to this
mechanism as the excess hydrazine, which is required to give good
yields of thiocarbohydrazide, favors each step of the postulated
mechanism.
II. ADDITION OF CYANAMIDE TO KYDRAZIDJ2S
Historical
The addition of cyanamide to hydrazides (equation l) is similar
to the condensations of cyanate and thiocyanate with hydrazides.
Certain analogies exist, namely, first, the reagents attach them-
selves to the p-nltrogen atom of the hydrazide group, secondly, the
- I ' v,
«
(-: ■• .•« - (
... \. -... .,,,»-. i. .. • ■,->•; •,< ' ■■ ■
' ■ - " * -. .. ■■-■■ •■: •■: - » .-. .- •*,.»•*'•'./
■ r .. *
:t>
^ *.-.*•
, . r • ■- .
* > •
r 11 *
S'k.J'..-'V
• . . . ',' *■»■ - . i
* !
■5 ■;•• • • ■■•;". ,i
*..§ ;■' - ,:
, M'
f-3-,4
■ iV i". ■•' • ■ ■ ■
-. i' : •.'•■ ■ ■
it - *
■, •-,,./ -
I
I »;
•"J
.. r, > - • : to
■
. .. <
> .- ■ <■ y
... .J. r:
«. ■ ,-"
e*r:v- '?
:r;;> j i.;o&) fito^t ::*:/•:':',' v-1
-?:.
. ...?... .;; v>' 2 J/J -'•.-'- ,r-- '
-99-
proposed mechanisms involve either rearrangement of an intermediate
salt or molecular addition of the hydrazlde across the carbon-nitro-
gen triple bond and thirdly, the condensations are pH deoendent as
a result of the stability of the reagents and/ or products. It is
more probable that cyanamide addition occurs according to the oro-
posed molecular addition mechanism as the existence of a cyanamide
salt in aqueous solution is improbable as a result of the weakly
acidic character of cyanamide. Furthermore, cyanate and thiocyanate
are relatively stable in alkaline solution as compared to cyanamide
which readily undergoes polymerization to dicyandiamide under these
conditions.
The orocedures employed to affect cyanamide addition to hydra-
zldes entail: l) the use of an aqueous solution of cyanamide, 2)
the employment of acidic aqueous solutions of sodium or calcium
cyanamide, 3) the decomposition of nitrosoguanidlne In aqueous
solution and 4) the hydrazinolysis of S-alkylisothiouronium salts.
Although there are numerous reports in the literature con-
cerning the addition of cyanamide to amines to oroduce guanldines,
the preparation of aminoguanidine and Its N- substituted derivatives
represents one of the two instances^ of direct cyanamide condensation
with hydrazine and substituted hydrazines. The preparation of bis-
guanylhydrazine dinltrate has also been described. (5)
Experimental and Discussion
The addition of cyanamide to semlcarbazide, thiosemicarbazide,
aminoguanidine and carbohydrazlde has resulted In the preparation
of the expected N-guanyl derivatives. These compounds have been
characterized by conversion to the corresponding Picrates and/ or
oicrolonates. The new compounds "which have been prepared include:
1) N-carbamyl-N^guanylhydrazine hydrochloride
2) N-thiocarbamyl-N»-guanylhydrazine hydrochloride
3) Bisguanylhydrazine d Hydrochloride
4) G-uanylcarbohydrazide dihydrochloride
III. BIOLOGICAL ACTIVITY OF THE HYDRAZINE DERIVATIVES OF THE 1
CARBONIC,' THIOCARBONIC AND AMKONOCARBONIC ACIDS
A number of hydrazine derivatives have been evaluated for their
antituberculin activity and also for their effect upon the blood
nressure of hypertensive animals; , however, none of the materials,
which have thus far been tested, display any significant activity.
The pharmacological evaluations have been carried out by personnel
in the laboratories of the Eli Lilly Corrmany.
Only a small number of hydrazine derivatives have thus far been
evaluated for their insecticidal activity by the Department of
Entomology at the University of Illinois. However, thlocarbohydra- '
zide and l-pVienylthiocarbohydrazlde have been found to be toxic
to roaches.
A group of twenty-four hydrazine compounds consisting mainly of
carbonic, thiocarbonic and ammonocarbonlc acid derivatives have been
subjected to an evaluation of their possible applications in the
l-f --, -.' -> If. ■ J ■
■ • f
_, ... ... ..
.■"■•.; ■ ■••■
7 ,,;)■..!>,,.
■> HA'.
- ••. 1 .■-,• t. r. .' * -.-.■■
* :"> ? \.
-100-
agrlcultural chemical field. The 86 tests have been carried out in
cooperation with the Department of Agronomy at the University of
Illinois.
Results have shown that methyl hydrazine sulfate, hydrazine
sulfate and thiocarbohydrazide are effective as contact herbicides.
Furthermore, methyl hydrazine sulfate, hydrazine sulfate, thio-
semicarbazide, thiocarbohydrazide and 1-ohenyl-thiocarbohydrazide
show promise as fungicidal or fungistatic agents. A significant
effect upon the growth of certain plants has been exhibited by
thiocarbohydrazide, thlosemicarbazide snd 5-aminotetrazole. A small
number of hydrazine compounds display crop plant defoliant activity*
An interesting observation which has been made during these biological
tests is the production of albinism by bisthlocarbamylhydrazine,
5-amino-tetrazole and 1, 2-diacetyl-.'% 5-diamino-l, 2, 3, 5-tetrahydro-l,
2,4-thiadiazole.
The results of these biological tests have repeatedly demon-
strated that the presence of a thiohydrazide group -C(S)N2H3, en-
hances the biological activity of the compounds which have thus far
been evaluated.
BIBLIOGRAPHY
1. Stolle and Bowles, Ber. 41, 1099-110? (l°08).
2. Scott, Doctoral Dissertation, Univ. of Illinois, 195?.
3. Guha and De, J. Chem. 3oc, 125, 1215-18 (1924).
4. Autenrieth and Hefner, Ber. 58, 2151-55 (l°25).
5. Thiele, Ann., 273, 133-44 (1893).
I. . •■>
-101-
Ion-Pair Formation in Acetic Acid
II. ::. Jones Feb. 24, 1953
In one of the earliest investigations on the
properties of acetic acid as a solvent, Raoult (l)
found that the freezing point depression of solutions
of alkali acetates correspond to ideal behaviour.
Later investigations by Piaoult and Recoura (2) and
Beckmann (3) confirmed and extended these observations.
At that time wo r hers in physical chemistry were chiefly
concerned with what was considered to be anomalous
behaviour of aqueous solutions of ionic compounds. The
theory of Arrhenius had not yet been completely accepted;
some of its glaring inadequacies were much discussed (4).
At the same time some of its triumphs, such r s the good
agreement between the degree of dissociation determined
by freezing point methods and conductivity studies (since
shown to be fortuitous), seemed to indicate that the
theory was in the main correct. The fact that the theory
seemed incapable of giving a. complete explanation for the
properties of ionic solutions led to the development of
tTro different schools of thought. The first took an
almost purely thermodynamic viewpoint and found an able
leader in G-. N. Lewis. The second school of thought
attempted to develop a solution on the basis of statistical
mechanics. The first work of any consequence along the
-102-
latter line was attempted by S. R. Ililner (5)-(8). The
first widely successful treatment, by Debye and Huckel (9),
offered several advantages over the earlier and somewhat
similar work, of I liner.
The theory of Debye and Huckel is familiar at least in
a qualitative sense to most chemists. Less frmiliar are
the attempts of others to overcome certain difficulties
and approximations which are contained in the Debye-Kucliel
theory. One of these approximations was their solution
of the Poisson-Boltzmonn equation, relating the distribution
of charge and the potential rt a point in the solution.
Soon after the appearance of the Debye -Huckel theory,
Bjerrum devised a method which did not require a solution
of this particular equation. This method of Bjerrum, and
subsequent developments of it by other authors, has become
known as the method of "ion-pairs" and is capable of
furnishing a consistent picture of ionic solutions.
It is most frequently applied to solutions in non-aqueous
solvents of low dielectric constant although it is by no
means limited to there.
Bjerrum assumed (10) that every ion within a
miniumum distance of another ion of opposite charge is
paired with that ion. Such an "ion-pair" acts as a single
particle. By determining the probability of such ion-pairs,
Bjerrum showed that such a theory could account for the
activity coefficients of salts and the radii of ions in
-103-
dilute aqueous solutions. The methods outlined in this
paper required tedious numerical 7rorI: and were much less
convenient in actual use than the simple formulae of
Debye and Huckel .
Fuoss and Kraue ( 11) (12) later used this principle
of Djerrum's in the explanation of the properties of
electrolytes in solvents of low dielectric constant.
By combining the Ostwald dilution law vrith the ideas
of Djerrum, Fuoss and Kraus developed a method of
estimating the degree of dissociation and the limiting
equivalent conductance of these solutions. Extensive
experiments by Kraus and coworkers have shown that the
general pattern of behaviour of electrolytes in a large
number of non-aqueous solvents is accurately predicted
hy this theory.
Although earlier studies in acetic acid attempted
to use the principles outlined by Fuoss and ICraus, the
first rigorous application of their theory to this
solvent was made by Griswold (1943) in work which has
not yet been published. Griswold collected very precise
information on the effect of one salt on the solubility of
another and also ms.de a careful survey of the literature for
similar data. In all cases the concept of ion-pairs
was capable of explaining the effects quantitatively.
Later Jones, under the direction of Griswold, provided
an independent method of checking the assumption made
by C-riswold. In addition, these later studies provided
-104-
quantitative vn.ln.es for both dissociation constants and
activity coefficients of the various solutes studied
earlier bv GrisTrold.
(1
(2
(3
(4
(5
(6
(7
(C
(C
(10)
F. II. Raoult, Ann. Chim. Fhys. (6), 2, 72 (1884)
F. II. Raoult and A. Recoura, Z. physil:. Ohem . , 5,
424 (1890)
E. Bechmann, Z. physilc. Chem . 57
G-. N. Lewis, %. physik. Chem. 70
S. R. Kilner, Phil. Rag. (6), 25
3. R. liilner, Phil. Vcg. (s), £5_
S. R. i'ilner, Phil. Hag. (o), 5J5
3. R. " ilner, Phil. Ran-. (6), 85
136 (1007)
212 (1910)
551 (1912)
742 (1915)
214 (1918)
552 (1918)
D. Debye and E. Kuckel, Physilc Z., 2A, 185 (1923)
1!. Bjerrum, K. Danske . Selsh.
(1923)
(ll) R. Fuoss and C. Rrau
!Iat. -fys. Redd. I, #9
J. At.. Chem. 3oc . , 55, 1019,
(12)
2587 (1933)
R. II. Fuoss, Chen. Revs., 17, 227 (1935)
-105-
R. L. Rebertus
larch 3, 1953
The Irondll ) Complex
Chemical Properties. The Irondll ) ver senates, NH4FeY • H20 and
HFeY, may be prepared according to the procedure of Brint zinger and
coT-rorhers^. A solution of the ion, FeY , is stable to acidified
permanganate. Only partial reactions occur upon sulfide or thiocynate
addition. Ho reaction takes place with o-phenanthroline , phthalate, or
benzoate . Alkali hydroxides precipitate ferric hydroxide, and cup-
ferron, CsH5IT'iI0,0NK4, also precipitates the iron. Solutions of FeY
are stable toward decomposition Trhen st®red in dark or red containers,
but in sunlight the solution becomes colorless due to the reduction
of the complex.
Physico-chemical Studies. The titration of HFeY ^rith sodium
hydroxide reveals that it behaves as a strong acid. However, at high
pH values a second inflection, accompanied by a color change, occurs
due to the formation of the complex, FeYOH", This titration curve is
compared Trith that of the T-real; dibasic acid, I!a2H2Y, in Fig. 1.
—
O.Of Mi i^Aiii
^
.
.— -*^
1
PH6
1
1
1
■7
,/.
1
... | >.' |V > I gpTfl
±
z
.J
Fig.l. Titration of HFeY and Ba&E2Y irith Na0H-
»• *,.w , hi
' i .
•106-
The absorption of FeY~ at pH 2.5 is such that the solution appears
yellow, but a pH 9.0 the orange color of FeY0"-I= appears. The method of
Job as modified by Vbsburgh* was used to determine whether FeY is the
only complex existing in acid solution. Two solutions of equal molar-
ities, one of ferric ion and one of ethylenediaminetetraacetic acid,
T'ere mixed in differing rrtios, and the optical densities of the
solutions were measured at selected wave lengths. A plot of Z, the
difference between the observed density and that calculated assuming
no reaction, against x, a composition function, reverled only the 1:1
complex.
The dissociation constant of the ion, FeY"", was determined by the
radioactive indicator method. A solution containing 0.025 molar FeY",
0.50 molar K, and enough sodium perchlor^te to give an ionic strength
of 2.02 was held at 25° to allow equilibrium dissocirtion. Then equal
volur.es of this solution and of 0.025 mol~r Fe* (010^)3 were mixed. An
instantaneous exchange of 15^ was observed. This inrt^ntaneous ex-
change results from the mixing of the radioactive ferric ion with
non-rr>dioactive ferric ion from the equilibrium, dissociation of the
complex. This vrlue is related to the fractional dissociation of the
complex hy
f = d(a + b)
a + db
where f is the fractional exchange, d is the fractional dissociation
of the complex in the original equilibrated, solution, and a and b are
the total concentrations of added ferric ion and complex ion in
solution after mixing. In this case a=b, and d calculates to be 0.081.
Thus in the original solution the ferric ion concentration is 0.002
molar. In highly acid solution the predominating equation for the
dissociation is
!" ■-: ' !,j " j •> ■
. I
I
■■' : '.
I
calculate K^:
-107-
ii
FeY~ + 4 H+^± Fe+++ + K4Y
From this equation the following equilibrium expression can be used to
(Fe+++)2 ' / "
[(FeYJ)0- (Fe+++)] [(H+)o"4 (Fe ' " »4 = ** *
where Kd is the dissociation constant of FeY*", Ka is the over-all
dissociation constant for the acid, H4Y, (i0~ ), and the zero
subscripts refer to total concentrations in the original solution be-
fore dissociation. The value of K^ is calculated to be lO"*^4. Schwarz*
enbach and Heller report a value of 10 i~° from an e.m.f. study of the
equilibrium constants.
Exchange between the complex FeY" and Fe* (0104)3 proceeds slowly,
the half-time of exchange being several hours. A detailed investigat-
ion of the kinetics of exchange yielded the following rate law;
R = Q. 4,7 (FeY-) (K^)3 + 0.015 ( FeY") (Fe+++) + 0.14(FeY~) (FeOH++)
In highly acid solutions the first term is the predominating one,
whereas in more basic solutions the second and third terms contribute
mostly to R.
Structure. The iron versenates are octahedral; each has
optical isomers.
C\ /!
J?**/ J // d*"™//
o
U. I ///
/
Fig, 2,. - Possible Structures of -Iron Versenates.
Commercial Application. — — Iron(ni) ver senate is commercially
available. This complex has the ability to liberate iron at such a
rate thrt it is easily assimilated oy citrus trees, and it has been
-108-
H±e Irp,n,(ll.) Complex
Chemical Properties -No solid salts of this complex have been
prepared. Oxygen in solution readily oxidizes this complex to FeY~*.
Ammonium hydroxide (pK 9) does not decompose it, but alkali hydroxides
give mixtures of ferric and ferrous hydroxides. The complex, FeY^, is
^Iso decomposed by sulfide ion and o-phenanthroline . Ho reaction vrith
ferricyanide ion occurs.
Physico-chemical Studies The acid, H3FeY, behaves as a strong
acid and only a single inflection in the tirrrtion curve is noted.
However, when the solutions are exposed to air, an inflection occurs
after only one equivalent of base is added due to the formation of
HFeY .
The spectrum of FeY= chows no absorption in the visible region.
The dissociation constant of the ferrous complex was determined
by a procedure quite different from that used to determine the constant
for the iron (ill) complex. The iron(ll) complex is of lover stability
so that in solutions of intermediate acid concentration there is
extensive dissociation with concomitant formation of weak acids from
the Y~ ion. A calculation of the dissociation constant is -oossible
from a determination of the actual hydrogen ion concentration in a
mixture of the complex and a known amount of acid and a knowledge of
the acidity constants for ethylenediaminetetraacetic acid6. The
following equations may be writtpn,
! + 2 x 10l6.42(K+)2 + 10l0.2s R+] '+ (Hs^_)
(Y~~) = (H+)o - (H+) - (HSO.)~
1021.69(H+)4 + 1019.56(H+)3 + 1016.72(H+)2+ 1010.2<^
^e+P =|I0 Hx^"4 „ (Y==)[1021.09(H+)4 + lol0.09(H+>3 + 10l6.42
(K+)2 + 1010.26(H+) + 1]
•-• i °i. 1 :■">: ■: ;>!
' ' ' '■ '■• •' C ■ ' . •
Hi; '• '■ v
'-.'•:•' ("■> ■
Ito'j to
•• ' . -.
t ■■■ ■ • • L Vt ...
f .-, i~
"to ' ■.■-,.
' : ' ' ■ '■''■'.■ ■■""; . :'
; ' ' : ■ ' * ::.'■ ■• -• ■
r — '
■ I...:* >-><• -.
.1 ■:.- r '•
) „,. . .. '
• : / f ■
;."!, ' r
■\ -: .
rr-W'
-109-
(FeY=) = (FeY=)0 - (Fe++)
In practice, sodium sulfate, sulfuric acid, and sodium ethylenediamine-
tetraacetato-ferrate(ll) are coexistent in the solution, and the
dissociation constant may. be calculated from the above relations.
Kd " <*>*+) (*~) __ m 10-14
(FeY=)
The half-time of exchange between FeY= and Fe*++ is less than a
minute. This is additional evidence for the lesser stability of the
FeY~ ion as compared to the FeY~ ion.
Li terrture References
1. S. S. Jones and F. A. Lonr, J. Phys. Chem . J56, 25 (1952).
2. K. Drintzinger, H. Thiele, and U. luller, 2. anorg. allgem . Chem.
2£1, 285 (1943).
3. P. Job, Ann. Chim. do) 9, 113 (1928).
4. "'. C. Vosburgh and C-. R. Cooper, J. Am. Chem. Soc, 63, 437 (1941 )
5. G-. Schwarzenbach and J. Heller, Helv. Chim. Act? 34,, 576 (1951).
5. G-. Schvrarzenbach and H, Achermann, ibid. , 30, 1798 (1947).
" *" ..."
'
. i • .'■■■■■
r..|>Ii>V
>.T. '
■ "• , ■
>-
-110-
■Amphoterism in Hon- a cue our Systems
Harold J. Ilatsuguma llarch 3, 1953
I . Introduction:
.Amphoterism has been studied very intensively in the
field of aqueous solutions. The actual mechanism of the dis-
solution of hydroxide precipitates in excess base has long
been a point of argument. Three concepts have been formulated.
(2)
The first, advanced by riant zch f considers the process
to be one of peptization. Solutions formed by such a
process would be colloidal in nature, and indeed both true
?nd colloidal solutions of chromium (ill ) hydroxide can be
prepared under suitable conditions. Bredig, on the other
hand, considered the amphoteric compound capable of acting
either as an extremely weak acid or bare. These two theories
(3)
hrve been superceded by the third, developed by Pfeiffer
This latter theory and the one most TTidely supported at the
present time explains the dissolution of amphoteric hydroxides
as a coordination phenomenon in which the .amphoteric element
goes into solution usually as a soluble anionic hydroxo complex
(4,5,6,7,8,9,10,11,12,13)
Scholder and coworkers have carried
out an extensive series of investigations on the amphoteric
behavior of several elements and their work confirms Pfeiffer' s
theory .
i ' '• ! i -
-111-
Recently amphoterism has been investigated in other
ionizing, water-like solvents. It has been found that
aluminum, zinc, tin, lead and chromium, all amphoteric
in water, also possess this property in various other solvents.
II . Airohoterism in Specific Solvents :
The phenomenon of amphoterism has been noted in almost
every non-aqueous solvent investigated. Many elements carry
over amphoteric character from one solvent to another. In
each case of amphoterism a coordination mechanism has been
postulated to explain observations.
The Copper (il) ion x' ' exhibits amphoteric character
in liquid ammonia. Copper (il) nitrate dissolves in liquid
ammonia to give a. deep blue solution. VJhen potassium amide
is added an olive-green precipitate is formed; this nrecipitate
dissolves upon addition of pxcess amide. Similar behavior is
noted with the amides of the following elements; strontium (II ) ,
beryllium (II ) , zinc(ll), silver (i), lead(ll), aluminum (ill )
and gallium (ill ) . Tost, if not all, of these substances are
initially only slightly soluble in liquid ammonia, but will
dissolve when an excess of pota.ssium amide is added. Some of
the compounds which hr.ve actually been isolated as crystalline
compounds are listed below:
Enrol ri cal Formula. Resolved Formula
K2[Ga(NH2)5] C-a(l!Ha) 3 . 2KNH 2
Ka [Zn (NHa ) 4] Zn (lIK2 ) 2 . 2KNHa
Ka[Na(NH?)»j Na(rIHa) .2KNH2
K [ Sr (NHa ) 3 J Sr (lIHa ) a . KNHa
. 7
i ...
-118-
The cyanides of iron (ill), silver (i) and mercury (il)
(14,16)
possess amphoteric character in liquid hydrogen cyanide
All of these elements form insoluble cyanides in liquid hy-
drogen cyanide, but in every case these precipitates dissolve
when an excess of the appropriate base analogue is added, e.g.;
FeCl3 + 3[(CaH5)3I-IH]CN~-Fe(CN)3i + 3[ (CaH5)3NH]Cl
Fe(CN)3 v + 3[(C3H5)3I?H]CN - •"■ [(CaHs)3ITK]3CFe (CN)63
Zinc(ll) and copper (il) acetates are amphoteric in ab-
(14)
solute acetic acid . Both elements precipitate as the
acetates when sodium acetate is added to solutions of these
elements. However, when an excess of the base is added the
precipitates dissolve to form compounds TThich are analogous
to the compounds formed in aqueous solutions when an excess
of hydroxide is added to precioitates of the hydroxides.
ZnOlg + 2NaCH3C00^Zn(CH3COO)2 + 2NaCl
Zn(CH3GOO)3 + 2NaCH3C00^Na2[Zn(CH3C00)4]
In absolute nitric acid uranyl nitrate and cadmium (il)
(17)
nitrate have been found to be amphoteric . Addition of
an excess of tetramethylammonium nitrate to solutions of
uranyl nitrate in absolute nitric acid gives the soluble
tetramethylammonium trinitratodioxouranate (VT ) complex.
UOa(N03)a + [(CH3)4N]N03 ^ [ (CH3) 4II] [U0a (N03) 3]
Potassium nitrate, another base in absolute nitric acid, gives
a complex compound lrhen added to solutions of cadmium (il)
nitrate in this solvent.
Cd(U03)a + xKN03^Kx[Cd(N03)2+x]
The actual composition of this compound has not yet been
determined.
I . .■■
-113-
Aluminum (ill ) and tin(lV) sulfites and the oxides of
bismuth (ill ) , gallium (ill ) and antimony (ill ) and (v) have
(14,18)
also been found to be amphoteric in liquid sulfur dioxide solvo ,
'/hen bis-tetramethylammonium sulfite is added to a solution
of aluminum (ill ) chloride in liquid sulfur dioxide, aluminum
(ill) sulfite precipitates. Tnen an excess of the base is
added the precipitate dissolves to give tris-tetramethylammonium
trisulfitoaluminate (ill ) . This compound reacts further TTith
the acid, thionyl chloride, to give the insoluble sulfite
once again.
2A1C13 + 3[(CH3)4N]2S03 ^ AL2(S03)3i. + 6[ (CK3)4N]C1
A12(S03)3^ + 3[(CH3)4N]3S03 ~ 2[(C:-T3)4N]3[A1(S03)3]
2[(CH3)4N]3[Al(S03)3] + 3S0C12 ==^ Al2(S03)3l + 6[(CH3)4N]C1
+ 6S02
Tris— tetramethylammonium trisulfitoaluminate (ill ) is also
formed when a solution of tvro moles of aluminum (ill) chloride
and seven moles of bis-tetramethylammonium sulfite in liquid
sulfur dioxide is titrated with thionyl chloride.
Arsenic (ill) sulfide in liquid hydrogen sulfide beha.ves
analogously to arcsenic(lll ) oxide in water . Arsenic (ill)
sulfide is only slightly soluble in this solvent, but when an
excess of triethylammonium monohydrogen sulfide is added to
the solution the soluble tris-triethylammonium trithio-
arsenate (ill ) complex is formed. The reaction is;
As2S3 + 6[(C2H5)3IIH](HS)^2[(C2H5)3NH]3AsS3 + 3H2S
Aluminum (ill ) and chromium (ill ) are amphoteric in liquid
(14)
hydrogen flouride . Addition of exciss b-^se to solutions of
the fluorides of these elements usually results in the formation
-114-
of the hexaflouro complexes.
III. Summary :
Anphoterism has been found to exist in many non-aqueous
solvents. In almost every case the reactions of amphoteric
compounds in various solvents lead to the formation of complex
compounds. Pfeiffer's theory, describing the process of dis-
solution of amphoteric hydroxides in aqueous solutions, has
been found to be valid in the field of other ionizing, water-
like solvents, whereas those due to Kantzch and Bredig cannot
be successfully applied.
References
1. G-rttttner, Barbara and Jander, C-.: Z.anorg. u. allgem.
Chem. 268, 229 (1952).
2.
3.
Hantzsch, A.: Ibid. 30, 289 (1902).
Pfeiffer, P.: Ber. 40, 4036 (1908).
Scholder, R. and i" at sen, R.: Z. anorg, u. allgem. Chem.
22,0, 411 (1934). •
5. Scholder, R. : Ibid. 220, 209 (1934).
6. Scholder t R. and 7eber, H.: Ibid. 216 T 159 (1933).
7. Scholder, R. and Pat-sen, R. : Ibid. 216, 176 (1933).
8. Scholder, R., Felnenstein, R. and Anel, A.: Ibid. 216r
138 (1953).
9. Scholder, R. and .re,ber, H.: Ibid. 215., 355 (1933).'
10. Scholder, R. and Pats ch, R.: Ibid. 217, 215 (1934).
11. Scholder, R. and Staufenbiel, S. : Ibid. 247, 259 (l94l).
12. Scholder, R. and Kolb, Anneliese: Ibid. 264, 207 (1951).
13. Scholder, R.: Z. Angew. Chem. 49, 255 (1936).
14. Audrieth, L.F. and ICleinberg, J., "Non-Acueous Solvents",
John 'tfiley and Sons, Inc., New York, 1953. pp. 80-1, 139,
159-62, 201, 224-5.
15. Fitzgerald, F.F.:„ -J. An. Chem. Soc. 29, 656 (1907).
16. Jander, G-. and G-ruttner, Barbara; ^er. 81, 114 (1948).
17. Jander, G-. and 'Jendt, Hildegard: Z. anorg. u. allgem.
Chem. 258, 1 (1949).
18. Jander, a. and Immig, K.: Ibid. 233, 295 (1937).
19. Jander, C-. and Schmidt, K.: Wiener Chem. Ztg. 46., 66 (1943)
-115-
Place in the Periodic System of the Heaviest Elements
"rvin Colton March 10, 1953
The heaviest elements refer to elements of atomic numbers 89-08,
Inclusive, i.e., actinium through californium. The suggested electronic
configurations ("beyond radon and xenon) for gaseous atoms of the act-
iniae and Irnthanide series are shown belotr for reference:
At . No, . Element .Conf i~urati on At . !To . El em e n t Co.n figuration
6d17s2
6d-7s^(or 5f 6d 7s )
5f25d17s2(or 5f16d27s2)
5f35d17s2
5f57s2(or 5f46d17s2)
5f67s2(or S^ea^a8)
5f77B2(or 5f33d17s2)
5f76d17s2
5f86d17s2
5f95d17s2
Arguments according to Sea horn; (l,2):.
Present evidence points to the conclusion that it is the 5f
electron shell that is being filled in these heaviest elements.
Further, the evidence seems to suggest a second r^re-earth like
series, beginning with actinium in the same sense that the "Ian-3,
thanide" series begins with lanthanum. Such an "actinide"series
5.s suggested on the basis of the following observations: (A) chemical
properties, (b) absorption spectra in aqueous solution and crystals
89
Ac
90
Th
91
Pa
92
U
95
Np
94
Pu
95
Am.
96
Cm
97
3k
98
Cf
57
La
5d16s2
58
Ce
4f26s2
59
?r
4f36s2
60
lid
A 2
4f'6s~
61
?m
4f56s2
■t-v
St.:
4f66s2
65
Eu
4f?6s2
64
Gd
4f75d16s2
65
Tb
4f96s2
65
^7
4f 6s
- '. ' ...--.„
■■
-116-
(C) cry st allograph! c structure data, (d) magnetic susceptibility and
(E) spectroscopic data.
A. Cher.icrl properties: A table of oxidation states of the lanthanide
and actiniae elements is shown below. Values in parentheses have been
reported but are unstable:
atonic no.
89
90
91
92
93
94
95
96
97
98
element
Ac
TV*
Pa
U
Np
Pu
Am
Cm
Cf
oxid. state
(+2}
(+2)
+3
(+3)
(+3)
+3
+3
+3
+3
+3
+3
+3
+4
+4
+4
+4
+4
+4
+4
+5 +5 +5 +5 +5
+6 +S +6 +6
atomic no.
57
58
59
60
61
62
63
64
65
66
element
La
Ce
Pr
Md
Pm
Sm
Eu
Gd
Tb
Dy
oxid. state
+2
+2
+3
+3
+4
+3
+4
+3
+3
+3
+3
+3
+3
+4
+3
The regularity of the +3 "tate in the lanthanide series is not
so well pronounced in the actinide series. The +4 oxidation state,
as well as the +3 state, seems to characterize the actanides. It
becomes increasingly difficult to effect oxidation to higher valence
states with increasing atomic number among heavier elements. The
following table shows some oxidation potentials of the actinides in
1M acueous solution:
element III to IY IV to VI
U +0.63 v. -0.60 v.
Up -0.14 -0.94
?u -0.95 -1.11
Am v^ -2 . G
The metals of the elements Th to Am bear striking resemblance to
those of the rare earth metals. A marhed similarity is also noted in
7 ?
that americium (5f 7s ) n-nri pu-o-niiim (4f7«o2\ x ^ ^ „ * • «.,
; . no. europium v^ti 6s ) both have densities
much lower than those of their neighbors.
-11/7-
B. Absorption spectra in a.cueous solution n.nd crystals:.
Sharp absorption bands, to a large degree in the visible region,
are characteristic of the lantlianide series - a series irhose proper-
ties are ascribed largely to the 4f electrons. The absorption spectra
of aqueous tripositive actinide elements are strihingly similar to
the spectre of the tripositive lr.nthanide elements; thus lending
support to the concept that 5f electrons are involved in the building
up of the actinide group. It has also been observed that the absorp-
tion spectra become greatly simplified as the mid-lie of each of the
two series is approached. The spectra for gadolinium and curium, each
with seven f electrons, show only one sharp peak between 2000 and
11,000 A0, Trhereas for the other elements of both series the spectra
are more complex. Absorption spectra \rith crystals, especially those
of emericium halides, reveal sharp lines Trith widths comparable to
the sharpest rare earth spectra.
C. Cr^stallorra^hic structure data: Zacharie.sen has observed the iso-
morphism of the compounds Th02, P^02, T.T02, Np02, Pu03 and AmC2, and a
Regular decrc^se in radius of the metallic ion in these oxides. He
has also used x-ray diffraction studies to determine the structure of
•". large number of compounds of Th, U, and the transuranium elements,
fche fluorides ThF4, UF4, NpF4 and ?uF4 are of identical structure
types as are the chlorides UC13, NpCl3, PuCl3 and AmCl3. Calculations
of ionic radii show a progressive decrease in size with increasing
.ntomic number, analogous to the well-hnown lanthanide contraction.
The compounds of the rare-earth elements are in turn isomorphous t-rith
I the corresponding compounds of the actinide elements. The following
[table illustrates these considerations:
-an-
ionic radii of aetinide and lantha.nl de elements
No. of
or 5f <
4f
sins .
Aetinide
III state
series
IV st,
=ite
u
inthani&e series
0
1
Ac+3
(Th+3)
0
1.11 A
(1.08)
Th+4
Pa+4
0.05
0.91
A°
t +3
La
Ce+3
1.04 A°
1.02
2
3
(Pa+3)
U+3
(1.06)
1.04
U+4
0.09
0.08
Pr+3
Nd+3
1.00
0.99
4
r,
6
Np+3
Pu+3
An4"3
1.02
1.01
1.00
Pu+4
Am+4
0.86
0.05
Pm+3
Sm+3
Eu+3
(0.98)
0.97
0.97
D. harnetic susceptibility: One would expect magnetic measurements or
compounds of the heaviest elements to give information on the quantum
states of the responsible electrons. However, the situation is complex
and the exact behavior of the heaviest elements on the basis of either
5f or 6d electrons has not yet been worked out. Yet, qualitatively
speaking, observed paramagnetic behaviors of a number of ions in
various oxidation states are similar to those of the lanthanide group.
The following figure brings out these qualitative similarities:
5000 I-
, n 4000
lolar
suscep-
tibility3000 .
i . g . s .
m\}^ 2000
t 10s
1000 f
Eu(III)
Am(III)
2 3 4 5
No. of f electrons
... „:..-,..-'
■4*-
S. Spectros co'ni.c data: Information In this field is still rather Scan
as fr.r as the heaviest elements are concerned. Investigation of the
spectrum of uranium atoms gave the lowest state of neutral uranium as
of Gd 7s"", a consistent configuration since uranium is the third elemer
j.
in the series. Observations on gaseous Th indicate that the 5f and
6d electrons are very close in the neutral, free thorium.. ^jalitative
comparisons of the emission spectra of americium and europium show a
strong analogy between the intensity of the lines, thus suggesting
strongly that the configuration for gaseous An in the ground state
should be 5f77s2.
Arguments against the "actinide" series (3.4.5):
On the basis of observations that the heavier elements exhibit
both +3 and +4 oxidation states, Zachariasen pr-fers to speak of a
"thoride" series for the tetravalent states and of an "actinide" serie;
for the trivalent states.
Thorium has been shown to be a true homologue of zirconium and
hafnium. The great instability in acueous solution of the triidides
of thorium, zirconium and hafnium is in marine cl contrast to the stabili'
of cerium triiodide . Protactinium is similarly a homologue of niobium
r.nd trntalum. Ease of oxidation of uranium to the +S state brings out
its strong resemblance to tungsten and emphasizes its lach of similar-
ity to neodymium. Uranium is most stable in the +6 state; if it were
to be regarded as an "actinide", similar to the lanthanides, it should
exist primarily in the +3 oxidation state.
Seaborg's arguments leading to the designation of the series as
the actinide series, insofar as they nre based on direct chemical evi-
dence, would involve an extrapolation baclcwrrd from the proposed of
configuration for americium +3 and 5f7 for curium +3. Some of the d: z
Sm
Eu
Gd
Tb
Dy
Pu
An
Cm
Bk
Cf
crepancies between chemical properties and electronic structure come
about because of the slight difference in energy betreen the 5f r>nd Gd
levels. Coryell (5) suggests that these discrepancies may be resolved
by recognition of a delay in the filling of the 5f subshell.
Conclusion: On the basis of the present evidence, both chemical and
electronic, and the opinions of various qualified Trorl:ers, the vie^s of
Coryell seem to be more rational. The heavier elements constitute a
"uranide" series rather than an "actinide" series. The periodic class-
ification, in part, would appear thusly:
La Ce Pr Nd Pm
Hf Ta W
Th Pa U Np
Ac
Note: As a tool in remembering the elements in the tT.rc series, the
following mnemonics may prove useful:
hr nc.er. prince need,s 61 small European £pHn to be dved wholly early
tomorrow. VJhx h? ludicrous?
Actual thoughts pp.ss under Neptune's public amnesty, calming bal^iy
.coffles •
.Bibliography
1. C-. T. Seaborg, Nucleonics. 5 (No. 5), 16 (1949).
2. C-. T. Seaborg, J. J. Katz and T.J. II. Manning, The Transuranium
Elements. National Nuclear Energy Series, vol. IV-14D, p. 1492.
IIcC-raw-Hill Book Co., Inc., New York (1940 ).
3. II. Kaissinsky, J. Chem . Soc . , 1049, S241 .
4. II. Kaissinsky, J. chlm. ohy s . , 47, 415 (1950).
o. C. D. Coryell, Rec. Chem . Prog. , Spring issue, 55 (l95l).
121
Complex Anion Determination by Ion Exchange
II. K. Snyder March 10, 1953
'.Ihile measuring the potential of cadmium amalgam electrodes,
Leden (3; found evidence for the formation of an anionic cadmium
sulfate complex. Fronaeus (2) calculated that the copper in a
(X01 M copper sulfate solution made 0.5 II with respect to sodium
sulfate is present as an anionic complex
per cent. In order to clarify this
these complexes with ion exchangers.
ium
a
i-
to the extent of eighty
situation, Leden (4) studied
For the study, Amberlite IRA 400 was used. In the first
experiments, solutions of cadmium perchlorate, sulfate, chloride,
and iodide were put through a column containing the resin in the
perchlorate, sulfate, chloride or iodide form. Then th^ resin was
washed with 10 ml. of water followed by successive 25 ml. portions
of water. The number of 25 ml. portions necessary to remove all
of the cadmium from the resin was used as an indication for the
formation of anionic complexes. The results pre shown in Table 1.
Table 1
Ex-ot .
Resin was
No".
Saturated
by:
1.
3M NaC104
2.
2M Na2S04
3.
211 Na3S04
4.
3K. NaCl
5.
3M NaCl
6.
311 Nal(?)
Cone, of 10 ml. of Cd
solution influent
C.Oin Cd(C104)2
O.Olh CdS04
0.011 Cd304 and 0.5h Na3S04
C.Glh CdCl3
O.Olh CdCl3 and 0.5H NaCl
No. of
Portions
1
1
1
6
8
O.Olh Cdl2
In experiment six of this series, the cadmium could not be
removed at all. It was also shown that cadmium is entirely removed
from a solution of cadmium iodide when it is shaken with the resin.
Since these data indicate that cadmium sulfate forms an anionic -"com~
pi ex to no greater extent than cadmium perchlorate, it seems likely
that previous indications for the formation of such a species were
erroneous.
Similar experiments ''ere carried out with copper perchlorate,
sulfate, chloride, and acetate,
was done with successive 10 ml.
are n-iven in Table 2.
with the exception that the washing
portions of ir.~ter. The results
122
Table Z
Expt .
Resin was
Cone, of 10 ml.
No.
Saturated
of Cu. solution
by:
influent
1.
5U NaC104
0.01U Cu(C104)2
2.
0.7H Na2S04
0.01N CuS04
3.
0.7II Na2S04
O.Oin CuS04 ana.
0.511 Na2S04
4.
514 NaCl
o.oin cuci2
5.
5H NaCl
0.01II CuCl2 and
0.5U NaCl
6.
3M NaAc
0,0111 CuAc2
7.
3M NaAc
O.OlIi CuAc2 and
0.5H NaAc
% Cu in successive 10 ml.
nortions of effluent
1*2 3 4 5 6 7
5
75
20
15
60
25
1
35
40
20
3
0.3
10
80
10
15
60
25
1
5
65
25
5
1 0.1
0
55
30
10
3 10
is
For the influent in experiments 2, 3, 6, 7, Fronaeus (2) has
calculated that the per cent of co^oer in an anionic conrolex is
0.l£, 80^, 0.5^, 50^ respectively/' Similarly for the influent in
experiment 5, Bjerrum (l; has calculated that 0.02^ of the copper
present as an anion. Apparently there is little formation of an
anionic complex of copper sulfate, at least not as much as Fronaeu
had previously a.ssumed.
Salmon (5) used ion exchange methods to study the complexes
which are formed between ferric ion and orthophosphate . Since both
cationic and anionic complexes have been reported in the literature,
both types of exchangers were tried. However, Permutit Zeo-Karb 225
the cation exchanger, removed neprly all of the iron from solution,
but no phosphate. Either no cationic species were present, or if
they were present, they were very unstable. In further preliminary
studies, unsaturated solutions of iron in Phosphoric acid were put
through columns with the Zeo-Karb and then through the chloride form
of the IRA 400. The process was then reversed. Results are
summarized in Table. 3. These data seem to indicate an easily dis-
placed equilibrium between ferric ion or a cntionic complex and an
anionic complex.
SOLUTI ON
Table 3
COLUMNS
% Fe in
HP205 NFe203 Ml. Used 1st
2nd
IRA
Iffluent
0.16
0.001
25.00
ZKH
100
0.16
0.001
50.00
ZKH
97
1.13
0.066
5.85
ZKH
ZKH
100
0.7
0.015
10.20
ZKH
IRAC1
98
2
nil
1.8
0.067
5.10
IRAC1
96.5
3.5
nil
1.5
0.046
5.00
IRAC1
ZKH
82
18
nil
1.13
0.056
5.01
IRAC1
ZKH
79.5
18.4
2
1.8
0.067
5.00
IRAC1
ZKH
02
17
1
1.13
1.13
1.13
0.065
0.065
0.066
6.00
6.65
:-5.00
IRAP04
IRAPOj
IRAPC4
ZKH
I RAP 04
IRAPO*
30
70
99
100
1
123
Further anion exchange was carried out Trith the phosphate form
of IRA 400 which was left in contact with saturated ferric phosphate
for 1-8 weeks. After filtering, the resin was washed in a column
with water and then eluted Trith 2ii KC1. The eluate was analyzed for
iron and ■oho si-hate. Table 4 shows the results.
SOLUTION
Tab
le 4
ADSORBED OH
RE
SIN
;1?205
^Fe303
P305
111.
mg.P20s
mg. Fe
3o3
P205
Fe203
Used
Fe26,
31.3
4.77
7.37
5.00
548
181
3.4
29.1
4.95
6.60
7.25
708
238
3.4
26.6
5.31
5.65
6.70
645
260
2.8
18.0
2.25
9.01
6.00
563
179
3.5
9.08
G.59
17.3
6.01
575
143
4.5
6.07
0.11
60.5
6.00
598
76
8.8
The minimum value approached by the ?205/Fe203 ratio at
approximately 25^ P205 was interpreted as an indication that the
anionic complex under these conditions contains three phosphate
groups for each iron atom. The increase beyond 25^ P205 may be due
to one of two things, viz., the increase in concentration of H3P04
results in increasing competition of the phosphate for the resin or
complexes with increasing numbers of phosphate groups may be formed.
BIBLIOGRAPHY
1. Bjerrum, J., Kgl . Danshe Videnshab. Selshab. llat.-fys. Medd.
22, No. 18, 43pp. (1946).
2. Fronaeus, S., Acta Chem. Scand. 4, 72(l950).
3. Leden, I., Acta Chem. Scand. 6, 97(1952).
4. Leden, I., Svensh. Kern. Tidskr. M, 145(1952).
5. Salmon, J. E., J. Chem. Soc. 1052, 2316.
124
TH~ NATURE OF THE FRIEDEli* CRAFTS COIIPLEX
Clayton T. Elston March 17, 1953
Since the original publications of Friedel and Crafts (l) the
literature dealing with the reaction has grown to quite extensive
proportions and, fcas been the subject of numerous reviews (2).
Early investigators assumed that the reaction involved an inter-
mediate org? no-aluminum compound, followed by coupling of this
material Trith the alkyl halide. This and related theories w<=re
later shown to be untenable. Further studies also revealed that
many other compounds are capable of catalyzing the reaction between
all:yl halide s and aromatic compounds. Examples are: FeCl3, SnCl4,
TiCl4, BiCl3, ZnCl2, BF3, and H2S04. In general, any compound which
has strong acid character (Lewis Definition) exhibits catalytic
activity. In an attempt to explain this activity the binary systems
aluminum halide-alkyl halide and aluminum halide-aromatic hydro-
carbon and the ternary system aluminum halide-al!:yl halide-aromatic
hydr oca. rbon , wi 11 be di scu s se d .
Carefully purified ethyl bromide is practically nonconducting'
(k= 3 x 10"9mhos). On the addition of AlDr3 the conductivity
increases with increasing concentration to about k= 1.21 x lCT^mhos
for a 20;' solution of A13r3. Flotinkoff (o) was the first to study
the electrolysis of solutions of AlBr3 in ethyl bromide. Using
aluminum electodes he found that metallic aluminum deposited on the
cathode. Similar results "frerP obtained by other investigators.
'.rertyporoch (5) studied the same system using Platinum electrodes
but his results vPre inconclusive. He found that the aluminum
concentrations irt the neighborhood of the cathode and anode were
almost ecual. He also reported that a noticeable separation of
aluminum occurred at the anode. A recent study on this system has
given somewhat different results (g). It was found that the
principal electrode reactions r.re the deposition of metallic a.lumini
at the cathode and liberation of bromine at the anode. The aluminun
concentration in the anode corn'orrtment remains almost constant; a.
reduction in the aluminum concentration occurs in the cathode com-
partment. On the basis of these data it would appear that the
aluminum is ore sent in solution in both the anionic and cation! c
form. The simplest explanation would involve equilibria such as t
following.
A13r3 < > [Al?.r2]+ + Br"
2AlBr3 < » [AlBr2]+ -+- [AlBr4]"
R-Br + AlBr3 ^=* R Br«AlBr3 <£Z> R+ [AlBr4]*"
Vapor pressure studies on the system metal halide-alkyl halide
have given evidence for complex formation of the type.
CH3C1 + GaCl3 r ■> CH3ClCraCl3
he
125
Brown (7) reports the following vapor pressure-composition diagram
for the system G-aCl3-CK3Cl. Van Dyke (G) obt"inedno evidence for
26.0
cm.
3, a
■ «-
Temp. -78.5°C.
. » m » I
t^ — ~$T
TT
~. » .mm-- mm
r~ — $r — <r
oles CH5Cl/0aCls
such complex formation with the
pre inconclusive .
■ystem AlT*,r3-G2H5Br but his results
Positive evidence for the equilibrium, H3CC1 + A1C13
R3CC1»A1C13 is furnished by the rapid racemization of optically
active alkyl halides in the presence of A1C13. Similarly when
A1C13 containing labelled chlorine was used as a catalyst for the
reaction of benzene with t-butyl chloride it was found that complete
interchange of chlorine atoms had occurred (0).
Although several workers (6). (9) have reported the formation of
complexes between aluminum halides and aromatic hydrocarbons the
existence of stable complexes is still questionable. Such complexes,
if formed, ^ould appear to involve only weak attractive forces.
Their role in the Frie del- Crafts reaction is uncertain but is
probably a very minor one.
Benzene has a very small specific conductance (k= 1 x 10 mhos
at 25°C.)-and ur>on addition of aluminum halide there is no notice-
able increase in conductivity. However, addition of alkyl halide
or halogen acid to such a. solution produces a very marked change
(101). The solution becores colored and as the concentration of the
alkyl halide (or halogen acid) is increased a second liquid •ohase
separates. The lower phase is highly colored and strongly conduct-
ing (k= 1 x 10 ^nihos) while the upper phase is only slightly colored
and weakly conducting (k= l.Q x 10~5mhos). The reactions involved
were found to be reversible since removal of the halogen acid yield-
ed a homoger&us system iThich could be further separated into pure
aluminum halide and hydrocarbon.
aluminum chloride-toluene-hydroChloric-
t A1C13 dissolves in toluene in the
rilliant green solution. Relating
pressure of HC1 -^oove the solution at
limiting case a/v^roximately one mole of
le of A1C13 which goes into solution,
s taken up for every two moles of A1013.
el-Crafts complexes are organic salts
AICI4 and HA13C17. The high solubility
In studying the system
acid, Brown (ll), found tha
presence of HOI to give a b
solubility of A1C13 to the
"80°c. they found that as a.
HC1 is taken up for each mo
At -45° C. one mole of MCI i
They suggest that the Fried
of the hypothetical acids H
Ar + HC1 + AICI3
of A1C13 in such complexes would tend to indicate that complexes of
* Arl-T A1C14
-.V —
126
a higher order are also possible. The general formula would be
ArH C^n^on+1^""' It: should be noted that results of various work-
ers (7) (8) have shown that there is no evidence for an acid of the
type HAL CI 4 or HAlBr4. Studies of the system A1X3-HX over a *'ide
range of temperatures (-120° to 300°C.) revealed that the pressure
of HX is not affected by the presence of ALX3. It would thus be
extremely improbable if a detectable concentration of the free acid
could exist under the conditions of the Friedel-Crafts reaction.
The exchange between labelled A1C13 and HC1 at low temperatures
offers a means of estimating the equilibrium concentration of HAIGI4
and such an investigation is now being conducted by Professor Sensor
AICI3 + HOI t * AICI3 * HC1*
at the University of Southern California.
The electrolysis of the ternary complex aluminum bromide-
ethyl bromide-benzene using an aluminum anode has been shoT,rn to be
an efficient method for aluminum plating (12). The overall elec-
trode reaction is the solution of aluminum at the anode and the
deposition of aluminum at the cathode. It seems reasonable to
assume that AlBr3 could behave as a l:-3 electrolyte when dissolved
in the highly polar complex phase. T/ertyporoch (5) studied elec-
trical transference in the system aluminum bromide-ethyl bromide-
hexaethylbenzene and found that aluminum concentrates in the anode
compartment, while bromine and hexaethylbenzene concentrate in the
cathode compartment. Such data substantiate the formulas proposed
by Brown.
Bro-i-Ti (ll) assumes that the formation of the ternary complex
is involved in the rate-controlling step of the Friedel-Crafts
reaction. The highly polar complex phase, which is capable of
dissolving an excess of any of its three components would presum-
ably furnish an excellent medium for the ionic reactions involved.
BIBLIOGRAPHY
(1) C. Friedel and J. li. Crafts, Comot . rend., 100, 692 (1885) .
(2) D. V. Nightingale, Chem . Revs., 25, 329 (1C39).
E. Berliner, Org. Reactions, Vol. V, page 229.
(3) '.;. A. Plotnikoff, J. Russ. Phys . Chem. Soc . , 3, 466 (1902).
(4 J H. E. Patten, Trans. Electrochem. Soc, 6, 9 (1904).
(5) S. vJertyporoch and A. TJohl, Ber., M, 135? (l93l).
(S) R. E. Van Dyke, J. Am. Chem. Soc, 72, 3619 (1950).
(?) H. C. Brown, H. Pear sail and L. P. Eddy, J. Am. Chem. Soc,
72, 5347 (1950).
(C) F. Fairbrother, Trans. Faraday Soc. 37, 763 (l94l).
(9) V. A. Plotnihov and N. II. Gratsianskii, Bull. acad. sci . ,
U.S.S.R., Classe sci. chim . , 101 (1947) .
C.A. , 4£, 4480 (1948).
(10) D. D. Eley and P. J. Xing, J. Chem. Soc, 1952, 2517.
(11) H. C. Brown and. H. IT. Pearsall, J. Am. Chem. Soc, 24, 191
(1952).
(12) R. D. Blue rnd F. C. Ilathews, Trans. Electrochem. Soc,
69, 519 (1936).
127
ACID-EASE BEHAVIOR IN INERT SOLVENTS
Harch 24, 1053 Daryle H. Busch
Introduction. At the turn of the century, the ionization
theory of acids and bases was commonly believed to apply to all
solvent systems. Substances such bs hydrogen chloride were not
considered to be acids in solvents which produced no ions. In
1902, Kohlenberg (31 ) reported several ionic tyne reactions in
benzene. Shortly thereafter, Vorlander (43) carried out the
titration of aniline with hydrogen chloride in benzene using methyl
yellow as an indicator. These and similar developments led a
number of investigators to suspect that ionization is not a necess-
ary condition for acid character.
Hethods Employed in the Study of Aci dr and Eases in A">rotic
- "edla.
Hantzsch was of the opinion that a more broadly valid criterion
for acid strength than "hydrogen ion concentration" could be found
in the relative tendencies of acids to form salts with indicator
bases. A technique was developed for describing the relative
strengths of acids in terms of the stabilities of such salts. (El)
(22) (29) Kantzsch and his co-worliers «lso measured the catalytic
effects of different acids on the rate of inversion of succrose and
the rate of decomposition of diazoacetic ester and related these
rates to the intrinsic strengths of the acids. (21) (22) (26) They
found that the acids "ere com:, only stronger in this respect in inert
solvents such as benzene raid chloroform than in basic media like
water.
The investigations carried out by Laller a.nd Do^nes (36) (3?) (38)
are especially significant in demonstrating the use of indica.tor
methods in determining the relative strengthr of acids and bases in
aprotic media. In the ideal cape, HA is considered to react Trith a
base B to produce a. much wealier acid KB in the presence of an indie a"
or I. The indicator is partially converted to its acid form HI.
From a. consideration cf the acidity constants for the acids HA, HE,
and HI the following expression is derived:
log CA]/[HA] = log Cl]/[HI] - P%A+ ?KKI
A plot of log [A]/[HA] against log [l]/[Hl] then gives a family of
parallel straight lines TTith unit slopes •"ho re intercepts on the
log [I] /[HI] axis give a measure of the strengths of the acids as
compared to the indicator HI .
An extensive investigation of the reaction of various organic
bases T"ith indicator acids in inert media has been undertaken by
Davis and her associates. The technique employed is best character-
ized as a spectro-pho tome trie titration of an indicator acid with a
basic material. The relative strengths of a series of bases have
been determined a.nd association constants corresponding to the
reaction shown below were measured.
B + HA 'z-z, BH+ A"
Two indicators were synthesized during the ea.rly phases of these
studies (o); bromphthalein magenta E (tetrabrom.ophenolphtha.lein
128
ethyl ester) and bromphthalein magenta B (tetrabromophenolpth^lein
n-butyl ester). These indicators are mor° soluble in aprotic media
and often give simpler color changes (10) than the raorp familiar
indicators.
A parameter representing the true strengths of bases in combin-
ation T~ith a standard indicator acid is found in the association
constant, Ka, for the reaction mentioned above.
Ka = [EH A]
[B][HA]
From this expression the eauation belo*r is obtained.
log [BHA]/[HA] - log [3] = log Kft
The association constant may be evaluated by plotting log [BKA]/[HA]
against - log [B]. In order for Ka to be valid, the plot should be
a straight line with a negative slope equal to unity. The validity
of the results may be judged from Table 1 below. (9;
Organic Base hog K-, Ka Slope of Curve
di-n-butylamine 4.19 1.5xl04 -1.17
triethylamine 4.36 S.SxlO4 -1.03
piperidine 5.08 1.2xi05 -1.08
dipheny Iguani di ne 3 .35 2 . 2x1 0^ -1 . 03
ditolylguanidine 5.80 6.3xl05 -1.23
Another type of information which has been obtained by Davis
et al stems from the nature of the color changes which the indicator
undergoes upon addition of a base. These color changes provide some
insight into the nature of the chemical reactions which are taking
place. All of the changes observed may be explained on the basis of
three assumptions:
1. The yellow color of solutions of bromphthalein magents (BPIl) in
aprotic media is characteristic of the neutral, unsolvated, and non-
ionized acid form of the indicator. 2. The blue color of solutions
of tetraalkylammonium salts of 3PII in aprotic solvents is character-
istic of the BPM anion ^Tlien it is associated Trith the positive con-
stituent of the salt only through coulombic attraction. 3. The
magenta color of solutions of BP1I containing en excess of a tertiary
amine is characteristic of a highly polar addition compound formed
by incomplete removal of the proton from the indicator anion ( a
hydrogen bridge exists between the amine and the anion).
The shift in color from that represented by the primary addition
compound formed by a primary, secondary, or tertiary amine with the
acid indicator to the color characteristic of the coulombically
bonded tetraalkylammonium salt in basic solvents or upon the additior
of a small amount of r basic solvent to a solution in an inert med-
ium may be explained similarly with the generalized mechanism.
129
B + HA ±-:BH+..,A~ + B< fe(HH+...Bt) A~
It also follows that the stability of the color characteristic of
the primary reaction, when an indicator acid is dissolved in a basic
solvent of very large steric requirements, stems from the reaction
of one molecule of b,°se to form a highly polar addition compound
whose hydrogen bridge is so shielded by the large solvent molecule "
that the secondary reaction, breaking of the hydrogen bond between
the anion of the acid .and the cation by the rction.of the second
solvent molecule, cannot be accomplished. This sort of phenomenon
finds support in the studies of Brown and his co-workers who showed
that when the steric requirements of an acid and base pair are too
1 rge no reaction can occur, (o)(lG)
Theoretical Slrnif ic-nce of Studies in Aorotic Solvents . Al-
though it is probable that no solvent of practical value is complete-
ly devoid of acidic or basic character, solv-nts such as hydrocarbon'
halogenated and partially halogenated hydrocarbons may be considered
to have neither acidic nor basic properties since their relative ten-
dencies to behave either as acids or as bases are much smaller than
is the case with the substances which are under study.
The factor which would be most significant in its effect on the
measurable properties of acids and bases and their salts in aprotic
media is the dielectric constant of the solvent. The effect of
dielectric constant on the behavior of acids and bases and their
salts has been studied by Bronsted, (7) by Hamnett, (20) and by
Kraus and Fuoss.(32) ICraus (34) reported that the properties of an
acid in an inert solvent nr° largely dependent on the quantum forces
between the proton and the anion. The role played by the dielectric
properties of the solvent was found to be relatively slight. Other
investigators (l)(2)(l6) report that a solvent of higher dielectric
constant than is known would be required to dissociate an acid with-
out the formation of a neT- chemical bond. ICraus (34) and Hammett
(20) have shown that the degree of dissociation of a salt depends
on the dielectric constant of the medium . In benzene, ion pairs
would be expected to form. (4) (33) The interionic attraction theory
relates the dielectric constant of the medium to activity coeffic-
ients. (5)
The, hechani sm of Acid- Base Be actions.. Bronsted (7) envisions
the combination of t,ro hypothetical half reactions to produce neutral-
ization. No mechanism for the exchange of the proton is offered.
The Lewis theory arrives at the same final result by presuming
that the reaction involves contact of the Base A with the acid HB.
T.ie intermediate compound is supposed to have only a transient exist-
ence since it is considered unrealistic to suppose that a proton can
share tiro pairs of electrons simultaneously, (39)
Demonstration of the formation of the hydrogen-bonded addition
compound as the first product of the reaction of a hydrogen acid
irith an amine provides a link between the Le^ris and Bronsted theories
In fact, this scheme presents a unique cnse of neutralization in the
Lewis concept only because of the high polarity of the compound form-
ed. The evidence cited in support of the existence of this addition
130
compound is: (13) (a) Ionization is unimportrnt in the reaction;
(b) It is possible to calculrte an equilibrium constant for the
reaction assumed; (c) Spectropho tome trie evidence mahes it poss-
ible to identify the species with the conditions under which it
p::ists; (d) The "salts" formed in these reactions have dipoles
quite similar to those found for the more obvious products of Lewis
acid-base reactions.
The two-stage mechanism proposed by Davis and co-workers is
represented by the scheme below. (13)
r> :t 4. tjt .:._,. rj t.t-jt T~
113H ~ nj. -— ^ .'.31;.. • ♦ »J-
yellow magenta
Rs-h"1"...!" + r3n zz: (R3:iH+...iin3)i~
magenta blue
The first equation represents the formation of a hydrogen bond. The
second equation represents rupture of the old hydrogen bond, which
joined equal and opposite charges, and the formation of a new hydro-
gen bond with the incidental effect that ion pairs are produced.
From less thorough investigations, the effect of acid molecules
on the process of neutralization may be inferred. I-Iaryott (40 )
-ted system indicates complex
formation between carboxylic acids and their salts in inert solvents,
it was proposed that the anion, RCOO-, of the salt forms a dimer
with a molecule of acid. The species formed would b° somewhat
analogous to the ordinary dimeric structure of the free acid in an
inert solvent.
Inasmuch as no solvent has been found TThich possesses acidic
but not basic properties and since autoprotolysi s of such strong
acid solvents as sulfuric acid and hydrofluoric acid has been
demonstrated (3) the scheme appears to be quite reasonable. This
leads to the conclusion that the dissolution of a base in an acidic
solvent follows a two step equilibrium.
The Relative Strengths of Acids md. Danes. The recognition
that -viter and other basic solvents exert r leveling effect (18) (25)
on the strengths of acids in solution has been instrumental in caus-
ing various investigators to study acid strengths in aprotic solvents
The leveling effect may be demonstrated by the addition of water to
solutions of various acids (29) in aprotic media. Such an experiment
would reveal that the relative acidities of the we alter acids are
increased while those of the stronger acids nre decreased. In gener-
the restriction posed by the leveling effect is that no acid may
exist in a basic solvent whose proton donating tendency exceeds that
of the onium ion of the solvent. It is obvious in view of this
limitation that the strengths of the stronger acids must be measured
in solvents of very slight basic character.
131
The most common method of measurement of the strengths of acids
in aqueous solutions is Vr determining the electromotive force Trith
an electrode reversible to hydrogen. The electrode potential as a
criterion for acidity is independent of the interpretation given to
it by the Bronsted theory. It is the activity of hydrogen which is
me^sur^d in this manner, end this activity is not necessarily
parallel with the concentration of hydrogen or hydro nium ions. It
is, to the contrary, a measure of the reversible work required to
transfer a proton from one riven base to another. In the ideal
case, (5) (20) (33) the hydrogen electrode potential would provide
the me-^ns for extending the precise measurement of acid and base
strengths to all solvents. This is not feasible at the present be-
cause of such complicating factors as licuid junction potentials
(22) (23) and individual ion activities. (24)
In attempting to measure the relative strengths of acids in
aprotic media, methods based on colorimetric or spectrophotometries
techniques have proven most fruitful . (s) (l?) (21 ) (20 ) (38) The form-
ulation most often used is essentially that given by Lalier and
Downes; this analysis follows the Bronsted theory. Table 2 lists
series of acids in the order of their decreasing strengths as ob-
served by several investigators. The velues of pK are given for
tT.ro of the series. These parameters differ in their numerical valuer
primarily because different arbitrary standards were chosen by the
respective investigators. The general trends are the same.
Lalier and Downes (37) (38) have demonstrated an ingenious tech-
nique for setting up a series of relative acid strengths based on
indicator studies in inert solvents. This method also involves the
Bronsted scheme for the acid-base reaction. These investigators
found that their indicator method was limited, for any given indicat-
or, to those acids whose pll values differed by plus or minus one
unit, or less, from the pIC value of the indicator. They then pointer
out that, if the relative acidities of tTro indicator systems could
be established, the acid systems studied with these t7To indicators
could then be compared. For example, dichloroacetic acid is weaker
than dimethyl yellow but stronger than bromphenol blue. The
experimental curves for these two indicrtors can then be displaced
along the log [l]/[Hl] axis of a graph of log [A]/[HA] against
log [l]/[Hl] until they become continuous. By a stepwise correlatior
of indicators in this manner it is possible to erteblish a complete
numerical scele for all the acids.
A number of investigators (4) (14) (15) (27) (22) (30) (35) have
concluded that the primary reaction which tahes place between a base
and a proton acid culminates in the formation of r highly polar add-
ition compound. The extent to which this reaction proceeds may be
measured and an equilibrium constant may be calculated. Table 3
lists the systems which have been studied and qives the association
constants for this primarv reaction.
•
132
Investigator
Table 2
Relative Acid Strengths
Bronsted(G) Hall (19) Hantzsch(2l) (22)
Kantzsch(2C
Method
Indicator
pK&
Indicator
Inversion
of Sugar
Solvent
Benzene
'later
Chloroform
'later
hydrochloric
methyl red
dimethyl yellow
(ion)
trichloroacetic
dichloroacetic
picric
o-ni t robe n zoic
chloroacetic
salicylic
bromphenol blue
b-di ni t rophe nol
o-chlorobenzoic
neutral red (ion)
m-*chlorobenzoic
bromcresol green
benzylammonium
ion
formic
phenylacetic
benzoic
acetic
i soamylammonium
ion
bromcresol purple
piperidinium ion
brom thymol blue
-7.4
•oer chloric
hydroiodic
4.9
sulfonic acids
perchloric
3.5
hydrobromic
bydrobromic
0.7
hydrochloric
benzene-
sulfonic
1.3
nitric
hydro chlori
0.3
trichloroacetic
nitric
2.3
tribromoaceti c
trichloro-
acetic
2.9
maleic
sulfuric
3.0
malonic
chloroaceti
4.1
chloroacetic
formic
3.7
a-broraonropionic
acetic
bromoacetic
G.9
formic
3.8
b-iodo^roiiionic
4.7
acetic
9.4
3.7
4.3
4.2
4.7
10.6
6.3
11.1
7.0
Investirator
Laller and Doirnes (38)
Griffiths (17)
"ethod
Indicator
Indicator
Solvent
Benzene
Chlorobenzene
MIA
methyl red
trichloroacetic
propyl red
dimethyl yellow
dichloroacetic
salicylic
chloroacetic
bromphenol blue
bromcresol green
formic
benzoic
aSe^e1 red
diethyl ammonium ion
pIC Act c\
2&
0.55
bromphenol blue
3.46
0.55
trichloroacetic
3.37
0.7
dichloroacetic
2.52
0.8
chloroacetic
0.77
1.7
salicylic
0.74
2.6
di nit rophe nol
0.00
2 • 8
ben-oic
-0.58
3.2
acetic
-1.00
3.4
propionic
-1.08
3.7
4.5
5.5
133
Table 3
Association Constants, K , and Dissociation Constants, K^,
for Primary Acid-Base Reactions in Inert Solvents.
_?_££ Acid Solvent lCa K-a.
•.ribenzylamine(ll) trinitro-m-cresol bPnzene 450 —
riphenylguanidine (13) bromphthalein magenta E benzene 525 —
rlbenzylamine(ll) picric acid benzene 1,600 —
.I-n-butylamine(9; bromphthalein magenta E benzene 15,500 —
;riethylamine(9) bromphthalein magenta S benzene 23,000 —
-.riethylamine (9 J bromphthalein magenta B benzene 23,000 —
iperidine(c) bromphthalein magenta E benzene 120,000 —
Liphenylguanidine(l3) bromphthalein mare nt a E benzene 220,000 —
lirhenyl;~uanidine(9) bromphthalein magenta B benzene 220,000 —
ll-o-tolylguanidlne (9"j bromphthalein magenta E benzene 800,000 —
Li-o-tolylguanidine (9 ) bromphthalein magenta B benzene 860,000 —
:,N-dimethvlanillne (12) hydrogen chloride CHC13 67,000,000 —
'.niline(35) oicric acid C6E5IT02 125 2xl0~5
Limethylanillne(35) picric acid C6H5N02 4,000 4.1x10"
uridine (35) picric acid C6H5!T02 51,300 5.54x10*
irimethylamine (41) sulfur dioxide benzene
or OH 01 3 525 —
Equilibrium Constants, K2, for "the Secondary Reaction of
Acids and Bases
5ai?e Acid Solvent K-
iphenylguanidine (14) bromphthalein magenta benzene 15.5
riphenylguanidine (14) bromphthalein magenta benzene 22.5
^A second measurable equilibrium has been demonstrated by Davis
and Hetzer. (13) (14) This equilibrium measures the extent to ™hich
the addition compound 3H+...A" is solvated and dissociated into ion
pairs by one mole of base B1 , which may be the same as or different
from the base B. The values of the equilibrium constants re-ported
for the tT-o secondary reactions studied by Davis and Hetzer are
listed in Table 3.
134
Literature Cited.
1. Arthur and 'Jest: J. Chem . Phys., 2, 215 (1934).
2. Arthur and Vest: J. Chem. Phys., 5, 10 (1937).
3. Bell: "Acids and Bases," John Piley and Sons, Hew York, 1952.
4. Bjerrum: Kgl. Danske Vedenskl. Selskad, 7, C9 (1926).
5. Bronsted: Rec. trav. chem., 4j>, 718 (l92o),
6. Bronsted: Ber., £1, 2049 (1928).
7. Bronsted: Chem. Revs . , 5, 23 (1928).
8. Davis and Elderfield: J. Am. Chem. Soc, J5Jk, 1499 (1932).
9. Da.vis and Schuhmann: J. Research, P.B.S., 39, 221 (1047).
10. Davis. Schuhmann, and Lovelace: J. Research., U.B.S., 4JL, 27
(1948).
11. Davis and PcDonald: J. Research, N.B.S., 4°, 595 (1949).
12. Davis: J. Am. Chem. Coc , 71, 3544 (194977
13. Davis and Hetzer: J. Research, K.3.S., 16, 49S (l95l).
14. Davis and Hetzer: J. Research, N.B.3., 4J1, 331 (1952).
15. C-ordy: J. Chem. Phys., 7, 93 (1939).
16. C-ordy and hart in: J. Chem. Phys., 7, 99 (1939).
17. Griffiths: J. Chem. Soc, 1938, 818.
18. Hall and Conant: J. Am. Chem. Soc, 49, 3047 (1927).
19. Hall: Chem. Revs., 8, 191 (1931).
20. Hammett: J. Am. Chem. Soc, 50, 2666 (1928).
21. Hantzsch: Z. Elektrochen. , 29, 221 (1025).
22. Hantzsch: Z. EleP.trochem . , 30, 194 (1924 ).
23. Hantzsch: Z. Elektrochem. , 31, 167 (1925).
24. Hantzsch and Bucerius: Per., 59, 783 (1926).
25. Hantzsch: Ber., 60, 1933 (1927).
26. Hantzsch: Z. Physil: Chem., 125, 251 (1927).
27. Hantzsch. and Duriger: Z. Physik Chem. 134 400 (1928).
28. Hantzsch: Z. Physik Chem., 134, 406 (1920).
29. Hantzsch and Voigt: Per., 62, 975 (l929).
50. Izmailov: Zhur. Fiz. IChim., 24, 381 (i960); c£, Chem.
Abstracts, M, 6706 (1950) .
51. Kahlenberg: J. Phys. Chem., 6, 1 (1902).
32. Kraus and Fuoss: J. Am. Chem. Soc, 55, 1019 (1933).
33. Hraus: Trans. Electro chem. Soc, 66, 179 (1934).
34. Kraus: J. Phys. Chem., 43., 231 (1939 ).
35. Kraus and Pitschonke: J. Am. Chem. Soc, 69., 2472 (1947 ).
55. Lailer and Downes: J. Am. Chem. Soc, 53, 888 (l93l).
57. LaPer and Downes: J. Am. Chem. Soc, .55, 1840 (1933).
58. LaPer and Downes: Chem. Revs., 13, 47 (1955 ).
39. Luder and Zuffanti: "The Electronic Theory of Acids and Bases",
John Piley and Sons, Inc., New York, 1946.
40 . : laryott : J . Re sea rch , P . P . S . , 38 , 527 (1947 ) .
41. P'oede and Curra.n: J. Am. Chem. Soc, 71_, 852 (1949) .
42. Ralston, Pool, and Parrood: J. Am. Chem. Soc, 67, 775 (1945).
43. Vorlander: Ber., 33, 1435 (1903).
lob
METAL- DIAMINE CHELATE COMPOUNDS CONTAINING
HIGHSK-i:SK3ERSD PINGS
William E. Cooley Ilarch 31, 1953
I . Introduction
It is well known that nitrogen atoms in amines readily act as
electron pair donors in coordinating with certain metal ions. Many
stable complex compounds contain coordinated ammonia or amine mole-
cules. Even greater stability results from coordination of ethyl-
enediamine or its homologs, since formation of a chelate ring takes
place. Chelate rings having five or six members appear to be much
more stable than larger rings, with respect to their resistance to
hydrolysis in water solution and decomposition by heat.
II . Five-membered Rings
Ethylenediamine (en) and substituted ethylene diamine s form five-
membered chelate rings with metal ions.
.I!K2-~ CH3
MZ' 1 M= Co, Cr, Pt, Ni, Cu, Fe , etc.
'\rH2..._ 0H2
Substituted groups on the carbon atoms of ethylenediamine usually
have little effect on the c^lor, stability, or method of preparation
of these complexes. Solubility in water is sometimes increased by
the addition of such groups; the propylene diamine (pn) and 2,
3-butylenediamine (bn) complexes of cobalt are more" soluble than their
ethylenediamine homologs. (l) Cyolopentane&i amine (l), cyclohexane-
diamine (o), i_so-butylenedia~ine (ibn) (9), and me_so-stilbenedi amine
(l) have also been used as chelating ligands .
III. Six-membered Rings
*
Increasing to three the number of carbon atoms between the amine
-roups lowers the stability of diamine chelates. I'feiffer and
Haimann (10) were unable to prepare chromium complexes with tri-
ne thylene diamine (tn) by the same reaction which" they found to give
good yields of ethylenediamine end propylenediamine chelates. Bailar,
Hollinson, and Work (l, 15) likewise found that anhydrous ethylene-
diamine and anhydrous chromium (ill) sulf .' react to form [Cr en3]3 -
(S04)3, but that anhydrous trimethylenediamine does not undergo a
parallel reaction.
i.
iann and Pope (?) resolved into its optical antipodes the chelate
compound [C0.4 Pt NKa CH2 CH (NH2) CHa I!H2], Resolution of this
compound is possible only if it has a 5-ring structure.
i.-x2 U-2
Q.4 Pt I.
^■NH2 - CH
NH2 - CH2
Coordination in the 1 and 3 positions would yield a symmetrical 6-ring,
136
It is possible, however, to produce six-m eir.be red chelate rings.
Jschugaeff (17) reported that addition of trimethylenediamine to
nickel disuccinlmide 8-hydrate resulted in the formation of a violet,
water-soluble complex compound, [lTitn3l (succlnimide ) 2 * 2H20. Also
prepared by Tschugaeff (19) were the blue-violet crystalline compound
INi tn2]SC4 and the unstable red-violet CKitn3]S04- The chloroplat-
inite salt of [Nitn33"H" is more stable r,na nay be formed from nickel
(il) chloride in water solution, t rime thylene diamine, and potassium
chloroplatinite .
T'erner (21 ) prepared bjjsy-trimethylenediamine complexes of cobalt
(ill) by the same methods used for bis,-e thylene diamine complexes.
Tschugaeff (19) and Drew end Tress (4) found that addition of tri-
methylenediamine to a wa.ter solution of potassium chloroplatinitp at
room temperature produces yellow [Pttn CT2]. At higher temperatures
[?ttn2][Pt GLj is' formed.
Breuil(2) added trimethylenediamine to ferrous chloride, ferrous
bromide, and ferrous iodide in anhydrous methyl alcohol. T.rith the
work carried out in a hydrogen atmosphere, the corresponding tri s~
diamine complexes, [Fetn3]X2, were formed. These compounds decompose
in air.
Bailar and ViTork(l) have prepared trimethylenediamihe complexes of
cobalt, noting that [Cotri3~]~7"r~r may be prepared only in the presence
of a catalyst, such as activated charcoal. Oth°rTrise carbon dioxide
from the air forms carbonate ions, which coordinate more readily than
the diamine, yielding [Cctn2 C03] .
Chelates containing six-membered rin^s have also been prepared
with 2, 2-dimethyltrimethylenedi amine (l) , L-methyltrimethylenediamine
(6), and 2, 4-di ami noibe nt ane . (3,19)
IV. Rings of "lore Than Six Members.
Tschugaeff (17,18), Werner{2l), Drew and Tress (4) and Pfeiffer
and Kaimann(lO) attempted to form chelates containing tetramethylene-
diamine and pentamethylenediamine. Their efforts resulted in the
formation of insoluble residues, products without fixed compositions,
or metal hydroxides. In some instances no reactions at all took place
Pfeiffer and lubbe (12) also were unable to form chelates of seven and
eight-membered rings. These investigators studied de came thylene -
diamine and octade came thylenedi amine as well, but no chelates were
obtained. HcReynolds (8) found that coordination takes place between
de came thylenedi amine and cupric ions, but there is no indication of
the presence of chelates.
Pfeiffer (ll) succeeded in forming such complexes as
> V
x=4, 5, 10
in which a large ring is stabilised by the presence of two six-membere.
rings.
137
Pfeiffer, Schmitz, ■"nd Bohr. ( 14) have recently reported succeci
stabilizing diamine chelate rings of more than six members. Using
alcohol and ether solutions as reaction media, these workers have
isolated crystalline compounds whose compo notions correspond to
in
-- «
(GH2)n i ! X
M = Cu
++
++
m
/ m
n
x
,,2n"Hg++lCa
(m=2) and Ag (m=l).
6.
P-OS02C10H7
= 4, 5
= C104
,N03-
The mercury and silver compounds are usually colorless and in some
cases may be recrystallized from hot ^ater. The copper and nickel
compounds are blue and blue-violet; in general they "decompose rapidly
in n°tpr. The authors state that thpse compounds form upon merely
mixing the metal salts and diamines in alcohol or ether, then evaporat-
ing the solvent. The absence of water molecules appears to lessen
competition with the amine groups for coordination positions.
The work of Schle singer (16) suggests another possible structure
for the compounds prepared by Pfeiffer and his associates. Schle sing
found that the inner complexes of copper with bi,s-lmino acids,
CCu(c:-:2)y-(:i:-:c?.R'Coo)
he blue color, general solubility in
.ter, and characteristic conductivity of copper-amino acid complexes
ii
i s
or o .
If x is 1C
the Ois.-jmino acid complexes are violet
For intermediate values of x, both forms
and only slightly soluble,
appear. Schlesinger assumec at ci^- trans isomerism is present,
with the longer molecular chains spanning the trans, positions and the
shorter chains being confined to the cis, positions, as in the amino
acid complexes.
HN C— C
R 0
cis Form
-NH
trans Form
V. Diamines as Metal Deactivators
Small amounts of metals, especially copper, have been found to be
powerful catalysts for aut oxidation of petroleum products and many
other organic compounds. Oomplexing agents function as metal deact-
ivators, reducing the catalytic effect by forming stable complexes
with the metals present. In the case of diamines, deactivation ability
falls off rapidly as the number of carbon atoms separating the amine
groups increases beyond three. This tendency reflects the greater
stability of five*- and six-membered chelate rings. (20)
138
References
1. Bailar and Work: J. Am. Chem. Soo., 6£, 252 (1946).
2. Breuil: Compt. rend., 109, 298 (1934).
3. Dippell and Jaeger: Rec. tr.^v. chim., 50, 547 (l93l).
4. Drew and Tress: J . Chem. Soc, 1933, 1355.
5. Jaeger and ter Berg: Proc . Acad. Sci . Amsterdam. 40, 490 (1937);
Jreger and Bijkerk: ibid. 40. 12, 16, 316 (1937); Z. anorg. allgen
Ohem., 233, 97 (1937)7"
6. Ilann: J. Ohem. Soc, 1920 . 1251.
7. Mann and Poise : Nature, 119, 551 (1927); Mann: J. Chem . Soc, 1927,
1224.
8. LcReynolds: Thesis, University of Illinois, 1938.
9'. Tills and Ouibbell: J. Chem. Soc, 1955. 839.
10. Pfeiffer and Haimann: Ber., 36., 1063 (1903).
11. Pfeiffer: Ann., J5Q3, °7 (l95o).
12. Pfeiffer and Lubbe : J. prakt, . Chem., [2] 136, 321 (1935).
13. Pfeiffer: Naturwiss . , 55, 150 (1948).
14. Pfeiffer, Schmitz, and Bohm : Z. anorg. allgem. Chem., 270, 287
(1952).
15. Rollinson and Bailar: J. Am. Chem. Soc., 65., 250 (1943).
16. Schlesinger: Per., J383, 1877 (1925).
17. Tschugaeff: Ber., 59_, 3190 (1906).
18. Tschugaeff: J. prakt. . Che., [2] 75_, 159 (1907).
19. Tschugaeff: J. prakt,. Che., [2] 7£, 88 (1907).
20. VTatson and Tom: Ind. Tng. Ohem., 41, 918 (1949),
21. Verner: Ber., 40, 15 (1907).
139
Cycl ope ntadienyl Complexes
Victor D. Aftan&ilian Aoril 7, 1953
Gringnard reagents have been shown (l) to form hydrocarbons by
the coupling reaction:
2RKgX + i:x3 -* R-R + 2MgX3 + M
Brown and Lichtenwalter (2) obtained an almost quantitative yield of
biphenyl from phenylmagnesium bromide and ferric chloride. Ferric
chloride is initially reduced to ferrous chloride by the Grignard
reagent, which in turn reacts with the excess of the reagent to form
biphenyl. Several mechanisms have been proposed for this reaction,
the most feasible one being that of f ered" by" I-"ri zewski and Turner (l6).
The-- postulated that probably unstable organometallic intermediate
compounds are formed, which decompose givin~ hydrocarbons and the free
metal as shown below:
2RTlmX + MX3 -> R-K-R + 2!igX2
R-l>?. ~> 1-1 + ?T * '
As a possible route to fulvalene (t j ~~".>N>-- ^ "" j) synthesis
Keally and Pauson (9) applied this reaction to cclopentadienyl-
magnesium bromide, but instead obtained bis~cyclopentadienyliron(ll ) .
•(HMgBr + FeCla -+ Fe (C5K5)3 + KgCl3 + FgBr3
Less than a month before this discovery was reported lTille^ end
co-workers (ll) prepared Fe(CBHB)a by passing gaseous cycionentadiere
over reduced iron in nitrogen atmosphere at 300° p.nd atmospheric press-
ure. The formation of bis-cyclopentadienyliron(lI ) proceeds only~ f or
10-15 minutes, after which further reaction is only effected after
careful oxidation and re-reduction of iron at 450° C and in an atmos-
phere of nitrogen containing steadily increasing proportions of oxygen
and finally air. Kaplan and co-workers (s) modified the Keally and'
Pauson procedure by the addition of one- third of a molar equivalent o^
ferric chloride to cyclopentadienylmagnesium bromide both "in t«tra-
hydrofuran, and obtained oli- yield.
are 1
Some of the physical properties of bis-cvclopentadienyliron(lI )
isted in Table I .
140
X Gt Lit a. •
Physical Properties of 3is-cyclor>entadienyliron(lI )
or "Ferrocene" *
Formula
Formula weight
Melting point , °c
Boiling point , C
Triple point, °C
Fo ( Cp ) 3 * *
186
175-174 (9)
172.5-173 (10)
249 (8'
183 (8
Vapor pressure of the solid logp-^ 7.615 2470 (8)
Vapor pressure of the liquid logPgm~10.27 " ^f^ ^
Heat of sublimation of the solid 16.81 heal/mole
Heat of vaporization of the liquid 11.3 heal /mole
Heat of fusion 5.5 heal/mole
Heat of formation 25.5 heal /mole
Trouton's constant 21.2
E° Fe(0p)3^ [Fe(Cp)2] + + e -0.56***
^in acidic solution) u.^j
Bond distance, Ac
C-C
Fe-C
1.41
2.0
(8
(8'
(8
(3
(3
(12
(16
(5)
(5)
~~ * hoodword and co-irorhers (17) p ropose d th i s n am e .
** The symbol Cp, representing cyclopentadienyl unit,
will be used in this abstract.
*** Latimer convention.
The infrared absorption spectrum of ferrocene shows a single
band at 325-v.,«, which indicates the presence of only one type of C-"
bond in this compound. Wilkinson etal (l6) proposed structure (i )
for ferrocene based upon the above observation.
H.J2/
(i)
This structure, in which the iron atom is symmetrically placed
between two cyclopentadienyl rings, making a "sandwich" type molecule,
has been confirmed by x-ray crystal measurements. (5 and 6). In soite
141
of its high decree of formrl uneaturaticn, ferrocene does not possess
properties typical cf polyolefinic substances. For example, it does
not react with malelc anhydride in boiling benzene; it is not hydro-
genated under normal conditions over reduced platinum (ll) oxide; it
is resistant to the action cf acidic reagents', and has high thermal
stability. Woodward and co-worlcers (l?) carried out several typical
aromatic reactions with ferrocene end reported the following organic
derivative s:
Q - CH3 CfCOH
^/ c>
cL~ OH 3 "coQH
0
diacetylferrocene ferrocene dicarboxylic acid
fl" Q
C-CHaCH80l 8^ -CH=0H3
"&■
C - CH3CH3G1 d - CI->CH2
^8 8
bis-' '.-chlortpropionylf errocene bis-acryloylf errocene
8 f^ __Q }~Z 0G:-3
dimethyl ester of bis-Q-oarboxybenzoylf errocene
Perhaps a more convincing proof for the aromaticity of ferrocence
is the striding resemblance of" the infrared absorption spectra of its
derivatives to those of bsnzens. ■' ' ■ •
(See Table II ) .
A. .
Table II Infrared Band Positions for Opposite Cases
R=ferrocenyl R=phenyl
H-R 3.26 5.27
'.-Ca8C€H3R:rv 5.97 5.93
Q
CH3OC-R 5.82 5.81
0
o-c:-:3ooc-c6h4-c-r 6.02 5.97
Even greater interests are the ionization constants for the
carbaxylic acids, measured in two-to-one ethanol to water solutions.
Ferrocene dicarboxylic acid dKj 3.1x10
pIC2 2.7x10:^
Benzoic acid pi: 1 2.7x10 '
JL'-X.Cj
TJoodv.rord etal (17) postulated that "the very small difference g
between the two dissociation constants of ferrocene dicarboxylic
acid indicates that the carboxyl groups interact very little, and must
be very far apart, while the near identity of the first constant with
benzoic acid demonstrates that the ring carbon atoms of ferrocene,
and thence, necessarily the central iron atom as well, are subsequent-
ly neutral." This observation is of importance with respect to the
detailed electronic structure of ferrocene, since it excludes any form
of hybridization which leads to charge separation ,rithin the molecule.
For example, the possibility of coordination of iron with 18 electrons
(five from each cyclcpentadienyl unit, plus 3 from the iron) to have
effective number of 36, krypton structure, as in the case of ferro-
eyanide, is ruled out. In addition to the objection that the aromatic
properties of ferrocene make it seem most unlikely that all the elec-
trons of the cyclcpentadienyl rings can be involved in the filling of
the orbitals of the metal atom, it seems that a high negative charge
would also be placed on the central metal atom and this we s shown not
to be the case in the determination of the ionization constants of
ferrocence dicarboxylic acid. At present there is not sufficient
evidence to predict the actual electronic structure of ferrocene.
Eiland and Pepin sky (5) are investigating the nature of the bonding
and the electronic configuration of the iron atom in ferrocene by a
method of three-dimensional analysis.
Ferrocene is readily oxidized to the blue [Fe(Gp)2] ce.tion.
Oxidation may be effected anodically, by air in presence of acids
or by halogens, by ferric chloride or eerie sulfate. Especially con-
venient are aqueous silver sulfate or p-be zocuinone in organic solvent
in the presence of acids. The ferricinium ion is reduced by stannous
chloride. Ferricinium ion has been isolated in the form of crystall-
ine salts: [Fe (Cp) 2 ]GaCl4,. CFe (Cp5s]pic^.te, ; endpCFe. (Cp)-33.C&©4.c.-» - •
In view of the above postulations concerning ferrocene, it was
to have been expected that ruthenium (il) and cobalt (ill), -both of
which ere isoelectronic with iron(ll), end also other transition
elements, which have available d orbitals, would form similar
complexes. Table III lists the cy elope ntadienyl complexes which have
been reported in literature to date.
Table III
List of Cy elope ntadienyl Complexes,
Complex, Special Name References
Fe ( Cp ) a + Ferrocene 8, 9, 11
LFe(Cp)2J Ferricinium ion 16
Ru(Cp)3 + Ruthenocene 13
LRu(Cp)3] Huthenioinum ion 13
[Co \ Cp ) 2 1 Cob r. 1 1 i c i ni urn ion 14
Kl(0p)a m++ .15
[Ti(Cp}2j]7 15
[Ti(Cp)3r 15
[Zr(Cp)2]++ 15
[v(cP)2r+ 15
143
Cyclopentadienyl complexes have opened an interesting field of
study in inorganic chemistry; intensive investigations are being
conducted to determine their structures, to study their properties
and to find uses for these compounds.
BIBLIOGRAPHY
1. Bennett, G.M. and Turner, S.E. ; J. Ohem . Soc, 105, 1051 (1914).
2. Brown, R. D. ; Nature, 165, 566 (i960) .
3. Cotton, P. A. and Wilkinson, C-. ; J. Am . Ohem. Soc, 74, 5764(1952
4. Dunity, J. D. and Or^el, L. E.; Nature, 171, 121 (19537.
5. Eiland, R. E. and Pepin siy, M.; J. Am. Cher.. Soc, 74,4971(1952).
o. Fischer, S. 0. and Pfab, IT.; Zeits Nature, 7, 577 (1952).
7. C-ilman, H. snd Lichtem-ralter, M.; J. An. Chem. Soc, 61, 957 (1939
8. Kaplan, L.,Xester, \U L., and ICatz, J. J.; ibid, 74, 5531 (1952).
9. Keally, T. J. and Pauson, P. L. ; Nature, 168, 1039 (l95l).
10. Krizewski, J. ind. Turner, E. E.; J. Ohem. Soc, H£, 559 (1919 ) .
11 • ^^erA^' AV> Tekboth, J. A., and Tremaine, J. F.; J. Chem. Soc,
12. Page, J. A. and Wilkinson, C-.; J. Am. Chem. See, 74, 6149 (1952).
13. ..llhinson, G-. ; Ibid, 74, 5146 (1952) .
14. IJilkinson, G-. ; Ibid, 74, 6143 (1952).
15. Uilkinson, a., Pauson, ?. L. Brimingham, J. !I. , and Cotton, F. A.:
J. An. Chem. Soc, 75, 1011 (1953).
15. h'ilhinson, C-. , Rosenblum, h'., IJhlting. *'. 0 . , and tfoodwrd R B •
J. An. Chem. Soc, 74, 2125 (l952). "" ' " ■'
17- tlTA; si»-tiMl)flun' !;-' rnd ******* "• c--; J- *•• Che-
144
Di sulfur Hep t oxide
April 14, 1953 A. B. G-alun
!I. Berthelot (l ,2, 3) was the first to describe a material which
he called "acide pe r sulfur! cue " . It was obtained by passing either
sulfur dioxide or sulfur trioxide mixed with oxygen through an
electric discharge. The product was considered to be a compound of
heptavalent sulfur. However, Men&ele jef f (4-) pointed out that this
compound must be a peroxide. H. Griram(5; calculated the heat of
f o rm a t i o n :
S207 -> 2S03 + 1/2 02 + 9.7 cal.
A. Hoser(s) found later (1910) that S207 could be prepared in a dis-
charge tube even at 70 C (Berthelot claimed it decomposed at room
temperature). F. Meyer et al(7) repeated Berthelot' s experiments
in 1922 and obtained products corresponding to the formula S3OJ.J..
They assumed the material to consist of a mixture of S03 and S04 and
claimed that Berthelot 's compound T-ras of the same composition. This
point of view was supported by T. Mai sin (8 ) who obtained the same
compound in 1928.
Pure SO 4 was prepared b;,* R. Schwarz and H. Aachenbach (9 ) in
1934; mixtures of S03 and S04 were .also character! zed (lO) . V. "Ja.nn-
agat and C-. Ilennlchen (ll ) prepared pure S-;07 and heave characterized
it chemically* The results of their investigations are summarized
i n the di s cus si o n \ rhi ch foil ows .
Method of Preparation: A discharge tube was evacuated and
filled up to a certain pressure with S02. The tube was immersed in
liquid air to freeze out the S02, and then filled with oxygen gas.
The gases Tr-"'re allowed to mix overnight rat 20° c, and a discharge of
about 1mA and 11KV was passed through the tube. A solution of CaCl2
was used as "coating liquid" . A solid separated as a film on the
walls of the discharge tube .
Berthelot assumed that nitrogen prevents the crystallization of
S207 but does not interfere with its formation. The authors repeat-
ed the experiments in presence of nitrogen and obtained a product
which proved to be (l!0)2 S207 (12) . A. Moser(c) also carried out
his experiments in presence of nitrogen.
If equal volumes of S03 and 02 are employed, and S207 is formed
according to the equation: 4SC2 + o02 — > 2S207 (solid), the gas volume
should decrease to l/3 that of the original. It was found that the
pressure does drop as shown in figure 1, but never to Po/8.
145
pressure
Po
Pf
time
Figure 1
The residual gas was found to contain S03 and S03 but no S307.
and so 2 but
ecuillibria established under discharge
u3 + 1/'?. 02 and SOa^fSCs + l/2 02 . This post-
ulate was confirmed by subjecting S207 (which is quite stable up to
40° C) to a discharge
It was concluded that several
conditions: S20, *- 230
T)re ssure
■ore s sure change s t ar '
ino: with SC2 + 03
pressure change start-
ing with' S207
-time
Figure 2
Anal ~r st s : This ua.s carried out by titrating free iodine and H
after treatment of the substance with a ICE solution:
(S03)x0active + 2i~ + (x-l)Hs0-*5a + XSO4 + 2(X-l)H+
where X is the ratio between S03 and active oxygen. (For all perox-
ides the ratio I +H+ should be 2). SO:
'4 was determined as 3a SO,
In hydrolytic experiments the ratio H2S03::-:2S208-:H2S04, was
determined in presence of each other according to a special method
devised by K. G-leu (l2): Caro's acid is reduced by HBr and the bromine
is absorbed in excess standard arsenite. Excess arsenite is deter-
mined by titration with standard potassium bromate . H202 is then
titrated in the same solution with KhhC4, the endpoint being determin-
ed by appearance of free bromine. Finally, H3S20a is hydro li zed by
strong acid and Caro ' s acid and hydrogen peroxide determined as be-
fore .
Reactions of (S207):: ^) Concentrated HgSO*: Disulfur heptoxide
does not dissolve in concentrated H3S04> whereas S03, S04 and their
mixtures (10) as well as the products prepared by Berthelot and Meyer
dissolved readily. 2) Titanyl ions: Addition of an acid solution of
titanyl sulfate to (S207)x gives an intensive orange colored ring(ll)
S04 does not give a titanyl peroxide reaction. (9)
146
3) Ani line : Aniline is completely oxidized to a tar by (S307)x. S04
oxidizes aniline to nitrobenzene , 4) Mn(ll): UnSO* is attacked only
very slowly in presence of silver ions by (S207)v (at about the same
rate as the slow conversion of (S307)x to Caro ' s'"acid) . S04 oxidizes
-'n instantaneously to IlnOX 5) Ou(ll): (S207)v does not oxidize
Ou(ll) in alkaline solution to the alkali cuprate (ill ) . It is claimed
(9) that S04 brings about this oxidation. 6) Cr207=: ^g a.ichromate
ion is not oxidized by (S307)x ?) Water: fs207)x dissolves readily
in cold water with some gas evolution (probably ozone) and formation
of H2S208, H3SO5 and a minor quantity of HaS04 in solution.
Structural considerations': The proposed structure for (S207)
is
x
0 f 0 0 t 0
j .. .. i
HO: S: 0 J :S:C:C:S:0: I S:OH
X
'6 L b b Jv b
A two dimensional sheet-like structure may also be considered
("This is consistent with the film-like nature of (s207) "-authors).
x
BIBLIOGRAPHY
1. II. Berthelot Compt . rend. 86, 20, 277 (lS70).
2. h. Berthelot Ann. chim. et phys. lb) 12., 463 (1877).
5. M. Berthelot Compt. rend., 90, 269, 351 (l880).
4. 'D. L'endelejeff Protokoll d. Journal d. russ. phys. Ges. (l)
561 (1881).
5. E. 01 ram Compt, rend. 140. 1704 (l90~).
6. A. Moser and N. Isgarischew Z. Elektrochem. 16, 613 (1910) .
7. F. Meyer, Q. Bailleut and C-. Henkel Ber. 55, 2923 (1922).
8. J. Mai sing Bull. soc. chim. Belg, 37, 326 (1928).
9. R. Schwarz and K. Aachenbach Z. anorg. u. allgem. Ohem. 219.
271 (1934). '
10. B. Fichter and A. horitz Helv. Chim. Acta 22, 92 (1939).
11. V. 'fanna-at and C-. Mennichen Z. anorg. u. allgem. Ohem. 268
69 (1952). ° — '
12. E. Terres and II. Cons tantine sou Angew- Cliem. 47, 468 (1934).
13. i: . C-leu Z. anorg. u. allgem. Cher.. 195. 61 (1931 ).
147
SULFAIURIC CHL0RID3
Diamond
April 21, 1953
INTRODUCTION:
Franklin-*- in his development of the nitrogen system of compounds
set up the following tabulation of the nitrogen derivatives of sulfur-
ic acid, or the acuo-ammono sulfuric acids:
s
OH
SO:
_>k
,NH2
SO 3 ^
[hhso2] ^ri [?c-:3c2]3
\:th2
(c) (d)
1 V
so 2hh3
NH '
'^SOaNHp
(g)
( a ) su 1 f am i c acid
(b) su If amide
(c; sulfimide
(&) tri sulfimide
(e ) imidodisulfuric
acid
(f) nitrido tri sulfuric
acid
(g) imidodi sulf amide
SO* OH
(e)
1 !
N(S02OH)3
(f)
These rela.tionsh.ips are somewhat formal and do not imply that
the compounds are necessarilj' pre parable from each other by the
scheme outlined'. Various aspects of the chemistry of the acuo-ammono
sulfuric acids have been reviewed previously.^'0 This report will
consider the chemistry of tri sulfimide , specifically the acid chlor-
ide of trisulf imide .
HISTORICAL
Traube - isolated a number of salts from
chloride with ammonia. Analysis of the silve
and barium salts led him to believe that they
sulfimide. Hantzsch and Holl° established th
exist as the simple momer, [lTHS02], but as th
ring structure was assigned to the compound b
to cyanuric acid; the name sulfanuric acid ha
the compound. Hantzsch and Holl claimed to h
sulfimide but later found that the product wa
.midodisulf amide . ° Further attempts to preps
have been unsuccessful.
ien sulf amide is he
point, the following reactions are postulated
the rea.ction of sulfuryl
r , sodi urn , pota.s slum ,
were derived from
at sulfimide did not
e t rimer (NHS03)s. A
nsed on its resemblance
s been suggested for
ave isolated free tri-
s an impure sample of
re free tri sulfimide
ated above its melting
143
3so2(nh2)2 >(:l-:so2)3 + 3HH3
2so2(:ih2)2 yjK(sotflHa)a + nk3
It has been suggested that trisulfimide is formed from the decompos-
ition of sulfamyl chloride, which may he an intermediate in the
reaction of an aryl chloro sulfonate with ammonia.'
CSH50S02C1 + :IH3 — $NH8S0S<31 + CeH50K
(SOsNH)a + 3HC1
In an investigation of the reaction of phosgene with sodium amide
Perret and Perrot noted the formation of melanuric acid and its
chloride derivatives.0 Because of the similarities existing between
trisulfimide and melanuric acid, it seemed possible that trisulfimid*
may be obtained from the reaction of sulfuryl chloride and sodium
amide. By fractional precipitation of the reaction products with
silver nitrate, silver trisulfimide, (502!;Ag)3, was obtained.
Hantzsch and Hollc prepared trimethyl trisulfimide in the following
manner *
3CH3I + (S02:TiV:)3~^(S02:!GH3)3
Tribenzoyl trisulfimide has also been prepared in a similar manner.
SULJAITUI'IC CKL0?.IDS:
The reaction of sulfamic acid vrith phosphorous pentachloride
was investigated by Ephraim and Gurewitch-, who claimed to have
isolated a binary compound of sulfamyl chloride, based on the follow-
ing equation:
NH2S030H + 2PC15 — >C12 + HC1 + ?CC13 + C1S03HH3 ■ PC13
Kirsanov has shown recently that this reaction produces trichloro-
phosphazo sulfuryl chloride, according to the following ecuation:
NHa-SOgOH + 2PC15«~ — )3HC1 + P0C13 + C1302U=?C13
Trichlorophosphazo sulfuryl chloride is an extremely hygroscopic,
white, crystalline compound; melting point at 35-36°, It is soluble
in organic solvents and reacts with ammonia, amines, alcohol, and
phenol. It hydrolyzes according to the ecuation:
C1S03N=PG13 + 6H20 ^4VC1 + HOS02ONH4 + H3P04
--irsanov"--^ was aole to prepare the su±fur ana.log of cyanuric
acid, sulfanuric acid or l-3-5trichloroxo-l-3-5 trithiatriazine , by
the thermal cleavage of trichlorophosphazo sulfuryl chloride. The"
trichlorophosphazo sulfuryl chloride was distilled in vacuum and
the residue obtained was a transparent brown liquid which contained
the sulfanuric chloride. By fractional crystallization procedures.
two -products were obtained that corresponded to the formula (MS0C1;3;
these were designated by Kirsanov as a- and p- sulfanuric chloride.
The a- sulfanuric chloride is a white crystalline product with a
melting point of 144-145° and a boiling point of 270°; however, the
149
material explodes violent!" upon distillation- The (3-sulfanuric
chloride is also a white crystalline solid with a melting point of
42-43 .
Sulfanuric chloride is soluble in ether and benzene; in alcohol
the solv- tion becomes acidic, probably involving the formation of
esters of sulfanuric acid. Sulfanuric chloride reacts vigorously
Trith ammonia, amines, and is slowly hydroiyzed by water.
Since sulfur in the hexavalent state, lilte carbon tends
primarily to assume a tetrahydral configuration, sulfanuric chloride
may exist as two geometric isomers, represented as the a and (3 forms
N N
S
OC CI 0'
II
II
\ //
y
'Cl
Nci
//
Jl c
A
Cl '0
s
"Cl
Cl
trans
jIO GRAP^Y
1.
Fu
o
L.
Re
3.
C-.
of
4.
w.
5 .
A.
6.
A.
7.
L.
8.
A.
a
IT1
10.
A.
11.
A
. Sisler and II. J. Butler; Ohem
TT
mv
C. Franklin; The Nitrogen System of Compounds,.; Reinhold
blishing Corp., Hew Yorl: (1935) .
F. Audrieth, II. Sveda, H
v. 26, 49 (1940) .
R. Johnson; Seminar, Division of Inorganic Chemistry;
111., Nov. 4, 1952.
Traube; Ber. 25., 2472 (1892).
Hanzsch and A. Holl; Ber. 34, 3450 (l90l).
Hanzsch and Steur; Ber., 30, 1022 (1905).
Denivelle, Bull. soc. claim., E 5, 2143 (1936) .
Ferret and R. Perrot; Compt . Rend., 199 f 955 (1934).
Zohraim and II. G-urewitch; Bar., 43, 138 (l910).
V. Kirsanov; J. den. Chem . , (USSR) 2J3, GG (1952).
V. Kirsanov; ibid, 22, Gl (1952).
-149a-
RCLL CALL
THALLIUM TRIIODIDE
Robert H. Marshall April 21, 1953
Thallium triiodide was first prepared by tickles in 1864.
iSarly methods of preparrtion involved the digestion of Til and
I 2 in ether, methanol, or ethanol solutions, followed by evapo-
ration. A long period of digestion is necessary, however, owing
to the low solubility of Til and an intermediate iodide, TI3I4.
A better method has been reported by Sharpe , in which Til is
dissolved in a solution of i2 in concentrated hydr iodic acid(l).
Evaporation at room temperature yields the Tll3 .
The isomorphism of Tll3 with Rbl3 and Csl3 appears to estab-
lish the structure of the solid as thallium (I) triiodide.
However, the absorption spectrum in methanol solution shows
different maxima than those characteristic of the I3~ ion.
When such a solution is treated with aqueous i\fa2CC3 solution,
T1203 is precipitated. This reaction may be considered to occur
in the following stages:
2 T1I3 ==^ 2T1+ + 2I3"
2 I3~ + 40H~ ===x 2C1- + 41" + 2H30
2 Tl+ + 20H" + 201" + SK20 ==- 2T1(0H)3 + 21"
\
2 Tl (0H)3 ====N T1203 + 3H20
2 Til a + 60H" ===-= T1203 + 61" + 3H20
As evidence for such a series of reactions, it has been
found that when I2_ and alkali are added to a solution containing
Tl+ ions, T1203 is precipitated quantitatively. In addition,
KI3 solutions are immediately decolorized by alkali.
Thermal decomposition of the Til3 yields Til and I2 as
final products. X-ray powder photography, has shown T13I4 to
be an intermediate product. Experiments were also carried out
which indicate the dissociation pressures of various poly-
halides to be:
T13I4 < Csi3 < Rbl3 <T1I3 < KI3.
1. A.G. Sharpe, J. Chem. boc. 1952. 2165.
150
PREPARATION7 OF COPPER HYDRIDE IN
AQU30US 1EDIUM
N, E. Bo jars April 28, 1955
A. INTRODUCTION.
Hyuouhosphorous Acid as. a Reducing Arent. Hypophosphorous acid
reacts with copper (11/ ionn in water, yielding a red-brown reduction
product. The constitution of this product has been the subject of
controversy until recently. It has now been shown! that, under
proper conditions, copper hydride (CuH) can be made in aqueous med-
ium .
Several alkali salts of the monobasic hypo Phosphorous acid can
be made from white phosphorus end the corresponding aqueous alkalir-'
A number of heavier metal salts has also been made.1" The free acid
can be prepared from barium hypophosphite and aqueous sulfuric acid,^
or in resin exchangers. ° Hypophosphorous acid is a colorless, cry-
stalline compound, m.p„ 25.5°C, readily soluble in water. Two hy-
drogen atoms are attached to phosphorus (A).
:*d:
H:P:0*:H H : ? : 0 : H
(A)
:0
(3)
Satisfactory proof has been advanced that a tautomer (b) does not
exist^'"^. However, two forms of the ecid in a slow equilibrium with
each other have b^en postulated-^* H for aqueous solutions. The
active reducing agent, according to Steele-^, is a hypothetical com-
pound H5P02, which is produced with measurable velocity. These
assumptions are necessary to explain the rate of certain reactions ,13
where hypophosphorous acid is the reducing agent* The active form is
produced, when the equilibrium amount is diminished 10* 14. Some ob-
jections against the current structural formulas for H3P03 have been
raised*!^ The ion H2P03 is tetrahedral .- A study involving radio-
active phosphorus has shown that mixtures of phosphates and hypo-
phosphites do not exchange phosphorus , even upon heating to decom-
position temperatures in closed vessels.--'
Hypophosphorous acid is oxidized by many oxidizing agents and
heavy metal ions2;3,4.10 ,11- in certain cases an induction period of
the reaction is observed17 • iL . Hypophosphorous acid is useful for
the production of nickel and nickel hydride alloys1^* "^ via nickel
hypophosphite, for electrode plating, end in ceremics, and as a
stabilizer for inon (ll ) chloride-1. The acid
some metal ions.06
_ — , — .. , _ —
forms complexes with
Hypophosphorous acid is a useful reducing agent0In organic
chemistry. A well-known example is the deemination/'0 involving the
151
replacement of the diazonium group by hydrogen. A method was found
at the University of Illinois-4 for introducing deuterium into
aromatic nuclei, in a similar way . The influence of various metal
salts upon the deamination reaction has been studied and the mech-
anism has been discussed-5. Hypophosphorous acid is also useful in
the production of aryl phos'ohinic aci'ds^ and other organic compounds
The use of the hypophosphorous acid in medicine0 and biology0 ,
as a stabilizer for some polymers01 >32^ and as a sensitizer for
fumigating mixture s°^ can be mentioned.
3. The Reaction of Conner (ll ) Ions with Hvponhosphite Ions.
It has long been known°4 that copper (ll ) ions are reduced by
hypophosphite ions. 'Turtz°5 obtained by such reduction a red-brown,
amorphous precipitate, which he identified as copper hydride.
2 Gu++ + 3 H3P02- + 3 K30 — > 2 CuH + 3 H3P03~ + 4 H+ (i )
Copper hydride then decomposes vrith evolution of hydrogen.
2 CuH. -> 2 Cu + Ha (II)
If excess of hypophosphite is present, hydrogen is evolved even dur-
ing the initial stage of the reaction; with an excess of copper (ll )
salt copper only is precipitated0^;0?. The constitution of "the
product of the reaction (i ) has b<=en a subject of controversy0® until
recently. Thus, it has been maintained*'*0" that spontaneous decompos-
ition and X-ray diffraction studies indicate that "the product is a
solution of hydrogen in copper as is the case >rith palladium, and
that the red-brown precipitate is not a hydride analogous to the
alkal i hydride s .
The nature of metallic hydrides has been investigated recently*-^
T'Jiberg and co-workers4(3>41>^:2 have developed methods for the prepar-
ation of hydrides in non-aqueous media. Lithium eluminum hydride40"
was found to be useful for thp preparation of dry corvner hydride.
4-CuI + LiAlH4 -* Lil + A1I3 + 4 CuH (ill)
Copper hydride c.°n be dissolved in pyridine, end reprecipitated by
ether. Thus it becomes evident that copper hydride is, after all, a
definite chemical compound.
C . The Procedure for the Preparation of Co'oner Hydride in
Acueous Hedium,
To obtain the. desired results, the old method of Wurtz0*^ was
properly modified1. According to the equation (l ) , stoichiometric
amounts of 0,4 M acueous copper (ll ) sulfate and 0.3 II aqueous sodium
hypophosphite containing 1-25$ free sulfuric acid were allowed to
r-act in a hydrogen atmosphere at 20°C for 15 to 20 hours, Frecautior
were made to exclude traces of oxygen, since otherwise some copper (i )
oxide appears as an impurity in the product.
T-Jet copper hydride was precipitated: all attempts to dry it led
a decomposition1*44, even below 45°C. However, a satisfactory analys:
was obteined by a special procedure1. In more than 100 separate
experiments a light red-brown product was always obtained1. It was
shown by analysis to contain 98.45$ Cu and 1.55$ K. in excellent
152
agreement with the calculated values for CuH.
The earlier workers had obtained analytical values up to 1.50^
K and about 98.50/3 Cu . An aqueous suspension of copper hydride has
been shown to act as a reducing agent; an attempt to devise a quan-
titative titrimetric procedure has not been entirely successful be-
cause of difficulties in estimating the end point1.
Thus it has been definitely proved1 that copper hydride can be
made in aqueous medium. Silver (l ) ions yield silver metal only under
similar conditions*^.
BIBLIOGRAPHY
1. G-. Wagner and E. Lohr, Oesterr. Chem . Ztg. 54, 16 (1953).
2. K. A. Hoffmann end U. R. Hoffmanrj Anorg-rnische Chemie, Vieweg,
Braunschweig, 1945, p. 249.
3. .J. A. N. Friend, Textbook of Inorganic Chmistry, G-riffin and Co.,
Ltd., London, 1934, Vol. 6, pt . 2, p. 135 f f .
4. T. Moeller, Inorganic Chemistry, J. Wiley and Sons, New York,
1952, p. 539 ff.
5. R. Element, Z. anorg. Chem. £30, 2S7 (1949).
6. A. Simon and F. Feher, Z. anorg. allgem. Chem. 230. 269 (1937 ) .
7. T:.T. H. Zachariasen and R. C. L. Kooney, J. Chem. Phys. 2, 34 (1954).
8. A. Ferrari and C. Colla, Gazz. chim. ital. 57_, 294 (1937).
9. H. Erlenmeyer and H. Gaertner, Helv. chim. Act,-1 17, 970 (1934 ).
10. A. D. Mitchell, J. Chem. Soc. 117, 1322 (l92C).
11. R. 0. Griffith and A. McKeown, Trans. Faraday Soc. 30, 530 (1934-).
12. B. D. Steele, J. Chem. Soc. 91, 1641 (1907),
13. A. D. Mitchell, ibid. 119, 1265 (1921 ) .
14. I. II. Kolthoff, Pharm. ."eekblad 53, 909 (1913); 61, 954 (1924).
15. V. 11. Plets, J. Gen. Chem. U.S.S.R. 2, 84 (1957).
16. V. D. Ionin, A. F. Lukovnikov, K. 3. Neimen, and A. N. Nesmevanov,
Doklady Akad. Nauk S.S.S.R. 67, 463 (1949).
3. Hath Sen., Gazz. chim. ital. 78, 423 (1948) .
B. Nath Sen., Collection Czechoslov. Chem. Commun. 10, 321 (1938).
P. Kersch, J. Chem. Education 20, 376 (1943).
R. Scholder end K. Heckel, Z. anorg. allgem. Chem. 108., 329 (l93l)
G-. C. Walker, W. J. Maxted, and F. N. Hughes, Can. Pharm. J.
83, No. 7, 12, 14, 52 (l950).
S. Banerjee, Science and Culture 1_6, 115 (1950).
Organic Syntheses, J. Wiley and Sons, Inc., New York, 1941, vol.
21, p. 30.
E. R. Alexander and R. E. Burs:e, Jr., J. Am. Chem. Soc. 72, 3100
(1950).
N. Kornblum, G-. D. Cooper, and J. E. Taylor, J. Am. Chem. Soc. 72,
3013 (1950).
V. !!. Plets, J. Gen. Chem. U.S.S.R. 7, 84 (1937).
K. Schmidt, Chem. Ber. 81, 477 (.19487-
N. Kornblum and D. C. Iffland, J. Am. Chem. Soc. 71, 2137 (1949).
V. I-I. Plets, J. Gen. Chem. U.S.S.R. 7 90, 273 (1937).
A. Takayima, Actn Photochimica (japan) 1,3, Ho. 1, 1 (1942).
H. \I. Gray and G. H. Latham (to E. I. duPont de Nemours and Co.),
U. S. Patent 2 '516 '980, Aug. 1, 1950.
K. *.T. Gray (to E. I. d'U:ont de Nemours r>nci Co.), U. S. Patent
2'510'777, June 6, 1950.
153
33. J. Taylor and J. M. Holm (to Imperial Chemical Industries Ltd.)
U. S. Patent 2»532*349, Dec. 5, 1950.
34. K. Rose, Ann. 9, 225 (1827); ibid. 58, 301 (18-1-3).
35. A. tfurtz .Compt. rend. 18, 702 (18447: Ann. Chim. Phys. 11 [3 T,
250 (1844); ibid. 16 [3], 190 (1846).
36. A. Sieverts, Z. anorg, Chem. 6_4, 29 (1009).
37. 17. Muthmann and F. Mawrow, Z. anorg. Chem. 11, 268 (1896).
38. M, Berthelot, Compt. rend, 89, 1005 (1879).
39. N. Weichsenf elder, Ann. 447 t 46 (1926).
40. E. T;.riberg and Th. Johannsen, Die Chemie 55, 38 (1924).
41. 0. Stecher and E. Wiberg, 3er. 75, 2003 (1942),
42. E. 71 berg, Lecture, May 19, 1952, Frankfurt /Main, General con-
ference of German Chemists.
43. A. E. Finliolt, A. C. Bond, ?nd H, J. Sohlesinger, J. Am. Chem.
Soc. 69, 1200 (1947),
44. J. C. T.arf and 17. Feitknecht, Kelv. chim. Acta 33, 613 (1952).
154
INORGANIC P.APS3 GKRCHATOG-RAPHY
Ilary .Joan* Si ro tele May 5, 1953
Inorganic analyses on strips of absorbent paper were first
described over a century ago. V'eil and Williams "(l) credit the first
recorded experiments on capillary analyses to F. F. Runge, who in 1850
analyzed a mixture of dyes on blotting paper and whose interest extend-
ed to the possibility of using capillary ascent of solution in blocks
of wood for the separation of solutes. At a later date this was more
fully investigated by Schoeribein and his student, Goppelsroeder.
Schoenbein, who has usually been regarded as the pioneer of capillary
analysis of inorganic compounds, showed that if a strip of paper is
dipped into water containing inorganic salts, the water rises in the
paper and carries the salts with it. The solutes concentrated in
distinct zones as the solution was drawn into the strip. He predicted
its later use as an analytical tool (2). In 1873, Bavley (3) notpd a
similar phenomenon. Fischer (4) knew of this work and applied it to
the quantitative investigation of the separation of sodium and barium
and of certain double salts. By separating the cations in a double
salt, such as ferrous ammonium sulfate, he demonstrated the absence of
a complex ion.
The true discovery of chromatography is credited to TsTrett (5;
who separated the pigments in leaves by passing a petroleum ether
extract of leaves through a column of calcium carbonate. He named the
method and stated, erroneously however, that it was based on purely
physical reactions. Tswett vigorously denied that his work was relat-
ed io the worn of Schoenbein end loppelsroeder.
In 1909, Krulla (s) recommended adsorption on filter paper as a
simple qualitative test for salts which give a color taction and as a
measure of the concentrations of dilute solutions.
Chromatography was then nearly completely forgotten for almost
. iff-rs; fUt lately mUcn has beer- ^one to develop the usefulness
ox tnis tool.
. ..Strain (7) has defined chromatography as the technique of
analysis and/or preparation in which there is a dynamic partition or
aistributlon of dissolved or dispersed materials between two immisc-
ible pnases, one of which is moving past the other.
j^Jwomptographlo separations may be run on either columns
or
paper stripe*
BT,nt ^ Picture for paper chromatography is essentially this: a
nnS - I n0j^°e °Ve n-nr-zed is applied to one end of the paper
*"Lt *f sloping solvent is made either to ascend or descend the
paper, with the result that different solutes appear in different
T55
Many variations of this simple procedure have appeared using
simple filter paper or blotting paper in the form of strips, disks,
sheets, and piles*. Cons&en et al (8) and Strain (9) have described
a descending method. V/illiams and Kirby (lC) and dockland and
Dunn (ll ) used an ascending procedure. Other variations of the
method have been developed by Brown (12), Iluller and Clegg (13),
Hutter (14), Ma and Fontaine do) and Datta et al (16). Attempts to
n.*OT)ly par>er chromatography to a larger scale have been made by
Mitchell* and Easkins (17), Porter (18) and Zanof sky et al (19 ) .
Burstn.ll and Kember (21 ) have given an excellent descriotion of
the practical problems involved in running a chromatogram. In
a neral the following things must be considered:
The, test solution
In general the concentration should be about 1 mg./0.05 ml.
for a one inch -fide strip.
b . Transfer of te".t solution to oao- r strip
For qualitative analysis, n C.05 ml. pipet may be used.
For quantitative analysis a micrometer syringe or capillary
buret should be used. Novellie (21 ) has described a simple
method of application.
c . "lyp en and r~ 1 s e 0 f ~o ■*• ^ e r
A strip one inch wide and eight to sixteen inches long of
T/Jhatman No. 1 or o is usually satisfactory. Scribner and
Wilson (22) have stated that paper cannot be considered
inert even towards water and have set up standards for
filter paper. Gas sidy (23 ) has found that some paper gives
bands due to the presence of copper. Different paoers may
even invert th - order of the spots. The ' machine direction
of the paper should also be considered. Flood (24) has usee1
paper Impregnated ~r*.th alumina.
d. Atmo spheric condi tl ons
^'r.e chromatogram should be produced in a closed tank so that
the atmosphere will be saturated with the vaoor of the
solvent to prevent evaporation.
e . Temper^ ture,
\l'-ze temperature should be kept somewhat constant. Gordon
(25) states that with- rising temperature -the adsorption; de-
creases' and the* salt becomes more 'motile.
f . Time o.nd distance of solvent movement
The time varies for fifteen minutes to thirty-six hours.
For complete sepa.rn.tion of a large group, the solvent front
should move twelve inches. For some separations, two to
three inches is sufficient.
g. Choice of solvent
The best method is to try one of each of the following
types: alcohols, ketones, ether, carboxylic acids, esters,
etc. When the best type has been found, different members
of its homologous series are trl - d, usually with the add-
ition of 1-2^ mineral acid. The solvent is usually saturat-
ed with water. La court et al (26) have shown the importance
156
of choosing the proper initial solvent.
Development and dete ctic.r. of z^nea
Physical methods include use of fluorescent indicators (27),
photoelectric cells, interferometry , and. radioactive tracers
(28). Chemical methods may consist of spraying or st re dicing
wi t h a color pro :h x c i ng re age n t .
A theory to be ndequa/fce should describe the formation nnd
development of zones, their rnte of movement, and their spatial as
well as concentration dimensions. It should make it possible to
calculate beforehand the state of development after a given amount
of developer has been applied and should clarify the findings of
experiment. Uiifortunrt.pl;; the time has not arrived when a new sub-
stance can be placed on a paper, a developer added, and the results
ac j1 irately pre di c t e d .
Mr.ny attempts heave been made to develop a theory. Wilson (29)
has devised a theory in terms of the adsorption isotherm which is
used to express the relationship between solute and adsorbent at
equilibrium. However, this does not tahe into account the effect of
one solute upon another at varying concentrations. He neglects all
diffusion and Inch of equilibrium.
In 1941, Martin and Synge (eC) introduced partition chromato-
graphy, in which the difference between partition coefficients in
the aqueous and nonaqueous pha.se s of the components of a mixture
gave refinement in the normal chromatographic separation which depends
upon adsorption characteristics. In 1944, Consden et al (8), extend-
ed this to the use of paper strips. Martin and Synge developed a
piste theory for chromatography analogous to that used for distillat-
ion columns. They consider the column or strip to consist of a
number of theoretical plates (H.S.T.P.) within each of which perfect
equilibrium occurs nn<?. relate the partition coefficient to the rate
of movement of the bands.
Thus, A = cross-sectional area of paper + nonmobile phase + mobile
phase
AT = cross-sectional area of mobile "iho.se i.e., solvent "oha.se
A3 = cross-sectional area of nonmobile phase i.e., water phase
a = partition coefficient = cone, in HPD phase,
cone . in solvent ph a s e .
R-p = movement of band = HAt A
x movement 01 advar.cina front of lieu id -— *•= ~ "pL
X
a -A - AL _ Ar
c:As
Rf%
AS V "*
- 1
Cther attempts at development of a theory have been made. Flood
(31) has related the zone radii to the cor centration- Hopf (32) has
developed Flood's formula further. Brimley (33) .assumes that the
spots spread by diffusion in a wa~~ analogous to the theory of heat
flow. Thomas (34) discusses n theory of nineties leading to a Lang-
muir type of isotherm at equilibrium.
However, no completely satisfactory theory has been devised.
157
IECHA1TI m
All chromatographic methods -re based on the differential migrat-
ion of solution through polyph.ase systems in Trhich the phases have a
preferential affinity for the solutes.
In considering the mechanism of chromatography, Arden, et al
(35; have listed these factors as requiring consideration.
Selective extraction of the salts of the organic solvent,
an effect prevailing at the test patch.
-u
l/here substantially immiscible solvents are used, partition
of the inorganic substance between organic solvent and
aqueous layer, since water is present as r. normal constit-
uent of the adsorbent paper and is usually added to the
organic solvent. This partition therefore takes place as
the liquid mixture moves down the paper, and separations
are due to slight changes in conditions at different points
on the adsorbent strip.
c. Adsorption of the metallic ions by the paper.
d. Formation of complexes with high solubilities in organic
media under specific conditions.
Feigl (36) has compered the mechanism to n process of fraction-
al precipitation. Iluller and Clegg (57) have described the mechanism
on circular filter paper. In the chromatographic field, it is often
a moot cuestion as to whether the separation is due to adsorption,
ion exchange, liquid-liquid distribution or ?. combination of tT'o or
three of these. The predominant force varies with the solvents and
the solute. Thus, Pollard, et al (27) state that for butanol mixtures
in cation analysis the separation is essentially partition chromatog-
raphy whereas in collidine, complex formation end selective extraction
play the dominant role .
A?"0 LI C ATI 0" T S
The most obvious application of paper chromatography is the
qualitative detection of substances. Although most of the ground worl
has been done with amino acids, recently applications to the detect-
ion of cations and anions have become more abundant, Lederer (38)
has devised a method for the separation of the noble metals. Arden
et al. (35) have found a method for the separation of several metal
ion groups, namely, Oa Sr Da; Al G-a In Zn; Co . Cu Fe Mn Ni ; and Pb Cu
Bi Cd Eg. A more intensive investigation by Pollard, et al (39) has
resulted in a scheme for the separation of Pb, Ag, Hg, As, Sb, Cr, lln,
Cu, Co, Mi, Pi, Fe Sn, Sr, Pa, Cd, Zn, Al , hg, Ca, Na, and K. Later
Pollard, et al (40) extended the scheme to include Ce , Li, ho, Tl , Ti,
U, V, and <■''. Anion analysis has developed more slowly, but Lederer
(4l) has devised a scheme for the separation of the chloride group
anions. Fillinger (42) has devised a simple scheme to be used in
elementary qualitative analysis. She makes use of alumina columns
but the procedure could also be aTDolied to nnoer.
153
A method of two dimensional chromatography has been discussed
■TH
ne
by Cons den (43) and To I ley (44)
method can be used for identification of nore difficult substances
iD3 produced by this
s ince
,v.ro
olvents may
used .
Methods for quantitative analysis have developed nore slowly.
For semi-quantitative work, the stain is compared with those made by
known amounts of substance. For more accurate work, the color band
Is cut out and the metal content determined accurately by color-
imetric, polaro graphic, and spectograp
-In'n
techniques. Fischer (45)
has described several simple methods for the quantitative assay of
amino acids. This method is applicable to other problems. Bull
et al. have plotted the % transmission on semi log paper versus
the distance along the strip and by means of a planimeter have
estimated the concentration of amino acids.
Anderson and Lederer (4?) have combined quantitative analysis
v;ith electrochromatography on paper. Strain (43) has used two and
three way electrochro.:)3togrsphy lor the resolution of mixtures.
An application for the determination of molecular structure
has been described by Strain (9)« He has done this work by study-
ing the adsorption sequences and relative rates of migration.
Strain (7) has also studied the eifect of ionic charge, radius and
uass on the degree of separability.
ADVANTAGES AND DISADVANTAGES
Some of the advantages of paper chromatography include its
use for the separation oi substances too unstable to be distilled,
the concentration of dilute solutions, and the determination of
the purity, homogeneity, and components oi solutions. It is also
3 rap id , si mole , and ine xpe :is iv e me t hod .
Unfortunately, at the present tine there is a considerable
degree of empiricism involved in chromatography. There is a great
deal of physical data to be worked out. Taper chromatography is
not applicable to large scale work. There is always the chance
of chemical change in the adsorbed solute.
Because there is some confusion in nomenclature, Dent (49)
has recommended the use of the word "papyrography" as a substitute
for
r chro na to ;ra-
■-v,
'U t
hi
u :r. qe s t i on ha s no t been c a rr ie 3
in this subject, especially in
.udies of its physico-chemical
nature, including kinetic studies so as to evaluate the various
factors involved in a nore quantitative and predictable fashion.
I-iuch work remains to be done
the way of further quantitative s
159
BIBLIOGRAPHY
1. H. Veil and T.I. vTiiliams, Nature 167, 905 (l95l).
2. C. ?. Schoenbein, Verhnnd naturfoesch Ges Basel III, 249 (l86l).
5. T. 3ayley, J. Ohem . Soc. oo, 304 (1878).
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19
20
n>n>i
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16C
■-D.
'1-6.
47.
4£.
49 .
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472
\ -4- V U* '-> /
161
INORGANIC COLUMN CHROMATOGRAPHY
S. A. 3artkievi.cz May 12, 1953
Chromatography, a simple and ingenious method of chemical anal-
ysis by adsorption, was invented in 1906 by II, Tswett, a Russian
botanist, while working with plant pigments. The method was used
sporadically in biological research; but, it was almost thirty years
Later before chromatography found a.ny serious application in the
inorganic field. (24,27,33; It has been used in inorganic chemistry
chiefly for the resolution and identification of mixtures of cations.
Chromatography has also been applied to some extent in the separation
of elements, detection of anions, purification of inorganic compounds,
and the separation of isotopes. Some attempts hove been made to apply
this method quantitatively.
Strain (24) describes the mechanism of chromatography as a dynamic
process which depends upon the continuous and simultaneous adsorption,
desorption, or elution and readsorption of the substances being
resolved upon the column. Concepts regarding the details of the
adsorption process itself are not entirely clear. (16,19,28,34) In
adsorption columns several types of adsorptive forces may come into
play. These may be of the physical type, Van der h'aal f s forces, or
ential chemical reaction, chemadsorption .
The choice of column size, adsorbent e,nri solvent, is determined
to a la.rge extent by the quantity and the adsorbability of the
materials to be resolved, the physical and chemical nature of the
adsorbent, the solubility of the materials to be adsorbed, and by the
activity and solubility of the adsorbent, (16,19,27,30) Final select-
ion is usually, however, made by empirical methods. (27) Le Rosen (37)
attempted to standardize chromatography, and suggested a method for
determining the characteristics of an adsorbent which Trould aid re-
searchers in reproducing the work of other people .
Terminology indicates limitation to colored substances only.
Modification of the process and the procedure, however, makes it
possible to work with colorless materials and to locate colorless
bands. (9,27)
It should be kept in mind that most of the above material has been
derived and postulated with reference primarily to organic materials.
Much of it undoubtedly does apply to inorganic materials as well.
The first extensive T'rerk done in inorganic chromatography was by
Schwab and Jockers. (l9) They examined a variety of substances which
could be used as adsorbents and used several of the standard reagents
for developers. Using A1203 as an adsorbent and primarily H2S and
0IH4)2S as developers, they obtained a definite series order of
adsorption for a number of cations. Working -r!th Dattler (20,21)
Schwab found that by treating A1303 with HN03 an anionic series could
be obtained — the cations being used as developers. An endeavor was
made to arrive at the quantitative possibilities of chromatography by
attempting to correlate band widths with concentration of the ion.
Srikantan and Krishman (25) applied the method to the Quantitative
162
determination of various constituents of alloys such as brass and
steel. They report that 99$ accuracy can be obtained by careful con-
trol of the method. Xu.bli (l? ) investigated the work of Schwab and
Dattler v20) with anions and extended the series order to almost twice
its original number. Schwab and Ghosh (22) investigated the possibil-
ities of the
Some success
Bishop recomme:
analysis and cites two examrd.es that car/ be used, ~Fi 111 nger"* and
Troftan (15; have reported the successful application of chromatography
to some degree in undergraduate qualitative analvsis at Rollins Coll1 "
ege, Virginia. Schwab and Ghosh (23) investigated the use o^ chrom-
atography in micro analysis and found that the sensitiveness of' t^e
metnod is as great and in some cases greater than by the droi) method.
±hey also made separations of the noble metals, u si tip- A1203 as nn
adsorbent. No developers -ere required since the compounds used were
SemSni5SS Cf^^' JT-?"'1 SRL?STWOr!-ers (29) 'found that the presence
of ,001/Co, 0.001/ of ,:i, 0.0017 of Zn, and 0.57 of Cu could be
determined using an A1203 column and (UK4)3S as n developer. ' Ventur-
er i±o and Saini ^32j report the separation of the Pt ^ovv noon Al C
using various developers. Venturello and Ao-liardi (31) investigated
tne work of Schwab and Jookers (19) and further extended the cationic
series, *rlenmeyer and co-workers also extensively investigated the
use of chromatography in inorganic analysis. Using 8-hydroxyouinoline
n.s the adsorbent Srlenmeyer and Dahn (l2) succeeded in resolving
series of ions. No developer was needed since the oxime itself "served
tne purpose. Working with Violuric acid as an adsorbent Erlenmev-T*
and Schoenauer (14) were successful in seoarat^ na ua+ arid K as well
as most of the alkaline earths. Erlenmeyer and SohmidUn (13)' extend-
ed ,ne work upon the Ma-group, Robinson (33) working with n mixture
of soarch and 8-hydroxycuinoline further extended the series order of
Erlenmeyer and Dahn. He also successfully determined the Zn con?en?"
nLn hnf- br"sin- rll°^rs °y the *et^ of band width measurements.
ffll uS1?g ? 2°^binVtlon of chromatography and colorimetry
successfully isolated and determined Cobalt in ferrous and non-ferrous
alloys. Several successful separations of ions have been reported
using cellulose as an adsorbent. (4,5,7,3,34) Kutzelni«r (18) usine
ZnS as an adsorbent successfully obtained' the following levies ordef-
Au, 0s, Ru, Pd, Kg, Ag, Bi, Cu, Fe , U, Pt, Rh, nPd Ir.
of sulfides produced sharply defined zones which can b«
fork
in/*4s f}f}dJs stl11 fragmentary due to the lack of proper
~. Ubj Jith uhe development of a more unWerssl adsorbent
153
..L »
o
T5T T
OG-RAPHY
1. Al-Mahdi and Hilson, Mikrochemie ver. I-likroehim. Acta, 40, 138
(1952).
2. Bach, Industria y quimica 12, 283 (1950).
3. Bishop, J. Chem., 3d. 22, 524 (1945).
4. Bishop and Liebman, Nature. 1,67 f 524 (l95l).
5. Buriel and Perez. Anal Chim Acta 3, 438 (1949)..
3. Burstall and Co-workers, Discussions Furaday Soc. ? 179
(1949). ' ~"
7. Burstall Kember and "Tells, J. Elechode^ositorr s Tech. Soc.
27, (1951).
Bur stall a nd "ell s , Analy s t . 23. , 396 ( 1 9 51 ) .
Cassidy, "Chromotograohy" Academy (l940)„
Dean, Anal Ohem. 23, 1093 (1951).
Dun.^bin and Co-workers, Nature 134, 915 (1949).
Erlenmeyer -^nd Dahn, Helv. Cher.. Acta 22, 1339 (1939).
Erlenmeyer and Schmidlin, Helv. Chem . Acta 24, 1213 (1941 ).
Erlenmeyer and Schoenauer, Helv. Chem. Acta 24-, 879 (1941).
Fillinger, J. Chem. Ed. 24, 444 (1947).
Jacobs ana Tompkins, Trans-Farad. Soc. 41, 388 (1945 ).
Kubli, Helv. Chem. Acta, 3_Q_, 453 (1947).
Kutzelnigg, Z. Erzleergban u. Metall. Hultenw. 3, 77 (1950).
Schwab and Jockers, Angew. Chem. J50, 546 (1937).
Schwab and Dnttler. Angew. Chem. 50, 69? (1937).
Ibid, 51 706 (1938).
Schwab and Ghash, Angew. Chem. 52, 633 (1939).
Ibid, 53 39 (1940 ).
Smith, "Inorganic Chromatography" D. Vnn llnstrand Company, Inc.
New York, 1953.
Srikantan and Kushman, J. Indian Chem. Soc. 26 415 (194^).
Strain, Anal. Chem. 22, 41 (1950).
Strain, "Chromatographic Adsorption Analysis," Interscience
Publishers, Inc., New York, 1942.
Sacconi, Nature, 164, 71 (1949).
Tanaka and Co-workers, Chem. Research (japan) 55 (1949).
Trueblood and Malmberg, Anal. Chem., 21, ic55 (1949).
Venturello and Agliardi, Ann. Chim. Applicata, 30, 220 (1940).
Venturello and Saini, Ann. Chim. Applicata, 39, 375 (1949 ).
Robinson, Disc. Fura. Soc., 7 195 (1949).
Williams, J. Chem. Soc. (1952) 1497, 3155, 3899.
Hilson, J. Am. Chem. Soc. 32, 1583 (1940 ).
Zechmeister, "Principles and Practice of Chromatography,"
John Hi ley and Sons, Inc., New York, 1941.
Le Rosen, J. Am. Chem. Soc. M, 1905 (1942).
' '■>•■ '■ ,
164
iMMONOLYSIS AND At-ilNOLYSIS CF SOME
SUBSTITUTED SI LI 301: KALIDES
A« s- Hay Kay 19 f 1953
In general the halosil'anes are uniformly more reactive than
their carbon analogs-. Trichlorosilane and tetrachlorosilane are
hydrolyzed even by moist air. As the halogen atoms are replaced by
alkyl or aryl groups the ease of hydrolysis de creases2'3 and the /
stability also increases as builder groups are -ut into the molecule ±.
'.Tith Increasing bond polarizability the ease of hydrolysis also
increases in the order F, 01, Br, I.
All of the halosilanes (with the exception of the fluoro
compounds) react with ammonia at room tern •■ ratvtre, or with liquid
ammonia--. In this manner the silicon analogs of primary and secondary
amines can be prepared0* ■ . Miner et al8 treated t-alkoxychloro si lanes
with ammonia and found that the type o^ product resulting from the
reaction depended on the ratio of chlorines to silicon in the molecule
The mono- and dichloro si lanes formed the corresponding mono- and
diaminosilanes, but the trichlorosilane s reacted with ammonia to give
only resinous condensation products. Trisilylamine prepared from
tetrachlorosilane and ammonia, is the only known analog of a tertiary
amine. They also found that the presence of- at least one tertiary
alhoxy group in the molecule is essential for the existence of "a "
staole compound containing two amino groups en the same silicon atom.
In contact with water these t-alkoxyaminositanes react rapidly
with evolution of ammonia to give silanols. They also react rapidly
with most primary and secondary alcohols at room/ temperature with"
bhe evolution of one mole of ■ mmonia.
(HO)2Si(rTH3)3+-10^(HC) 2(^0)81:^3+^3
l/ith methanol both amino groups react readily but heat is needed
for tue reaction with higher alcohols. At elevated temperatures with
primary alcohols the second amino grou-o reacts readily with secondare
alcohols the reaction is slower, and with tertiary alcohols the first"
amino group reacts slowly and the second not at all.
Generally trialkylchlorosilanes react with ammonia to give the
corresponding disilazanes but 'trie thylchlorosilane is reoo^ted to ViVe
some triethylaminosilane also. • ' " " " ° "
. Similarly the reaction of tetrachlorosilane with alkylamines
yields tne corresponding N-alkyidisilazanes, Trime thylchlorosilane
l?Lme^ll^ine rrC%t0 f°rm ^imethyl-N-methylaminosilane,, Reaction
witn a furtner molecule of trime thylchlorosilane gives heutametbyl-
cLisile. zane • J
(GH3)3Si01+aH3^IH^>(CK3)3SiNHCHff->[(0H3)3Si]3IICH3
Qm^v,9y Zlr'°tl0,\0f ammoni-°- with clime thyldichlorosilane Larsson and
Kanetaf ^^^i1?*}0^^1*™™ (I) ™d octamethylcycloretra-
-iJ.azane Ui ; . Botn are hydrolyzed by water.
165
NH Si ( CH 3 ) $ JJH
Si(CK3)s . Si(CK3)3
! |
NH Si ( OH 3 ) 3 * M
I II
Dimethyldichlorpsilane and methylamine or ethylnmine in a 1:4
ratio yield as the main products dimethyl-bis-(methylaminoVsilane and
dimethyl-bis-(ethylamino)--silane, respectively.
4RIH2+(CH3)aSiC!lg->(CH3)2Si{riKR)2+2HNH3Cl
Schwarz and -'Jpi^el10 prepared some amirosilanes and studied their
conversion to the disilazane s. They prepared triisopropoxyaminosilane
by introducing dry ammonia into rn ether solution of triisopropoxych"> o-
rosilane. It is very sensitive to moisture. Mien heated under reflux
ammoniP is evolved and after 40 hours a 50 percent vield o^ the '
disilazane was obtained. \Jhen triphenoxy amino si lane , prepared in a
similar manner, is heated it is converted :uantitatively to the
di si If zane .
The preparation of diisopropoxydiaminosilane ima attempted in n
manner analogous to that for the monoaminosilanes. It could not be'
prepared but instead a mixture of tetraisopro^cxydisilazane (ill ) and
hexaisopropoxycyclotrisilazane(lV) was obtained. In order to determine
whether the cyclic compound was formed by r further condensation of
[(RC)3Si]2NK <ro)2 Si— HH~*S3 (0R)2
\ /
II
NH
V
(0R)2
IV
the disilazane, some of the latter was heated for some time at 180 der
Jnaer tnese conditions - mixture of the cyclotrisilazane (iv) and
cyclotetrasilazane (V) was obtained. These compounds, as all other
(no) a si m si (or) a
* \
NH NH
f ^ I I
(R0)2 Si MH —Si (OR)
V
aminosilanes are sensitive to water, ammonia being split out.
pnp n^i^n°^ci:l0]TSi^ne "nd hex™e^ylenediamine react in methyl-
ene cnloride at the boiling point to give a mixture of triphenoxv-
aminosilane and triphenoxyhexamethyleniminosilane (VT). After six enyo
2
(0o)3si:
i
\
166
- -•» CH 2 " — CH 3
CH2— CH3— CH2
VI
in a desiccator over phosphorus pentoxide the cyclic compound crystall-
ized out of the mixture. In contrast to other amino si lanes and disil-
azanes described, it is comparatively stable to moisture. Several
hours in contact with water at room temperature produces only a slight
amount of hydros/sis.
In an attempt to prepare a molecule containing the grouping -
SillHCO - four moles of triisopropoxyaminosilane and one mole of
adipoyl chloride were condensed in refluxing ether solution. A mix-
ture of triisopropoxyadipimidosilane (VTl) and hexaisopropoxydisilazane
was obtained. The cyclic compound could be purified by crystalllzatlor
(R0)3si:
nil
■ 3 ^n 3
< r\ n-_* __ rfq
vi :
(R = (CH3)2CHO-)
from benzene. It is extraordinarily susceptible to hydrolysis and
gives by reaction with wr-ter, ammonia, adipic acid, and hexaisopropox;
disilazane which undergoes further hydrolysis to yield hexaisopropoxy-
disiloxane .
31 BLI OGHAPHY
1
2
3
/i
5
6
n
ri
D
9
10,
a*
R.
C.
a
^ «
E .
lman, and
o .
Gi
ha
A.
C.
3.
Lr
Sc
G-. Z
ur&hard, and
Emeleus, and
. Dunn, Chem. Revs., 52, 77 (1953
E. C-. Rcchow, -ibid. , 41., 97 (1947
D. S. \yne. J. Chem. Soc, 1590 (1947).
m.
lman, and R. N. Clark, J. Am. Chem. Soe.; 59, 1499 (1947).
rsden, and F. S. Kipping, J. Chem. Soc, 93, 198 (1908).
&raus,
Sauer,
Miner,
and
nnd R. H. Hasel
1 >
Jl
Pedlow,
rsson, and
Jr'
Ind.
Smit
P,V>-
)
arz, and F. T.!eigel,
952).
J. Am. Chem. Soc, J56, 197 u.
ibid., 68, 241 (1946).
R . ? . Holy s z , Jr . , and
hem., 1368 (l94v).
Scand., 3, 487 (.1949),
fn QX
).
Acta. Che-.
Z. anorg.
llgem. Chem., 268,
Thesis Report
Niels C. Nielsen Hay 26, 1953
The preparation and s tudy of the structures and properties
of some metal derivatives of azo and azome thine dyes has led
to interesting results v± th respect to the theoretical aspects
of coordination chemistry and may well be of practical value to
the dye industry which has used these "lakes" for many years.
One aspect of this problem concerns the donor properties of the
azo group.
The earlier characterization of the metal lakes of azo dyes
as Werner-type coordination compounds was accomplished by Morgan
and Drexv, and their associates and students, who prepared, analyzed,
and performed various substitution reactions on these compounds.
This work showed the necessity of having substituents such as -OH,
NH2, etc. in at least one position ortno to the azo group in order
to form stable complexes. The later workers also indicated that
the azo group was involved in the formation of only one coordinate
covalent bond.1
Caliis and Nielsen ,3 continued this investigation and, by
the use of magnetic susceptibility measurements, i^as able to
establish the configuration of a large number of complexes.
Further work by Liu" and the measurement of stability constants
by Snavely and Fernelius5 permits a rather complete evaluation
of the metal complexes of the azo dyei
■ k-> •
With respect to the azo group as a donor, Feigl6 has found
it possible to induce color changes in solutions of azo dye indi-
cators by the a ddition of Pd(CN)2. The dyes used did not con-
tain any ortho substituents, and the change, apparently, cannot
be produced by any other compound.
Since the azo group does contai \ enough atoms to furnish
two pairs of electrons, in addition to TT electrons from the
double bond, it may be possible to form two coordinate covalent
bonds with this group, in examining the absorption spectra of
various dyes and their metal complexes in neutral, acid, and
alkaline solution, some evidence has been found for the addition
of a proton to the azo group without completely rupturing the
coordinate bond previously established with a metal ion. The
resulting spectra are also being evaluated with respect to the
affect of chemical constitution on the color of these compounds.
In an effort to induce the azo group to form more than one
bond, the following dyes are being prepared:
OH ^OH HO
N/ , > N = N < , > N = N <
\ nu un - — ■ — ' n
t)H ""~ OH HO OH HO
A study of the metal complexes of these I ~es should provide some
clarification of the donor properties of the azo group.
;araman, "The Chemistry of the Synthetic Dyes,", Aca-
ss, New York, 1952, Chap ter_^IV, pp% 551-569.
is. N.C. Nielsen, and J\C. Bailar, Jr * , J. Am. Ohem,
1. K. Venkataraman, "The Chemistry of the S^
demic Prer-
2. C.F. Calll-
Soc. , 74, 3461 (1952) .
3. J.C. Bailar, Jr. and C.F. Caliis, ibid, 74. 6018 (1952).
4. J.C .-I Liu, Ph.D. Thesis, University orTilinois, 1951.
5. F.A. Snavely and W.C. Fernelius, Science, 117. 15 (19o3). v
6. F. Feigl and G.B. Heisig J. Am. Chem. Soc.T~73, 5633 (1Q51)
-168-
:hesis report
The Synthesis of Hydrazine
Roger W. Sanftner May 26, 1953
Interest in an economic synthesis of hydrazine has
been heightened of late by the use of hydrazine as a speci-
ality fuel. Furthermore many uses, e.g. in polymers or ag-
ricultural chemicals, for substantial quantities of hydrazine
are conceivable should a low cost procedure for manufacture
of hydrazine be developed. Many approaches have been de-
vised, but the Raschig: synthesis (l), modified but slightly,
still remains the principal technical process in operation
in this country today.
Recently Mattair .and Sisler (2) reported the isolation
of hydrazine from the interaction of gaseous chlorine v/ith
liquid and gaseous ammonia. They stated that the yield of
hydrazine is dependent upon the ammonia to chlorine ratio
and upon the ammonium chloride concentration. Wiberg and
Schmidt (3) report that chloramine and ammonia dissolved in
anhydrous ether give no hydrazine whereas, with water present,
hydrazine was reported to have been formed.
A study has been undertaken of the reaction of ether-
eal solutions of chloramine with liquid ammonia under vari-
ous conditions in order to determine the effect of ammonium
chloride on the ammonolytic reaction of cnloramine and liqiid
ammonia. The reaction between hydrazine and chloramine in
liquid ammonia was also investigated.
McElroy (4) and Oldham (5) have found that chloramine
will undergo an one electron reduction at. .the dropping mer-
cury electrode in liquid ammonia. Hydrazine could be pos-
tulated as one of the products of such a reaction. Since
chloramine undergoes such a reduction, it was suggested
that chemical reduction of chloramine might result in the
formation of hydrazine. The reaction of aqueous chlora-
mine solutions with a variety of amalgams was investigated;
it was found that under certain conditions small quantities
of hydrazine are formed.
In Raschig* s original work, a variety of matelals was
investigated in an attempt to find a substance which would
act as a catalyst for the reaction between hypochlorite and
ammonia. It was found that glue and gelatin are effective
in increasing the yields of hydrazine; it was first assumed
that the increased viscosity of the solution served to in-
-159-
crease the yield. Subsequent investigations by Bodenstein
(6,7) and Mueller (8) have shown this theory to be incorrect,
These investigations have shown that the presence of small
amounts of metallic ions catalize the decompositon of hydra-
zine during the course of the reaction; glue and gelatin
serve as metallic deactivators. Many substances have been
investigated as substitutes for glue and gelatin (1,9. 10).
These two materials still remain, however, the most satis-
factory substances from the standpoint of yield. None the
less many disadvantages arise from their use. An investi-
gation was begun to determine the role of gelatin in its
removal of metallic ions from solution and also to discover
an effective agent which would not posses the disadvantages
exhibited by gelatin. Many substances have been investi-
gated and several promising alternative materials have been
found. Furthermore, it is believed that the cause of ef-
fectiveness of gelatin has been found.
1. Raschig, "schwefel and Stickstof f studien" , Verlag. chemie
G-. m.b.h., Leipzig, Berlin (1924).
2. Mattair and sisler, J. Am Chem. Soc. 73., 1619 (1951).
3. Wiberg and Schmidt, 3. Katurforsch. 6b, 336 (1951).
4. IvicElroy, Investigate Methods for the Synthesis of Hydra-
zine DA-11-022-0RD 828, (19 52).
5. Oldham, Private Communication (19 53).
6. Bodenstein, 2. physik. Chem; 137, 131 (1928).
7. Bodenstein, g. physik. Chem., 139, 397 (1928).
8. Miller, Kgl. Danske Videnskab. Selskab. iviath-fys. IvCedd.
12 (No. 16), 1 (1934)
9. Joyner, J. Chem. Soc. 1923, 1114.
10. Ffeiffer and Simons, Ber. 80, 127 (1927).
-170-
v
P.A,H0RRIGAN May 26,1953
A Study of the Uncoordinated Carboxyl Groups
in Glutamic Acid Complexes
There have been reports that unooordii:" ted functional grpups
on ligands react one way when the ligand is coordinated to a
metal atom, and another when the ligand is "free". '
It was decided to prepare complexes of dobalt and copper with
glutamic acid and study the properties of the uncoordinated car-
boxyl groups. The following compounds were prepared:
\
;CH-CH2-CH2-C0CHj
0 3 -!
-: Glut
Let
O-C-CH-CH2-CHp-C00
Na,[co(Glut)Jl
Ac;ifCo(Glut)il
Co^'fCo(GlutKH
,0 DJz 0 \
O-C-CH-CHp-CHp-G \»2I.
*-im2 * 6c2h5
- — NH2
NapfCu(Glut)Pl
C u TC u ( Glut )pf
L' r -J
"3
p-irHp- \j — \j — Crip— '^~y
/ 0 0 0 \
Co 40G-CH -CHo-GHo-d-O-CHo-O-^Al
/ 0 - 0 0 \
u40-d-CH-CHo-CH2-d-0-CHp-c(-^""VBrl
As an extension of this study, the application of the "masking
effect" is being attempted ,fith some of t_9se complexes. The
"masking effect'^ first illustrated by A.C .Kurtz-5>\ who ore-oared
amino derivatives of such complexes as / 0
[ / u \
CujO-Cl-CH-CHp-CHp-CHp-NH2
V^^-iiHp l-p
l2
and,vmo then removed the metal with HpS. Thus, one reactive end of
the molecule was "masked" while the other was left free to form a
derivative. The phenacyl derivatives mentioned above are being * :■■
treated to yield the metal-free ligands.
This work has shown that the uncoordinated carboxjl groups in
cobalt and copper glutamate complexes exhibit no anomalous propers
ties .
1. V.I.Xuznetsov,Zhur.Obshchei Khim(J. Gen Chem) 20,807-15(1950)
2. R.N.Keller and L.J.Edwards, J.Amer. Chem. Soc . ,74,215-19(1952)
3. A. C.Furtz, J. Biol. Chem. ,122,477 (1937-8)
4. A. C.Kurtz, J. Biol. Chem. ,180, 1253-67(1949)
THE LITERATURE 0? INORGANIC CHEMISTRY
i» General Reference Work s
1. Abegg: "Handbuch der anorganischen Cheraie" (1905-1939).
2. Chemical Society (London): ,fAnnual Reports" (19C4- ).
3. Friend: "A Textbook of Inorganic Chemistry" (1914-1937).
4. Gmelin: "Handbuch der anorganischen Chemie" (1924- ).
5. Gmelin-Kraut: "Handbuch der anorganischen Chemie" (19CS-1932)
6. Kirk and Othmer: "Encyclopedia of Chemical Technology" (1947- )
7. Klenm (ed.): "FIAT Reviews of German Science. Inorganic
Chemistry" (1948J .
8. teller: ''A Comprehensive Treatise on Inorganic and theore-
tical Chemistry" (1922-1937).
9. Moissan-Pascal: "Traite de Chimie Hinerale" (iew Edition).
10. 1'horpe's "Dictionary of Applied Chemistry", 4th Edition(1938-)
II . Texts and General Reference Texts
1. Emeleus and Anderson: "Modern Asoects of Inorganic Chemistry"
(1952).
2. Ephraim' s "Inorganic hemistry" , 5th English Edition by Thorne
and Roberts (1948).
3. Freshman texts: Of the many available, those by the following
authors are suggested: Hopkins and Bailar (1951); Sneed
and Maynard (1942) Pauling (1947, 1952); Laubengayer (1949) ;
Sisler, Vander Werf , and -Davidson (1949); ^ichardson and
Scarlett (1951); Labor and Lehrman (1949); Felsing and Watt
(1951) ; Holmes (1951) .
4. Hildebrand and Powell: "Principles of Chemistry", 6th Edition
(1952) and Latimer and Hildebrand: "Reference Book of Inor-
ganic Chemistry", 3rd. Edition (1952). Combined Volume.
5. Huckel: "Structural Chemistry of Inorganic Compounds" Vol.1
(1950) , V0i. 11(1951) .
6. Hofmann-HUidorff : 0Anorganiscne Chemie" (1951)
7. rlaxted: "Modern Advances in Inorga ic Chemistry" (1947).
8. Morgan and Burs tall: "Inorganic Chemistry, & Survey of
Modern Developments" (1936).
9. Moeller: "Inorganic Chemistry" (1952),
10. Partington: "A Textbook of Inorganic °hemistry" (1950).
11. Partington: "General and Inorganic Chemistry" 2nd Edition 0-951)
12. Piemy: "Lehrbuch der anorganischen Chemie", 3 Auflage (1940,
1949 ) .
13. Sidgwick: "The Chemical Elements and Their Compounds" (1950).
14." Wells: "Structural Inorganic Chemistry", 2nd Edition (1951).
15. Wiberg: "Anorganische Chemie" (1951).
16. Yost and Russell: "Systematic Inorganic Chemistry" (1944).
Ill . Inorganic Preparations and Laboratory Practice
1. Booth (Ed.): "Inorganic Syntheses", Vol. I (1939).
2. Fernelius (Ed.): "Inorganic Syntheses11, vol. II (1946).
3. -^udrieth (Ed.): "Inorganic Syntheses1'. Vol. Ill (1950).
4. Bailar (Ed.): "Inorganic Syntheses", ol. IV (1953).
5. Archibald: " ±he Preparation of Pure Inorganic Substances" 093$
-172-
6. Biltz and Biltz (trans, by Hall and Blanchard): "Laboratory
Methods of Inorganic °hemistry" , 2nd Edition (1923).
7. Blanchard, Phelan, and jDavis: "Synthetic Inorganic Ghemistry",
5th Edition (1932) .
8. Brauer: "Handbuch der praparativen anorganischen chemie" ,
Vol. 1 (Pts. 1-6) (1951-2)
9. Grubitsch: "Anorganisch-pra"parative Chemie" (1950).
10. Henderson and Pernelius: "•& Course in Inorgaic Preparations"
(1935).
11. King: "Inorganic Preparations. A systematic uourse of Experi-
ments" (1936).
12. Noves and Bray: "Qualitative Analysis for the Rare Elements"
(1927) .
13. Sanderson: "Vacuum Manipulation of Volatile Compounds" (1948).
14. Vanino: "Handbuch der praparativen Chemie", 3rd Edition,
Vol. II (1925).
15. Walton: "Inorganic Preparations" (1948).
IV . Texts of General Interest and Application
1. Bethe: "Elementary Nuclear Theory" (1947).
2. Priedlander and Kennedy: "Introduction to Kadiochemistry" (1943)
3. Glasstone: "Sourcebook on Atomic Energy" (1950).
4. Hein: °hemische Koordination-lehre (1950)
5. Herzberg: "Atomic Spectra and Atomic Structure" (1944).
6. Latimer: "Oxidation Potentials" 2nd Edition (1952).
7. Lewis: "Valence and the Structure of Atoms and Molecules "(19 23)
8. Palmer: "Valency, Classical and Modern" (1944).
9. Pauling: "The Nature of the Chemical Bond". 2nd Edition (1948)
10. Pollard and Davidson: "applied Nuclear Phvsics" 2nd Edition
(1951)
11. Rice: "Electronic Structure and Chemical Binding" (1940) .
12. Rice and Teller: "The Structure of Matter" (1949).
13. Selwood: "Magnetochemis try" (1943 )
14. Sidgwick: "Electronic Theory of Valence" (1927).
15. Sidgwick: "The Covalent Link in Chemistry" (1933).
16. Smyth: "Atomic Energy" (1945).
17. Stillwell: "Crystal chemistry" (1938).
18. Stranathan: " The 'Particles' of Modern Physics" (1942).
19. Syrkin and Dyatkina: "Structure of molecules and Chemical
Bond" (1950) .
20. Van Arkel: "Molecules and Crystals" (1949).
21. Van Arkel (Ed.): "Reine Metalle" (1939).
22. Various texts on physical chemistry, especially those by
Daniels (1948); Prutton and naron (1951); Gucker and Mel-
drum (1944) .
23. Wahl and Bonner (Ed.): "Radioactivity Applied to Chemistry"
(1951) .
24. Wells: "Structural Inorganic uhemistry", 2nd Edition(1951) .
V . Volumes Dealing with Special Phases of Inorganic Chemistry
1. Audrieth and Kleinberg: "Non-aqueous Solvents" (1953).
2. ^udrieth and Ogg: "Chemistry of Hydrazine" (1951).
3. Bell: "acids and Bases" (1952).
-173-
4. Booth aid Martin "Boron Trifluoride and its derivatives" (1949).
5. .Davidson et al: "More Acids and ^ases" (1944).
6. Faith, Keyes, and Clark: "industrial Chemicals" (1950).
7. Fricke and Hllttig: "Hydroxyde und Oxyhydrate" (1937).
8. Furnas: "Rogers' s Manual of Industrial Chemistry" (1942).
9. Hall, et al: "Acids and Bases" (1941).
10. Haszeldine and Sharpe "Fluorine and its Compounds" (1951) .
11. Hopkins: "Chapters in the Wiemistry of the Less Familiar Ele-
ments" (1939).
12. Hurd: "Chemistry of Hydrides" (1952).
13. Jander: "^ie Chemie in wasserahlnichen LBsungsmitteln" (1949)
14. Katz and Eabinowitoh: "Uranium" (1950)
15. Kleinberg "Unfamiliar Oxidation States and their Stabiliza-
tion" (1950)
16. Kobe: "Inorganic Process Industries" (1948).
17. Luder and Zuffanti: "Electronic Theory of Acids and ^ases"
(1946).
18. Mar tell and Calvin: ,,L;hemis try of the Metal Chelate Compounds'
(1952)
19. Rankama and Salama: "Geochemistry" (1950).
20. Head: "Industrial Chemistry" (1943).
21. Riegel "Industrial Chemistry" (1942).
22. Hochow: "Chemistry of the .Silicones" (1946)
23. Schwarz-Bass: "Inorganic uomplex ompounds'1 (1923).
24. Seaborg, Katz, and Manning: "The Transuranium Elements" (1949).
25. Simons: "Fluorine Chemistry" (1950)
26. von Hevesy: "Die seltenen Erden vom Standpunkte des Atom-
baues" (1927) .
27. Walden: "Salts, Acids, and Bases" (1929).
28. winnacker-V/eingaertner: "Chemische Technologie" , Vol. I,
II (1950).
29. Yost, Russell, and Oarner: "The Hare Earth Elements and
Their Compounds" (1947).
VI . Journals
1. Abstracting:
Chemical Abstracts (C.A.);
British Chemical Abstracts (B.C,A»);
Chemishes .Zentralblatt (Chem. Zentr.)
2. Original Research:
Journal of the American Chemical Society (J. Am. Chem. Soc)
Journal of Physical Chemistry (J. Phys . Chem.);
Analytical ^hemistry (^nal. Chem.);
Industrial and Engineering ^hemistry (Ind. Eng. Ghem.);
Journal of the Electrochemical Society (J. Eiectrochem.
Soc.) ;
Journal of the uhemical Society (J. Chem. Soc);
Nature (Mature) ;
Science (Science);
Transactions of the Faraday Society (Trans. Faraday Soc);
Proceedings of the uoyal Society of London (Trans. Hoyal Soc
London) ;
Journal of Chemical Physics (J. Chem. Phys.);
-174-
Zeitschrift fur anorganische Chemie (Z. anorg. Chem. or
Z. anorg. ellgem. Chem.);
Angewandte Chemie (-^ngew. Chem.):
Chemische ^erichte (^hem, Ber. or Ber.);
Comptes rendus (Compt. rend.);
Helvetica Chimica Acta (Helv. Chim. Acta)
3. Keview:
Chemical Reviews (Chem. Revs.);
Annual Reports (Ann. Reports) ;
Quarterly Reviews (Quart, Revs.);
Journal of Chemical Education (J. ^hem. nlduc).
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