PERCHLORIC ACID AND PERCHLORATEI

by

ALFRED A. SCHILT
Department of Chemistry
Northern Illinois University
DeKalb, Illinois 60115
 Published by
THE G. FREDERICK SMITH CHEMICAL COMPANY
867 McKinley Avenue
Columbus, Ohio 43223

LIBRARY OF CONGRESS
CATALOG CARD NUMBER: 79-63068

© COPYRIGHT 1979
BY
THE G. FREDERICK SMITH CHEMICAL COMPANY
 Dedicated to the memory of

G. FREDERICK SMITH

An inspiring mentor, zealous researcher, delightful raconteur,

and faithful friend.
 ABOUT THE AUTHOR

Alfred A. Schilt is Professor of Chemistry at Northern

Illinois University where he has taught and conducted re-
search in analytical chemistry since 1962. He completed both
his undergraduate studies and a masters degree at the Uni-
versity of Colorado before joining the Eastman Kodak Com-
pany as an analytical chemist in 1951. Two years later he
resumed graduate studies at the University of Illinois where,
under the enthusiastic direction of Professor G. Frederick
Smith, he developed strong abiding interests in perchloric
acid and perchlorate chemistry, spectrophotometric reagents
for trace metal determinations, and applications of coordin-
ation chemistry in chemical analysis. After receiving his doc-
torate in 1956 he taught at the University of Michigan until
1962 when he joined the faculty at Northern Illinois Univer-
sity. In addition to numerous research articles, he is the
author of "Analytical Applications of 1,10-Phenanthroline
and Related Compounds" published by Pergamon Press in
1969 and co-author of "The Copper Reagents: Cuproine, Neo-
cuproine, Bathocuproine" published by The G. Frederick
Smith Chemical Company in 1972.
 PREFACE

This monograph is intended to serve as a ready, single

source of useful information on the chemistry and applica-
tions of perchloric acid and perchlorate compounds. It
should also prove helpful to those who would contribute
further to the knowledge and utilization of these versatile
and intriguing compounds. To this end a thorough search
has been made of the literature up to and including publi-
cations in 1977. Most, if not all, of the pertinent references
are cited. This approach, coupled with a desire for concise-
ness, has precluded an in-depth treatment of many of the
topics. Many papers related to the study of the compounds
for use in explosives and propellant mixtures have been
omitted. Also omitted for the most part are references to re-
search articles that employed or dealt only marginally with
perchloric acid or its salts.
Over the more than 160 years since its inception an aura
of fear highlighted by general ignorance or lack of reliable
information has surrounded perchloric acid chemistry.
Thanks to the pioneering efforts and energies of scientists
such as G. Frederick Smith, H. H. Willard, and their many
devoted students, perchloric acid and its compounds are
now widely and effectively employed. Fear and ignorance
have given way to respect and understanding, with due re-
gard for appropriate safety measures. It is sincerely hoped
that this monograph will help disseminate the knowledge
and safeguards to enable others to put perchloric acid and
its derivatives safely and effectively to work in solving their
particular problems.

ALFRED A. SCHILT
 CONTENTS

PREFACE iv

I. INTRODUCTION 1
Historical Highlights 1
Question of Natural Occurence 3
Structure and Properties of Perchlorate Ion 4
References 7

II. PREPARATION AND PROPERTIES OF

PERCHLORIC ACID 9
Aqueous Perchloric Acid 9
Preparation 9
Purification 11
Physical Properties 11
Chemical Properties 21
Hydrates of Perchloric Acid 22
Anhydrous Perchloric Acid 24
Preparation and Properties 24
Reactions 26
Perchloric Anhydride 26
References 28

III. PROPERTIES AND PREPARATION OF

PERCHLORATES 31
Representative (Sub Group A] Metal Perchlorates 31
Ammonium and Alkali Metal Perchlorates . . . 31
Alkaline Earth Perchlorates 38
Group IIIA-VA Metal Perchlorates 41
Transition Metal Perchlorates 43
Group IB and IIB Metal Perchlorates 47
Group IIIB and Rare Earth Metal Perchlorates 48
Group IVB-VIIB Metal Perchlorates 49
Group VIIIB Metal Perchlorates 50

vi
 Miscellaneous Perchlorates 51
Nitronium Perchlorate 51
Nitrosyl Perchlorate 51
Hydrazine Perchlorate 52
Phosphonium Perchlorate 52
Selenious Acid Salt 52
Perchloryl Fluoride 52
Halogen Perchlorates 53
Organic Perchlorates 54
References 56

IV. APPLICATIONS OF PERCHLORIC ACID AND

PERCHLORATES IN CHEMICAL ANALYSIS . . . . 64
Dissolution and Oxidation of Inorganic Samples .. 64
Soils, Clays, and Silicates 64
Metals, Alloys, and Ores 67
Miscellaneous 68
Wet Oxidation of Organic Matter 69
Efficiency and Products of Wet Oxidations .. 71
Procedures and Special Techniques 74
Determination of Nonmetals 77
Determination of Trace Metals 80
Titrimetric Reagents 81
Precipitation and Extration Reagents 83
Deproteinization Agent 84
Drying Agents 85
Miscellaneous 87
References 90

V APPLICATIONS IN ORGANIC SYNTHESIS,

INDUSTRY, AND COMMERCE 96
Catalysts 96
Reactants in Organic Synthesis 99
Solvents 100

vii
 Explosives, Propellants, and Pyrotechnics 100
Electrolytes 102
Electropolishing 102
Voltaic Cells and Batteries 103
Miscellaneous 106
Animal Feed Additives 106
Miscellaneous 107
References 108

VI. CHEMICAL ANALYSIS OF PERCHLORATES . . . 114

Detection Methods 114
Separation Methods 115
Reduction of Perchlorates 118
Determination of Perchlorates 120
Gravimetric methods 121
Titrimetric methods 122
Titration of Perchloric Acid 124
Photometric Methods 125
Ion-Selective Electrodes 129
Miscellaneous Methods 132
Analysis of Perchlorate Compounds and Mixtures. 132
References 134

VII. BIOLOGICAL EFFECTS OF PERCHLORATES . . . 140

Animal Studies 140
Physiological Effects 140
Perchlorate Metabolism 142
Pharmacological Studies 143
Toxicology and Pathology 144
Bacterial and Micro-organism Studies 146
Plant Studies 147
References 150

viii
VIII. SAFETY AND ENVIRONMENTAL CONCERNS . . 153
Some Reported Perchlorate Explosions 153
Precautions in Use and Handling of Perchloric Acid 155
Precautions in Handling Perchlorates 157
Perchloric Acid Fume Hoods and Eradicates . . . . 158
Environmental Pollution 160
References 162

APPENDIX 164
Physical and Equilibrium Data for Binary, Ternary, and
Quarternary Systems of Perchlorates 164
References 175

INDEX 181

ix
 CHAPTER I
INTRODUCTION

Historical Highlights

The first preparation of a perchlorate compound was de-

scribed by Count Friederich von Stadion in a paper entitled
"Von den Verbindungen des Chlorine mit dem Sauerstoff",
published in 1816.l He obtained a new salt, which he called
"oxygenated potassium chlorate", as a water-insoluble resi-
due from a mixture of potassium chlorate and concentrated
sulfuric acid which had been allowed to stand for 24 hours
with frequent agitation. In the same paper he also described
obtaining "oxygenated chloric acid" by distilling a mixture
of the new salt with its own weight of sulfuric acid (diluted
with one-third its weight of water) to 140"', collecting the
white vapors that condense in the receiver as a liquid acid.
His "liquid acid" was undoubtedly an aqueous solution con-
taining approximately 709^ by weight perchloric acid. In a
subsequent 1818 paper,- Stadion described electrolytic meth-
ods for preparation of "oxygenated chloric acid" and "oxy-
genated potassium chlorate".
Important contributions were made in the early 1830's
by G. S. Serullas who prepared ammonium perchlorate, a
number of metal perchlorates, and a solid form of perchloric
acid (mistaken for the anhydrous acid but later identified as
the monohydrate) .3_5 Serullas also introduced the term "per-
chlorate", using it interchangeably with Stadion's term
"oxychlorate".
In the thirty-year period following the work reported by
Serullas, a number of significant papers appeared describing
new preparations of perchlorates and perchloric acid. These
included contributions by Berzelius in 1835,(i by Penny in
1840,7 and by Kolbe in 1862.8 The preparation of ethyl per-
chlorate, an extremely hazardous explosive, was reported in
1841 by Hare and Boye.9 Anhydrous perchloric acid, another
very explosive substance, was prepared in 1862 by Roscoe,10
who also correctly identified the solid acid obtained by
Serullas as the monohydrate."

1
 The explosive nature of certain perchlorates undoubtedly
discouraged many from undertaking further studies. Still
others were attracted by the possible commercialization of
perchlorates for use in pyrotechnics and explosives. For
example, Swedish patents were issued to O. F. Carlson in
189212 and 189713 relating to the manufacture of explosives
containing perchlorates. The first commercial perchlorate
plant was constructed in 1893 at Mansbo, Sweden. Produc-
tion problems were numerous and slow to be resolved, but
by 1904 the plant was producing ammonium perchlorate on
an efficient, regular basis. 14 Not long afterwards manufacture
of sodium, potassium, and ammonium perchlorates began in
other countries, notably France, Switzerland, Germany and
the United States. Production was pushed to maximum levels
during World Wars I and II. Instead of falling as it did after
World War I, production of ammonium perchlorate con-
tinued to grow after 1945, increasing phenomenally in the
1950's because of its usefulness in solid rocket propellants. 14
The remarkable commercial growth for perchlorates from
1910 to the present was accompanied or preceded quite
naturally by comparable growth in research. Numerous
patents were granted, and countless research publications
appeared. Research activity, as evidenced by rate of publi-
cation, continues strong to this day.
In reviewing the history of any endeavor the contribu-
tions of certain individuals can be seen clearly in retrospect
to have influenced significantly further progress and work
by others. In addition to recognizing the pioneering efforts
of Stadion, Serullas and Roscoe, any historical account of
perchloric acid and perchlorates should include at least men-
tion of the persons and contributions cited in the following
sentences. Isolation of the anhydride of perchloric acid,
CUO, was achieved by A. Michael and W. T. Conn in 1900,15
Extensive measurements of physical properties of perchloric
acid-water mixtures were reported in 1902 by H. J. van Wyk
leading to the identification of five separate hydrates. 16 In
1912, H. H. Willard of the University of Michigan described
a convenient and practical method for the preparation of
perchloric acid from ammonium perchlorate. 1 ' Assisted by
various graduate students, Professor Willard made a number
of significant contributions to perchlorate chemistry in the

2
 course of his long and remarkably productive career.18 Possi-
bly the most significant influence that Willard exerted on
the advancement of perchlorate chemistry was to interest
G. Frederick Smith, one of his first graduate students, to
undertake his doctoral research on the subject. From his
earliest publication with Willard in 1922 to his last pro-
fessional publications in 1966, and even up to his death in
1976, Smith never lost his curiosity and enthusiasm for per-
chlorate chemistry research. With over 60 research papers
and publications devoted to the subject, he had truly a life-
long love affair with perchlorate chemistry. In addition to
advancing the use of perchloric acid and perchlorates in
chemical analysis, Professor Smith carried out fundamental
studies, devised new preparation procedures, and founded
the chemical company that bears his name and has provided
perchloric acid and perchlorates to the scientific community
since 1928.,9
In the early 1930's Ernest Kahane of Paris, France pub-
lished a number of important papers describing the use of
perchloric acid for the total destruction of organic and bio-
logical matter prior to determination of mineral constituents.
Many investigations and applications followed after Kahane's
original publication. 20 More recent contributions to per-
chlorate chemistry include the preparation of perchloryl
fluoride by Bode and Klesper21 and the preparation of
chlorine and bromine perchlorates by Schack, Pilpovich, and
Wilson.22 These perchloric acid derivatives have proven use-
ful in organic synthesis.

Question of Natural Occurence

The discovery of small amounts of perchlorates in

natural deposits of nitrates in Chile was reported by Beckurts
in 1886.23 This was confirmed a number of years later by
various groups of investigators concerned with the possible
harmful effects of perchlorate in Chilean nitrate used as
plant food.24'27 In 1896 Sjollema reported finding potassium
perchlorate as a contaminate of Chile saltpeter in amounts
ranging from 0 to 6.79%, varying even within the same lot.25
Maschhaupt found a maximum perchlorate content of 1.5%
in crude and about l</c in refined saltpeter.27 Questions as to

3
 how the perchlorate was produced by natural means and
why perchlorate has been found naturally only in Chilean
saltpeter remain disturbingly unanswered. It is possible that
the answers could provide new insights into geological and
biological processes and even lead to a new method of prep-
aration of perchlorates.
Discovery of perchlorate in sea water was reported in
1958 by Baas-Becking, Haldane, and Izard,28 who found levels
of 10 to 1000 ppm CIO* in samples collected at various locali-
ties. Their findings have not been substantiated but instead
rather strongly refuted by others. An extensive study by
Greenhalgh and Riley29 indicated no detectable concentration
of perchlorate in any of over 30 surface and deep-water sam-
ples collected from both northern and southern hemispheres™
They attributed the earlier findings of Baas-Becking and co-
workers to chloride interference. Johannesson :tl found no
evidence of naturally occuring perchlorates in samples of
fresh-seawater taken 50 miles apart off the coast of Welling-
ton, New Zealand in 1960. At least two other groups have
reported the absence of detectable levels of perchlorate in
sea water samples, based upon highly sensitive analytical
procedures.:'"':l:! Thus it appears that perchlorate is not an
important natural constituent of sea water.

Structure and Properties of Perchlorate Ion

X-ray diffraction measurements on crystalline hydron-

ium perchlorate by Lee and Carpenter'"1 indicate that the
perchlorate ion has nearly perfect tetrahedral geometry and
an average Cl-0 bond distance of 1.42 A. In water its limit-
ing conductance is 67.32 ± 0.06.'iS A difference in molar re-
fractive power between perchlorate ion and chloride ion
of 6.66 ± 0.06 has been deducted from refractive index
measurements of their aqueous solutions.3" The standard
enthalpy of formation of the perchlorate ion in infinitely
dilute aqueous solution at 25 has been determined by two
different groups of investigators; their results —30.53 ± 0.2037
and —30.87 ± 0.07 kcal/mole:iM are in close agreement.
X-ray photoelectron spectra™ and fluorescence spectra4"
have been investigated to study the bonding and electronic
structure of the perchloric anion. The results indicate that the

4
 3d-orbitals of chlorine participate in the 3dir(Cl)-2pir(0) and
3dff(Cl)-2s(0) bonds.The relative contributions of the chlorine
3d-orbitals to the 6- and II- bond levels were evaluated by
comparing relative intensities in the L-spectrum of chlorine. 40
The hydration number of the perchlorate ion has an
average value of 4, according to Symons and Waddington 41
who based their determination upon Raman and infrared
spectra of aqueous solutions. They found a similar value for
the fluoroborate anion. It appears that each oxygen atom, or
fluoride atom in the case of fluoroborate, may be hydrogen
bonded by a separate water molecule. Hydration of per-
chlorate ions in nitrobenzene has been studied using cryo-
scopic measurements to estimate the value of the first hydra-
tion constant. 42 Solvation in methanolic solutions has also
been studied by low temperature infrared spectroscopy. 43
Raman spectra of various aqueous perchlorate solutions have
been interpreted and evaluated as a function of concentra-
tion to study the state of perchloric ions and their interactions
with metal ions.44
The effect of perchlorate ions on the hydrogen bonded
structure of water has been examined by infrared spec-
trometry.43'40 Strong evidence of the breaking of water struc-
ture by perchlorate ions was found with the observation that
the optical density of the 2100 cm"1 band (combination band
of H2O] decreases with increasing perchlorate ion concen-
tration, while that of the 3590 cm"1 band (fundamental band
of H2OI increases. No detectable absorption was observed in
the 4000-600 cm'1 region indicative of hydrogen bonding be-
tween water and perchlorate ion, and the fundamental band
due to perchlorate ion was unaltered. This suggests that the
interaction which causes the water structure to be altered is
purely ionic.45
Nuclear magnetic resonance studies of perchlorate solu-
tions have provided some insights concerning the interactions
of perchlorate ions with solvent and with metal cations.
Berman and Stengle47 found that free perchlorate ions in
solution have a longer relaxation time than in a contact ion-
pair with cations, as indicated by NMR line widths for
chlorine-35. They report that the tendency for perchlorate
ions to form contact ion-pairs is favored by high charge to
radius ratio of cation, low dielectric constant of solvent, and
low basic strength of solvent. Craighead and Bryant48 found
evidence of weak complexation between manganous and
perchlorate ions. Based upon chlorine-35 NMR relation time
measurements they estimated an exchange rate of 3 x 104
to 3.6 x 107 sec"1 for the perchlorate-manganous ion exchange
process. Quadrapole relaxation studies of perchlorate and
other tetrahedral ions in aqueous solution have been reported
also by Reimarsson and coworkers. 49 A proton magnetic reso-
nance study of the interaction of perchlorates with anhydrous
perchloric acid has provided evidence of hydrogen bonded
H(ClCu)2' anion formation.50
Although the basicity of the perchlorate ion is exceed-
ingly weak, it is capable of metal ion complexation in the
absence of excessive competition by solvent or other ligands.
Considerable evidence now supports this once widely doubted
fact."'1"'" Further details are provided in Chapter III.
Oxidation-reduction and other chemical properties of
the perchlorate ion are described together with those of the
acid and other perchlorate compounds in the chapters that
follow.

6
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2. F. von S t a d i o n , A n n . c h i m . <:l p h y s . . 8, 406 (1818).
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h o l d , N e w York, N.Y., 1960, C h a p t e r 1.
15. A. M i c h a e l a n d W. T. C o n n , A m . C h e m . /.. 23, 444 (1900).
16. H. J. v a n W y k , Z. u n o r g . Chem., 32. 115 (1902.
17. H. H. W i l l a r d , J. A m e r . C h e m . S o c , 34, 1480 (1912).
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b i o l o g i q u e ) , " Part I a n d II, P a r i s . F r a n c e . H e r m a n n a n d Gie. 1934.
21. H. Bode a n d E. K l e s p e r , Z. [more., u. ullgcm. C h e m . . 266. 275 (1951).
22. C. J. S c h a c k , D. P i l i p o v i c h , a n d R. D. W i l s o n , U.S. P a l e n l 3694172
(1972).
23. H. B e c k u r t s , A r c h . P h u r m . . 224, 333 (1886).
24. M a e r c k e r , Lumiiv. Vers.-Sin.. 52, 34 (1899).
25. B. S j o l l e m a , C h e m . Z l s . . 20. 1002 (1896).
26. A. V e r w e i j , C h e m Weefcljlull, 1, 155 (1904).
27. ], G. M a s c h h a u p t , D i r c k l i e vim lien I.iini/liDiiiv. 1914, 17.
28. L. G. M. Bass-Becking, A. D. H a l d a n e , a n d D. Izard, Niilm-e, 182.
645 (1958).
29. R. G r e e n h a l g h a n d J. P. Riley, Niilure, 187. 1107 (I960).
30. R. G r e e n h a l g h a n d J. P. Riley, j . M u r i n e Biol. A s s o c Unilei/
K i n g d o m , 4 1 , 1 7 5 (1961).
31. J. K. [ o h a n n e s s o n . Ann). Chem.. 34,1111 (1962).
32. K. W . L o a c h , N n t u r e . 196, 754 (1962).
33. I. I w a s a k i , S. U t s u m i , a n d C. Kang, Bull. C h e m . S o c Jnpim. 36, 325
(1963).
34. F. S. Lee a n d G. B. C a r p e n t e r , /. Pliys. C h e m . , 63, 279 (1959).
35. J. H. Jones, J. A m e r . C h e m . S o c , 67, 855 (1945).
36. A. M a z z u c c h e l l i a n d A. Verciilo, Cuzz. c h i m . inil.. 55, 498 (1925).
37. M. E. Efimov a n d V. A, M e d v e d e v , /. C h e m . T h c r m o i l y n . , 7. 719
(1975).
38. E. P. K i r p i c h e v , Y. I. R u b t s o v , N. V. K r i v t s o v , T. V. S o r o k i n a , a n d
G. B. M a n e l i s , Zh. Fiz. Khim., 49, 1975 (1975).

7
 39. R. P r i n s , C h e m . Phys. Lett., 1 9 , 3 5 5 (1973).
40. A. P. S a d o v s k i i , L. N. M a z a l o v , G. N. D o l e n k o , A. A. K r a s n o p e r o v a ,
a n d V. D. Y u m a t o v , Z h . Strukt. Khim.. 14, 1048 (1973).
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71, 22 (1975).
42. T. T a r u i , Bull. Chum, Soc. Jtipcm, 48, 295g (1975).
43. I. M. S t r a u s s a n d M.'C. R. S y m o n s , C h e m . P h y s . Left.. 39, 47i (1976),
44. D. N. G l e b o v s k i i , V. A. L a t y s h e v a , L. A. M y u n d , a n d B. P. T a r a s o v ,
Mai. Fiz. Biofiz. Vod. Sist, 1, 04 (1973).
45. T. T a k a m u r a , H. K i h a r a - M o r i s h i t a , a n d H. Y o s h i d a , N i p p n Kugaku
Kciishi[2j, 1974, 375.
46. D. W . J a m e s a n d R. F. A r m i s h a w , Inorg. N u c i . C h e m . Let!., 12, 425
(1976).
47. H. A. H e r m a n a n d T. R. S t e n g l e , J. Phys. C h e m . . 79, 1001 (1975).
48. K. L. C r a i g h e a d a n d R. G. B r y a n t , Mol. p h y s . . 2 9 , 1 7 8 1 [1975].
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P h y s . Chem.. 8 1 , 7 8 9 (1977).
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Nouk SSSH. Sor. Khin? (2), 265 (1973).
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8
 CHAPTER II
PREPARATION AND PROPERTIES OF
PERCHLORIC ACID
AQUEOUS PERCHLORIC ACID
Preparation
Perchloric acid was first synthesized by von Stadion1'"
in 1816 by vacuum distillation of a mixture of sulfuric acid
and potassium chlorate, and by electrolysis of a saturated
aqueous solution of chlorine dioxide. Numerous other meth-
ods have been devised since, however the most important
commercial ones are based upon either the Kreider-Mathers 3
or the Willard4 method. These and other methods, classified
as to type, are described below.
Electrosynthesis. Electrolytic oxidation of hydrochloric
acid to perchloric acid, as described by Walker 5 in 1918 and
in greater detail by Goodwin and Walker'1 in 1921, proved
relatively inefficient for commercial purposes. Production of
chlorine rather than perchloric acid is favored at higher con-
centrations of hydrochloric acid, so it is necessary to employ
very dilute (0.1-0.5 N) solutions and thus extensive evapora-
tion to concentrate the product. Some improvements were
forthcoming. Rakov and coworkers 7 found that (1] perchloric
acid formation begins at 2.4 V and reaches a maximum value
at 2.8-2.9 V, (2) lowering the temperature to —20° signifi-
cantly accelerates the process, and (3] concentration changes
of chloride ion from 0.5 to 1.8 N and of perchlorate ion from
3 to 8 N are without effect except for decreased current effi-
ciency at the highest concentration of hydrochloric and per-
chloric acid. In another study, Rakov et al.8 found that plati-
num anodes afford higher yields of perchloric acid than did
Pt-Ti or Pt-Ta bianodes. Rakov and Shimonis," in a patent
granted in 1976, claim increased current efficiency is attained
in the electrosynthesis of perchloric acid from 1 to 4 N hydro-
chloric acid by using an iridium anode at a potential of 2.9-3.3
V and a temperature at —5 to —30°,
Preparation by anodic oxidation of chlorine has been
investigated. 10 At platium electrodes a reaction apparently oc-
curs between adsorbed chlorine and active oxygen which
leads to formation of perchloric acid through an adsorbed in-
9
 termediate product.11 Lead dioxide anodes are especially
effective, particularly if sodium flouride is added to the
electrolyte.12
Electrolytic oxidation of chlorates was investigated by
Newman and Mathers. 13 Using a two-diaphragm, three-com-
partment cell with a platinum anode and iron cathode they
were able to obtain an aqueous solution of about 2 N per-
chloric acid prior to concentrating by evaporation.
Double-decomposition. A solution of perchloric acid
can be made by warming an aqueous solution of potassium
perchlorate with a small excess of hydrofluosilicic acid,
HiSiFs, for an hour, and filtering the cooled liquid.14'15 The
slight excess of HaSiFa can be removed by adding a little
barium perchlorate or chloride. Another preparation involves
decomposition of barium perchlorate with sulfuric acid to
yield insoluble barium sulfate and perchloric acid.16 Kreider17
prepared an aqueous solution of perchloric acid by treating
sodium perchlorate with an excess of concentrated hydro-
chloric acid, filtering off the sparingly soluble sodium
chloride, and heating to 135° to drive off the excess hydro-
chloric acid. Mathers 18 investigated the Kreider procedure
and provided additional procedural details. The Kreider-
Mathers method of preparation proved suitable for com-
mercial use and led to the Pernet process, 13 patented in 1946,
which provides for continuous operation and nearly auto-
matic control.
A mixture of potassium perchlorate and sulfuric acid
when distilled in a current of steam yields an aqueous solu-
tion of perchloric acid, contaminated by sulfuric acid, which
must be redistilled or treated with barium carbonate to re-
move sulfate.20 A much superior method, devised by Willard, 4
involves treatment of ammonium perchlorate with an excess
of a mixture of nitric and hydrochloric acid, giving rise to
the approximate over-all reaction as follows:
34 NH4C104 + 36 HNOa + 8 HC1 -» 36 HC104 + 4 Ck
+ 35 N2O + 73 H2O. Gaseous products and any excess acid
reactants are eliminated on heating to concentrate the per-
chloric acid to its constant boiling composition. Yields in
excess of 99%, approaching theoretical, were obtained by
Willard.
Chemical Oxidation. Ozone can be used to oxidize
hypochlorous 21 and chloric acid22 to perchloric acid. Patents
10
have been issued for the production of perchlorates and per-
chloric acid by ozone treatment of chlorates in H C l O A n d
of gaseous mixtures of water and chlorine or hydrogen chlor-
ide irradiated with 2537 A.24
Chloric acid undergoes self-oxidation, especially if con-
centrated, decomposing into perchloric and chlorous acids. 25
Serullas obtained perchloric acid in low yield from thermal
decomposition of chloric acid.26
Preparation procedures for perchloric acid and per-
chlorates by oxidation of chloric acid or chlorates on heating
with silver oxide27 or by oxidation with lead dioxide in
55-70% H2SO428 have been reported but appear to afford little
advantage.
Purification
Aqueous solutions of perchloric acid can be concentrated
by boiling at atmospheric pressure to 203°, at which point an
azeotropic solution is attained containing 72.4% HCIO4. For
purification by distillation it is necessary to employ reduced
pressures (below 200 mm] to avoid partial decomposition to
chlorine, chlorine oxides and oxygen. Distillation procedures
are described in papers by Mathers, 3 Willard, 4 and others.29"31
High-purity perchloric acid (70% by weight], greatly
diminished in trace metal content, can be conveniently pre-
pared by a sub-boiling distillation method developed at the
National Bureau of Standards. 32 A production rate of about
600 ml of perchloric acid per day was achieved using a pure
quartz sub-boiling still. Analysis revealed the presence of
16 ppb total impurity elements as opposed to 3400 ppb for the
ACS grade starting acid and 100 ppb for a commercial high-
purity acid. Using a similar appartus, but with a constant-
level feed-control device, other workers 33 decreased lead im-
purities in 70% perchloric acid to levels of 0.2-0.4 ppb in a
single pass.
An interesting method for removal of impurities from 1 M
perchloric acid for use as a supporting electrolyte has been
described, 34 based upon adsorptive and electroctive properties
of a column of platinum sponge held at a fixed potential.
Physical Properties
Extensive measurements have been made of the physical
properties of aqueous solutions of perchloric acid, particu-

11
 larly by H. J. Van Wyk and by L. H. Brickwedde. These are
summarized briefly with literature references in the follow-
ing paragraphs, and many of the results are compiled in
Tables 1 through 8,
Densities of aqueous solutions at various temperatures
have been reported by Van Wyk,35 Brickwedde, 38 Clark,3T van
Emster,38 Markham,39 and Smith.40 Tables 1 and 2 include
most of the results.
Freezing point data were determined and plotted versus
composition by Van Wyk35'41 and by Brickwedde 36 to identify
hydrates.
Vapor pressures and activity coefficients were measured
by Pearce and Nelson,42 Robinson and Baker,43 and Robinson
and Stockes.44 Results are compiled in Tables 3 and 4. Boiling
points of various aqueous compositions are listed in Table 5.
Viscosities are reported by Brickwedde, 36 Van Wyk,3-'1
Clark,37 and Simon.43 Table 6 includes the data.
Surface tension data of Neros and Eversole,46 given in
Table 7, indicate that a maximum occurs at a composition
corresponding to HCIO* • 3H2O.
Refractive index measurements are reported by McLean
and Pearson47 for perchloric acid solutions of 0 to 72 per
cent concentration at 20° and 30". Their data are compiled
in Table 8 and provide a convenient means of determining
concentration of aqueous perchloric acid solutions. A plot of
molar refraction against acid concentration yields a smooth
curve with a minimum at 18% HCIO4.48"49 From density and
refractive index data, a radius for the perchlorate ion of
1.82 AS1 and apparent molal volumes48"50 were calculated.
Electrical resistivity and conductance measurements
have been reported by Brickwedde 36 of solutions 10 to 70
weight per cent perchloric acid from —60 to +50° C, by
Usanovich and Sumarokova 52 for 0 to 100% HCIO4 at 50°,
and by Klochko and Kurbanov.53 Discontinuities in the
specific conductance curves corresponding to tri- and tetra-
hydrates and a maximum at 37 weight per cent HCIO4 were
found52'53. Viscosity and electrical conductance of perchloric
acid solutions from —50 to +90°C were reported recently
by Maksimova et ai. 5 ' Conductance, transport numbers, dif-
fusion coeffients, and related quantities have been reported
for aqueous perchloric acid at 25° C by Haase.55

12
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TABLE 2. SPECIFIC GRAVITY OF
PERCHLORIC ACID SOLUTIONS'
HCIO4 Specific Gravity
Mole % Weight % 2(TC 50'C

100 100 1.7676 1.7098

92.75 98.62 1.7817 1.7259
76.15 94.69 1.8059 1.7531
63.85 90.78 1.7690
50.00 84.80 1.7756
43.4 81.06 1.7619
35.67 75.57 1.7386 1.7023
27.96 68.41 1.6471 1.6110
21.44 60.37 1.5353 1.5007
15.46 50.50 1.4078 1.3779
10.56 39.72 1.2901 1.2649
6.23 27.05 1.1778 1.1574

14
TABLE 3. VAPOR PRESSURE OF
PERCHLORIC ACID AQUEOUS
SOLUTIONS AT 25°C4!
lality HC10< Vapor Pressure, mm
0.0 23.752
0.20064 23.593
0.60655 23.254
1.01589 22.870
3.1512 20.192
5.4347 16.308
7.8719 11.490
10.5139 6.838

15
 TABLE 4. MEAN ACTIVITY COEFFICIENTS
AND OSMOTIC COEFFICIENTS OF AQUEOUS
PERCHLORIC ACID SOLUTIONS AT 25 °Ca
HC104 Activity Coefficient O s m o t i c Coe
Molarity f* fo
0.2 0.766 0.951
0.4 0.754 0.966
0.6 0.763 0.988
0.8 0.784 1.013

1.0 0.810 1.041

2.0 1.039 1.210
3.0 1.420 1.406
4.0 2.018 1.622

5.0 2.94 1.860

6.0 4.41 2.106
7.0 6.90
8.0 11.32

9.0 18.0
10.0 29.0
11.0 47.3
12.0 77.7

1 16
TABLE 5. COMPOSITION OF LIQUID AND
VAPOR PHASES OF AQUEOUS PERCHLORIC
ACID AT DIFFERENT BOILING POINTS
Boiling point Pressure Weight per Cent HCIO4
'C mm Liquid Vapor

203 760 72.40 72.40

198.7 760 70.06 40.11
181.2 760 65.20 6.06
162.3 760 61.2 0.90
148.0 760 56.65
114.8 760 38.90
107 18 70.5
92 18 79.8
70 18 84.8
35 18 92.0
24.8 18 94.8
16.0 18 100
Sfi 100
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 TABLE 7. SURFACE TENSION OF
PERCHLORIC ACID SOLUTIONS46
Surface Tension at Various
HC104 Temperatures (°C)
Weight % 15° 25° 50"
0.00 73.51 71.97 68.16
4.86 72.52 71.18 67.60
10.01 71.66 70.34 66.97
20.38 70.46 69.21 66.12
30.36 69.82 68.57 65.66
40.37 69.72 68.49 65.74
53.74 70.33 69.02 66.60
60.70 70.88 69.69 67.40
63.47 70.77 69.73 67.44
67.59 70.67 69.71 67.41
70.43 70.07 69.54 67.26
72.25 69.96 69.01 66.85

19
TABLE 8. REFRACTIVE INDEX OF AQUEOUS
PERCHLORIC ACID47
HCICU Refractive I n d e x
Weight % 20'C 30'C
0.00 1.3330 1.3320
9.73 1.3395 1.3381
20.05 1.3470 1.3452
30.75 1.3580 1.3559
40.31 1.3680 1.3665
50.28 1.3813 1.3800
60.08 1.3983 1.3960
62.81 1.4034 1.4010
64.00 1.4054 1.4028
68.52 1.4129 1.4103
70.06 1.4151 1.4130
72.62 1.4190 1.4159

20
 Raman spectra156'57 and proton magnetic resonance spec-
tra 58 have been recorded to study the degree of dissociation of
perchloric acid in aqueous solutions.
Apparent molal expansibility, volume, and compressi-
bility of aqueous solutions of perchloric acid at 25° were
determined recently59 to study their dependence on com-
position and the nature of anion interactions with water.

Chemical Properties
The chemistry of perchloric acid in aqueous solutions
can be characterized briefly as that of an exceptionally strong
acid with essentially no oxidizing strength except towards
active metals that normally displace hydrogen from acids.
However, if the perchloric acid be both hot and concentrated
its oxidizing properties are considerable, and in certain cases
dangerously vigorous. It is also a strong dehydrating agent
when hot and concentrated. Under such conditions it serves
as an highly efficient oxidant for destruction of organic mat-
ter. When dilute or cold, however, its oxidizing strength is so
greatly diminished that it is suitable for use without inter-
ference in the study of many redox reactions. Likewise, the
very weak tendency for perchlorate ions to coordinate to
metal ions favors the use of perchloric acid as a non-inter-
fering acid in complexation studies.
Perchloric acid is one of the strongest acids known. Its
effective strength depends, like any other acid, on the basicity
of the solvent in which it is dissolved. In water it is com-
pletely ionized in the most concentrated solutions and com-
pletely dissociated up to a concentration of about 4 M, where
the first appearance of undissociated HClCU has been de-
tected.60 In concentrated sulfuric acid it is strongly disso-
ciated but weaker than disulfuric acid (H2S2O7).61 In organic
solvents it behaves as a strong acid, highly ionized but not
necessarily highly dissociated, depending on dielectric and
solvation strengths of the solvent. As a strong acid it is
especially useful in measuring pK values and relative
strengths of very weak bases. Hammett acidity functions
for this purpose have been measured by various workers, in-
cluding Gillespie,62 Yates et al.,63'64 and Attiga and Rochester.65
Values of the acidity function, Ho, for perchloric acid in acetic

21
acid-water mixtures have been determined by Wiberg and
Evans.66
Certain strong reductants are oxidized by perchloric
acid at room temperature. Increasing the concentration of
perchloric acid favors its reaction rate as an oxidant, sug-
gesting that the protonated form rather than the perchlorate
ion is the active species.67 Some of the reductants investi-
gated in reaction rate studies include ruthenium(II), 68 titan-
ium(III),69 molybdenum(III), 70 chromiumfll], 70 titanous chlo-
ride or sulfate,17 platinum cathode, 72 hydrogen and simple
organic fuels,73 and various metals. 74 Reductants for perchlo-
rates are further discussed in Chapter VII.
Perchloric acid solutions are decomposable thermally or
by exposure to ionizing radiation. Studies of its thermal de-
composition indicate an activation energy for decomposition
of 22.56 kcal/mol 75 , an autocatalytic effect for 65-100 wt %
HCIO4,78 and heterogeneous catalysis by oxides of iron, cop-
per or chromium.77 Radiolysis studies reveal that perchloric
acid decomposes only as a result of direct action of the radia-
tion, yielding chlorate ions as major product, and that it
does not react with products from the radiolysis of water.78'80

HYDRATES OF PERCHLORIC ACID

A total of five different hydrate compositions have been
identified from studies of liquid-solid equilibriums in the
binary system perchloric acid-water. Several different in-
vestigators confirm these findings. Van Wyk,35'41 the first
to investigate the system, found the following hydrates:
HC104 • HsO, HCIO4' 2H2O,, HC104 • 2.5H2O, a- and 0-forms
of HC104 * 3HsO, and HCIO4 * 3.5HaO. Brickwedde 36 obtained
a freezing point-composition plot very similar to that of
Van Wyk's but with different, presumably more accurate,
temperatures. Zinov'ev and Babaeva81 confirmed the existence
of all five hydrates, noting also that the mono- and dihydrate
exist in aqueous solution up to a temperature of 75° and
that the other three hydrates are much more highly dissoci-
ated. Mascherpa and coworkers, 8a on repeating the study of
Van Wyk over the concentration range of 58-74 wt. per cent
HC104, found hydrates of 1 molecule of HC1CU with 2, 2.5,
3, and 3.5 molecules of HaO, exhibiting melting points re-

22
 spectively of —20.65°, —32.1°, —40.2°, and —45.6° (all
± 0.1°). A sixth hydrate composition, corresponding to
H2O • 4HC104, has been reported by Mascherpa, 84 based upon
a thermal analysis study.
Perchloric acid hydrates involve considerable hydrogen
bonding. A bond-valence analysis of five of the hydrates by
Brown84 shows that normal H-bonds account for only about
one-half of the bonding of the perchlorate ion. Each H-atom
also forms an average of four additional weak interactions.
The infrared spectrum of crystalline perchloric acid tri-
hydrate is consistent with Hs0+ central pattern complexed by
two water molecules through short asymmetric hydrogen
bonds. 85 Perchlorate ions participate to some extent in hydro-
gen bonding in the hydrate HCIO4 * 3.5 H2O according to
Almlof.8* For this hydrate, X-ray diffraction data86 at —188°
indicates two independent H i O r complexes having one water
molecule in common. The two perchlorate ions show only
minor deviations from tetrahedral symmetry, with mean Cl-0
distances of 1.437 and 1.443 A.
The most extensively studied hydrate is the monohydrate.
It exists entirely as the oxonium salt H.iO+ClO-4 and thus can
be correctly referred to as oxonium perchlorate. It contains
84.78 per cent HCIO4 by weight, forms long needle-like crys-
tals, melts at 49.905° with considerable expansion, and ex-
hibits a heat of fusion of 2.46 ± 0.08 kcal/mole. Methods for
its preparation, described by Roscoe87 and by Klages,88 in-
volve preparation of the anhydrous acid and treatment with
an appropriate amount of either water or perchloric acid
dihydrate. Oxonium perchlorate dihydrate can be safely
stored at room temperature indefinitely without decompo-
sition if protected from reactive substances.
The crystal structure of oxonium perchlorate has been
determined by X-ray diffraction,89 A reversible phase change
occurs —23.4° accompanied by an increase in density.90 Be-
low this temperature, at —80° Nordman 91 reports that the
crystal structure belongs to the monoclinic space group and
consists of H-bonded layers of perchlorate and hydronium
ions. Above the transition point the structure is orthorhombic
and more disordered. The perchlorate ions are nearly per-
fect tetrahedra with an average Cl-0 distance of 1.42 A.
Proton magnetic resonace spectra"2 indicate that the oxonium

23
ion has a pyramidal structure. The Raman spectrum 93 is con-
sistent with the X-ray and NMR findings, and valence force
constants have been calculated.
Dioxonium perchlorate (perchloric acid dihydrate, 73.60
wt. % HCIO4) boils at 203 = at atmospheric pressure, is hy-
groscopic, fumes in moist air, and freezes at —17.8°. Recom-
mended as a primary standard for acidimetry, 40 it can be
prepared by distilling 70-72% HCIO4 at 2-7 mm and discard-
ing the first-half or more of the distillate before collecting
the dioxonium perchlorate.

ANHYDROUS PERCHLORIC ACID

Preparation and Properties
The preparation of anhydrous perchloric acid can be
readily achieved by fractional distillation of a mixture of
concentrated sulfuric acid and 65-70% perchloric acid.94'95
Magnesium perchlorate,"" phosphorous pentoxide, 97 and sul-
fur trioxide97 have also been described as dehydration agents
in place of sulfuric acid. Anhydrous perchloric acid can also
be obtained by extration into methylene chloride from a
mixture of one part by volume of 70% perchloric acid and
four parts 25% fuming sulfuric acid.88 It has been recom-
mended that only small amounts of reagents be used in this
procedure and that the extraction flask be kept cold because
of the risk of an explosion.98 Procedures for the preparation
of anydrous deuterium perchlorate and DaOClO-i have been
described by Smith and Diehl.™
The anhydrous acid is colorless, hygroscopic, volatile,
extremely reactive, and explosively unstable. In contact
with skin it produces serious and painful wounds. On addi-
tion to water it generates considerable heat and a hissing
noise. It freezes at —112 and boils at 18 mm at 16° without
decomposition. It can not be distilled at ordinary pressures
without decomposition and explodes at about 90°. On stand-
ing at ordinary temperatures the pure acid gradually yellows
and eventually (within 10 to 30 days) explodes spontane-
ously, Storage time can be extended up to 60 days by use
of liquid-air temperatures without formation of colored de-
composition products. Extreme caution is necessary to avoid
contacting the anhydrous acid with wood, paper, or other

24
combustible matter because explosions invariably result,
even at ordinary temperatures. Except for special use or need
anhydrous perchloric acid should not be made or stored.
Anhydrous perchloric acid undergoes partial autodis-
sociation to chlorine heptoxide and oxonium perchlorate
above its melting point, and the rate of attainment of equilib-
rium increases rapidly above —30°. 100 The existence of the
equilibrium 3 HCIO^CUO? + HCIO*' H2O has been demon-
strated from viscosity,101 kinetics of decomposition and vapor
pressure, 102 and electrical measurements. 103 Rosolovskii102 re-
ports values for the equilibrium constant (as written above)
of 0.80 x 10"4 at —10°, 1.30 x 10"4 at 20°, and 1.94 x 10"4 at
70 c . Bout and Potier103 report a value of 0.60 x 10"6 for the
self-dissociation equilibrium constant of the following re-
action:
3 HC104 *=* CIO2O7 + H*0 + + CICV
This value differs from Rosolovskii's because of the dif-
ference in formulating the oxonium perchlorate product. The
extent of self-dissociation is thus indicated in either formula-
tion at approximately 1 per cent.
Heat capacities and thermodynamic functions for an-
hydrous perchloric acid from 5 to 55°K indicate that HCIO4
exists as an independent species in the crystalline s t a t e . m At
298.15 °K its heat capacity, CP, is 28.80 cal/mol degree. A
triple point occurs at 172.0 °K.
Raman and infrared spectra of anhydrous HCIO4 and
DCIO4 have been extensively studied for all three physical
states.105'7 There is no evidence of any molecular association
in the liquid state. Bond energies, bond lengths, and standard
thermodynamic quantities have been estimated from the
data.105
The thermal decomposition of anhydrous perchloric
acid has been studied extensively.108"116 Proceeding through
the formation of CI2O7, the following decomposition products
are formed at 60-80°; O2, HClCVHaO, CIO2, and CI2.111 At 200-
439° the products are CI2, H2O, and O2.110 In addition to CI2O7,
other probable intermediates include ClOs and CIO4. At 10
mm and 300-400°, the decomposition is first order with re-
spect to HCIO4. Catalysts, in order of their activity, are
C02O3 > Mn02 > CuO > Fe 2 0 3 >CuCr 2 04 > AI2O3 > SiOa.nG

25
Reactions
An extremely powerful oxidant, anhydrous perchloric
acid causes explosions upon contact with most organic sub-
stances. It can, however, be dissolved in chloroform, methy-
lene chloride, acetic acid, and certain other organic solvents
for somewhat safer employment in chemical reactions and
syntheses. On dissolving it in acetonitrile, it interacts with
the solvent by H-bonding initially and then very slowly
transfers its proton to the acetonitrile. 117 Applications of the
anhydrous acid in various organic solvents have been re-
viewed by Burton and Praill.118
Neither hydrogen chloride nor hydrogen bromide react
with anhydrous perchloric acid but hydrogen iodide as well
as sodium iodide ignite on contact with it. Thionyl chloride
also ignites on contact but not sulfuryl chloride. Phosphor-
ous oxychloride dissolves it without reaction, while phos-
phorous pentachloride reacts to give chlorine heptoxide. A
solution of the acid in chloroform explodes violently if
poured upon phosphorous pentoxide. Reaction with iodine
yields deliquescent needles of HlOsfh 3, which when warmed
give off iodine and leave iodic acid as a residue.
Graphite reacts slowly with the anhydrous acid at room
temperature to give C2-1.9CIO-1 and ClOa.1" However, a drop
added to wood charcoal produces a violent explosion. Paper
and wood are ignited by the acid. The gas-phase reaction of
perchloric acid with hydrogen is first order in either reactant
and yields hydrogen chloride.120 Methane or ammonia is with-
out effect on the gas-phase reaction with hydrogen.121 In the
gas-phase reaction of perchloric acid with ethylene the
major products are HC1, CO, and H2O; some 1,2-dichloroe-
thane and vinyl chloride is also produced together with
minor amounts of numerous other organic compounds, in-
cluding Cs and d halides. 1 "

PERCHLORIC ANHYDRIDE
Chlorine heptoxide, the anhydride of perchloric acid,
was first isolated by Michael and Conn30 by adding anhydrous
perchloric acid very slowly to phosphorous pentoxide cooled
to —10 5 , and after one day at —10°, distilling the product
from the mixture at 82 \ They carefully reported that "the

26
 apparatus may be virtually pulverized by violent explosion,
and personal precautions must be taken accordingly." More
recently, Kolarov et ah123 described a method which involves
heating a mixture of anhydrous magnesium perchlorate with
phosphorous pentoxide in a ratio of 1 to 1-2 at 100-160" and
1-2 mm for 2-5 hours to obtain a pale yellow-to-orange prod-
uct as distillate collected in a receiver cooled to —78°. The
distillate is free of phosphoric acid but contains lower oxides
of chlorine. Use of hydrated magnesium perchlorate led to
an explosion. Replacement of magnesium perchlorate with
potassium or ammonium perchlorate gave no product.
Chlorine heptoxide. CI2O7, is a colorless volatile oil
which decomposes spontaneously on standing for a few days,
turning greenish-yellow. According to Babaeva,124 only 33.5%
decomposed in 502 min. at 80^. In the presence of 1% HCIO4
the rate of oxygen evolution is increased and the final prod-
uct is oxonium perchlorate.
Perchloric anhydride is soluble in and slowly attacks
benzene, reacts slowly with water to form perchloric acid,
reacts with iodine to form iodine pentoxide, and explodes on
contact with flame or by percussion. It volatilizes rapidly,
so that a small amount dropped onto paper or wood may
evaporate before it can react violently. Reaction with olefins
yields impact-sensitive alykl perchlorates.12'"' It is soluble in
phosphorous oxychloride, producing normal freezing point
lowering indicative of the molecular state CI2O.126

27
 REFERENCES
1. F. v o n S t a d i o n , G i l b e r t ' s A n n . , 52,197, 339 (1816).
2. F. v o n S t a d i o n , A n n . Cliim. P h y s . , 8 , 400 (1818).
3. F. C. M a t h e r s , J. A m o r . C h e m . S o c , 32, 66 (1910).
4. H. H. W l U a r d . /. Amor. C h e m . S o c , 34,1480 (1912).
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28
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63. K. Yales a n d H. Wai, /. A m e r . C h e m . S o c . 86, 5408 (1964).
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65. S. A. A t t i g a a n d C. H. R o c h e s t e r , J. C h e m . S o c . P e r k i n T r a n s . , 2,
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66. K. B. W i b e r g a n d R. J. E v a n s , /. A m e r . C h e m . S o c , 80. 3019 (1958).
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68. J. E. E a r l e y a n d T. W. Kallen, Jnorg. C h e m . , 10, 1152 (1971).
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70. G. B r e d i g a n d ). M i c h a e l , Z. p h y s . C h e m . , 100. 124 (1922).
71. A. L. G r e e n b e r g a n d G. H. W a l d e n , J. C h e m . P h y s . . H.645 (1940).

29
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73. R. Gilbert a n d P. W. M. J a c o b s , C o m b u s t . F l u m e . IB, 327 (1971],
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80. C. F e r r a d i n i , j . P u c h e a u l t . R (ulien, I, Gilles, a n d B. I.esigne,
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107. A. I. Karelin a n d Z. I. G r i g o r o v i c h . Spcclruchim. Acta, Part. A, 32A,
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31
 CHAPTER III

PROPERTIES AND PREPARATION

OF PERCHLORATES
REPRESENTATIVE (SUB GROUP A)
METAL PERCHLORATES
Ammonium and Alkali Metal Perchiorate
The anhydrous salts of the alkali metal perchlorates are
isomorphous with one another and with ammonium, thallium,
and silver perchlorates. Their crystal structures have been
determined by X-ray and optical methods. 1 With the excep-
tion of lithium perchiorate, all exhibit dimorphism under-
going transitions from rhombic to cubic forms at charac-
teristic temperatures. 2-4 Their structural transformations have
also been examined by differential thermal analysis. 5,8
Thermochemical data for various aqueous solutions of
the ammonium and alkali metal perchlorates, heats of solu-
tion, heats of formation, and thermochemical data for the
anhydrous salts have been determined by numerous in-
vestigations.7"12 Solubilities of the perchlorates in various
solvents were determined by Willard and Smith.13 Magnetic
susceptibilities 14 and molar refractions 15 have also been de-
termined. Many of the results of these various determina-
tions are compiled in Table 9.
Additional thermochemical data, not included in Table
9 nor in complete agreement with previous data, have been
reported recently. Birky and Hepler report heats of solution
at 25° for lithium, ammonium and potassium perchlorates of
—6.31, 8.02, and 12.31 kcal/mole, respectively." Vorob'ev
et al. report that the enthalpies of formation of sodium and
potassium perchlorates are —90.68 and —-101.9 kcal/mole,
respectively.17
Guenther 18 found that the solubilities of rubidium and
cesium perchlorates in the presence of univalent strong elec-
trolytes (ionic strengths 0.068 to 0.46 M) are accurately pre-
dicted by simple Debye-Huckel theory. Values obtained for
the pKsp (—log of the solubility product) of potassium,
rubidium, and cesium perchlorates at 25° are 1.944, 2,542,
and 2.380, respectively.

32
33
 Rosolovskii and coworkers 19 report that the solubility
of the perchlorates in anhydrous perchloric acid at 0° de-
pends on the radius of the cation but not its nature. The
following solubilities (in g of solute per 100 g. HCIO4 were
found: 0.106 LiClCk, 0.628 NaClCU, 4.256 KCIO*. 22.56 RbClCU,
and 68.40 CsClO-i.
Thermal decomposition and differential thermal analysis
studies of the crystalline solids have been carried out by
various investigators.3'"'20'2'" Lithium, sodium, and potassium
perchlorates yield oxygen and the respective chloride on
thermolysis."0 Gas evolution measurements and DTA22 in-
dicate that lithium perchlorate melts at 241'', decomposes at
489 , and does not undergo a phase change. By similar means
potassium and sodium perchlorates were found to undergo
endothermic phase change at 306 and 304", melt at 575°
and 468°, and decompose at 620° and 561°, respectively."
Alkali metal perchlorate solutions have been studied
extensively to gain information about the perchlorate ion.
From electrical conductance measurements, Jones23 found a
value of 67.32 ± 0.06 mhos for the limiting equivalent con-
ductance of the perchlorate ion. The standard enthalpy of
formation of the perchlorate ion in infinitely dilute aqueous
solution is —30.87 ± 0.07 kcal/g, according to Kirpichev.24
The crystal radius of the perchlorate ion is 1.85 A, determined
from results of solvation studies of alkali metal perchlor-
ates.2;i Calculations based on Raman and infrared spectra of
aqueous solutions indicate a hydration number of four for
perchlorate ions.26 The tendency for perchlorate ions to form
contact ion pairs was studied as a function of cation and
solvent by chlorine-35 nuclear magnetic resonance. 37 Factors
favoring contact ion pairing are high charge to radius ratio
of cation, low dielectric constant of solvent, and low basic
strength of solvent. The possibility of perchlorate ion co-
ordination to various metal ions and other species has been
explored by vibrational spectroscopy and other techniques
and will be discussed in a later section with transition metal
perchlorates.
A convenient general preparation method for any alkali
metal perchlorate consists of the treatment of the hydroxide,
oxide, or metal with perchloric acid followed by isolation
of the neutral salt by crystallization. Sodium and potassium

34
perchlorates can be obtained by treating their chlorides or
fluorides with nitrosyl perchlorate, NOCIO4.28 Commercial
quantities of sodium perchlorate are best prepared by elec-
trolysis of aqueous solutions of sodium chloride.39,3° Com-
mercial quantities of ammonium, lithium, potassium, ru-
bidium, and cesium are best prepared from sodium per-
chlorate by metathesis (double-decomposition) reactions, tak-
ing advantage of the much higher aqueous solubility of the
sodium perchlorate than the other perchlorates in order to
fractionally crystallize the desired perchlorate. 3 "
Ammonium Perchiorate. The anhydrous salt is pre-
pared by reaction of sodium perchlorate with ammonium sul-
fate or chloride. No hydrates are known but a triammine with
a dissociation pressure of 2 mm at —79° has been reported.*1
Its solubility in ammoniacal solutions increases slightly with
increasing ammonia concentration. 33 In aqueous perchloric
acid solutions its solubility decreases with increase of per-
chloric acid concentration and increases with increasing
temperature.™ Densities of aqueous solutions of ammonium
perchlorate at 15 J and 25° have been determined.'14 An in-
crease in volume results at 25" when ammonium perchlorate
is dissolved in water, but a decrease occurs at 15' .
Thermal decomposition of ammonium perchlorate has
been studied extensively because of its effectiveness as an
oxidant in rocket propellants and as an explosive. It is stable
at 110'", decomposes at 130 ; , and explodes at 380°.35 Below
300,J the decomposition reaction is described by the follow-
ing equation:35"™
4 NH4CIO4 -> 2 Ck + 8 HaO + 2 N2O + 3 O2
Above 400° most of the nitrogen is evolved as nitric oxide.35
The activation energy for decomposition changes from 18.9
to 29.6 kcal/mole on increasing the temperature above 240 c ,
a change that coincides with the crystal transition from or-
thorhombic to cubic at 240*. From 400' to 440°, under nitro-
gen pressure of 20 torr to control sublimation, the decom-
position of ammonium perchlorate exhibits an activation
energy of 23.4 kcal/mole and probably involves decom-
position of the vapor rather than solid phase.™
A differential thermal analysis study of the ammonium
perchlorate-lithium perchlorate has been reported, indicating

35
a simple eutectic exists at 182° and 69.5 mole % LiClCk.39
Studies of ternary and quarternary systems involving am-
monium perchlorate are cited in the Appendix.
Lithium Perchlorate. The trihydrate is obtained by
crystallization from a mixture of lithium carbonate and 70%
perchloric acid,40 or by electrolysis of a solution of lithium
chlorate. 41 Conversion of the trihydrate to anhydrous lithium
perchlorate, with a molar heat of hydration of 14.2 kcal to
overcome, requires prolonged drying at 300°.40 Pruntsev and
coworkers have patented a process for obtaining anhydrous
lithium perchlorate by heat treatment at 170-2000.42
In addition to forming a very stable hydrate, lithium
perchlorate forms a di-, tri-, and pentammine, exhibiting
dissociation pressures of 2 mm at 20°, 39.5 mm at 20', and
31 mm at —79', respectively. 31
Lithium perchlorate is thermally stable at and above its
melting point at 247', according to Markowitz. 4 Thermal de-
composition, yielding oxygen and lithium chloride, occurs
at an appreciable rate at 400 s and is catalyzed by the chloride
product.43,44 The autocatalysis is affected by addition of silver
nitrate in the lithium perchlorate melt.44'45
The solubility of lithium perchlorate in water and the
densities of the saturated solutions between 0° and 40° have
been determined, 46 also the densities of non-saturated aque-
ous solutions at 15 °.47 Solubilities in acetone and in methyl
ethyl ketone have been reported for the —50° to +50°
range.48
The structure of lithium perchlorate trihydrate has been
determined in detail by an X-ray and neutron diffraction
study,4" and its hydrate water structure has been examined
by proton magnetic resonance. 80 The lithium atoms are co-
ordinated by an almost regular octahedron of water mole-
cules (2.133 A average Li-0 bond distance), the perchlorate
ions have a regular tetrahedral structure (1.440 A Cl-O bond
distance], and each hydrogen atom forms a weak bond
(2.044 A) with one perchlorate-oxygen atom and a very weak
bond (2.617 A) with a second perchlorate-oxygen atom.
Sodium Perchlorate. Electrochemical oxidation of aque-
ous sodium chloride solutions is the most common method
for preparing sodium perchlorate. Many different studies,

36
 electrodes, and procedures have been described over the
years. Improvements recently reported include use of lead
dioxide on carbon or titanium as efficient anode materials,51'52
decreased temperatures, 53,S4 and addition of sodium fluo-
ride.55 For the anodic oxidation of sodium chlorate to sodium
perchlorate, increased current efficiency is claimed if po-
tassium persulfate is added.56 Another approach uses a lead
dioxide anode for chlorate concentrations of 250-300 g/1
and then changes to a platinum anode when the chlorate
concentration has been depleted by electrolysis to one-half
its original value.57
Sodium perchlorate forms a monohydrate, with a heat
of hydration of 2.01 kcal/mole, which can be completely de-
hydrated at 130°. The anhydrous salt is stable to tempera-
tures up to 471° before loss of oxygen and formation of
sodium chloride occurs.38 Sodium perchlorate also forms a
tetrammine. 31
The solubility of sodium perchlorate in aqueous solu-
tions,50 perchloric acid solutions,59 and in water-dimethyl sul-
foxide media60 has been studied. Conductivity of the salt
has been measured in water, 61 methanol, 62 ethanol, 63 di-
me thy If ormamide,64 nitromethane, 65 hydrocyanic acid,63 and
hydrazine. 67
A detailed crystal structure of sodium perchlorate mono-
hydrate has been determined by Berglund and coworkers, 68
who subsequently performed a neutron diffraction study09
and a deuteron magnetic resonance study of the solid.70 A
chlorine-35 NMR relaxation study of aqueous sodium per-
chlorate solutions has also been reported. 71
Studies of ternary and quaternary systems involving
sodium perchlorate are cited in the Appendix.
Potassium Perchlorate. The thermal decomposition of
potassium perchlorate is complex and has been extensively
studied because of the interest in its use as an oxidizer in
solid rocket propellants. Decomposition of the pure salt has
been detected at a temperature as low as 530°,7a although
others found it to be stable at higher temperatures. 58 ' 73 The
kinetics of its isothermal decomposition 74 under constant
oxygen pressure 75 and under its own evolved oxygen pres-
sure 76 have been investigated.

37
 No hydrates or ammoniates have been reported for
potassium perchlorate. Solubility data up to a temperature
of 265° have been determined. 77 Activity coefficients, de-
termined from solubilities in various salt solutions, have
been reported. 78 Conductivities in hydrazine, 87 hydrogen cya-
nide,86 and dimethylformamide 84 have been measured.
Alkaline Earth Perchlorates
Anhydrous perchlorates of the alkaline earth metals can
be prepared by heating ammonium perchlorate with the
corresponding oxides or carbonates. 79 They can also be
prepared by heating their hydrates with pyridine, eventually
driving off pyridine without decomposition of the per-
chlorate.80 The hydrates are prepared by treatment of the
metal oxides or various salts with aqueous perchloric acid.18
All of the alkaline earth perchlorates form hydrates, am-
mines, and pyridine adducts. The hydrates are identified in
Table 10 with data reported for their heats of hydration.8,81
The ammines correspond to M(C10-i)2 ' nNHa, where n = 2,
6, and 7 for Mg; 2, 3, 4, 6 and 7 for Ca; 1, 2, 6, 7, 10 and 12
for Sr; and 2, 5, 6 and 9 for Ba.82"84 The pyridine complexes
correspond to M(C104)i • nCCflHUN), where n — 2 and 4 for
Be; 1, 2, 4 and 6 for Mg; 1 and 2 for Ca and Sr.80,85 Their
thermal stability increases with decreasing number of pyri-
dine groups.
The basic salts M(OH]C104 (where M is Mg, Ca, or Ba)
have been prepared and characterized, 8 "
Alkaline earth perchlorates are unusually soluble in or-
ganic solvents, as evidenced by the data of Willard and
Smith13 compiled in Table 10. Curiously, calcium and barium
perchlorates are more soluble in methanol than in water. The
solubilities of magnesium, strontium, and calcium perchlor-
ates in water have been measured for the 0° to 50° range.87
Magnetic susceptibilities 14 and apparent molar volumes, 88
the latter deduced from density measurements of aqueous
solutions, are given in Table 10. Vapor pressure 8 " and den-
sity measurements 90 have also been reported for aqueous
solutions of the metal perchlorates at various temperatures.
Beryllium Perchlorate. The dihydrate has been pre-
pared by heating a mixture of beryllium chloride and oxo-
nium perchlorate to 60" in vacuum.01 Infrared spectra re-

38
 co N ^ in rf M

N
!D H
O (D
(D CO
CO (D
O CO 2o '—
Ol &'
a. o i N o d d c
o n co * in
D O ^
n °5
< CO W rn t-( r-< r-< V

< -S

< r °-
ft" °
o -S r-l CO CM O CO rH r^
O CO CD O CO H OI
^ (/) CO ' cococn^cocnNtD

225
.2 2

"as
X"

m 11
<
II
HS
ill!
»£ft " • S ' 3 - - " * ' S B » J Sob1
-§£
x x <2:
39
vealed that the product contains Be(HsO)4+2 and Be(C104)4'2
ions. It melts below 80°, forms an oxo- or hydroxy-containing
compound at 150°, loses HCIO4 and forms a solid residue
Be40(C10<)6 in the temperature range 190° to 165°, and
eventually yields beryllium oxide on further heating to 290°.
The tetrahydrate, when heated in an atmosphere of helium,
decomposes in the liquid phase to form Be(OH) (ClOt) which
then decomposes to beryllium oxide and anhydrous per-
chloric acid.92 Novoselova and coworkers have prepared
cesium, rubidium, and potassium perchloratoberyllates of
general formula M2Be(C104)4 and interpreted their infrared
spectra obtained as Nujol mulls in the 200-4000 cm'1 region.93

Magnesium Perchlorate. The anhydrous salt serves as

an extremely efficient drying agent.94 It also strongly absorbs
ammonia, ether,95 and many other polar organic vapors, 96
Vapor pressure measurements have been made of the water-
magnesium perchlorate system, indicating the existence of
a di-, tetra- and hexahydrate, but no trihydrate. 97
A thermogravimetric study indicated a gradual weight
loss with no sharp break in the temperature-weight curve for
magnesium perchlorate hexahydrate decomposing to mag-
nesium oxide.98 Another study identified (MgClJaO as the
final product of the thermal decomposition. 99 Markowitz100
reports that a mixture of magnesium oxide and chloride is
formed. A kinetic study by Novoselova and coworkers 101
indicates that magnesium perchlorate undergoes two-stage
decomposition: first to MgO • Mg(C104)2, and subsequently
to either MgO and Mg(C104)2 (below 165°) or MgO, Cb and
O2 (above 165°).
Conductances have been measured for solutions of
magnesium perchlorate in water 102 and various nonaqueous
solvents.103"105 Osmotic and activity coefficients have been
reported.108 The structure of magnesium perchlorate has been
investigated by X-ray107 and Raman108 techniques.

Calcium Perchlorate. Studies reported for calcium per-

chlorate include conductivity in acetone, 1M adiabatic com-
pressibility of aqueous solutions,109 enthalpies of dilution and
relative apparent molar enthalpy of aqueous solution,110 and
kinetics of thermal decomposition.111

«s
 Strontium Perchlorate. Conductances in methanol-
acetone solutions'11'"' and adiabatic compressibility 109 of aque-
ous solutions of strontium perchlorate have been measured.
Barium Perchlorate. Preparation and use as drying
agents of the anhydrous salt94 and of the trihydrate 112 have
been described by Smith. The ammoniates and their equilib-
rium dissociation pressures have been reported by Smeets. 113
The conductance of barium perchlorate in several different
organic solvents has been measured. 114 Vibrational spectra
of the crystalline trihydrate have been analyzed,115 indi-
cating insufficient lattice dissymmetry for a measurable
piezoelectric effect or optical activity.

Groups IMA - VA Metal Perchlorates

Although the perchlorates of all of the representative
metals have been prepared, those of Groups III to V of the
periodic table have not been studied as thoroughly or as sys-
tematically as those of Groups I and II. Perhaps the most
distinctive features of metal perchlorates as a class are their
relatively large solubilities in organic solvents, their appre-
ciable affinities for water and hydrate formation, and their
proclivities to undergo thermal decomposition and sometimes
explosions. Most of the literature published on each com-
pound is cited or briefly summarized below.
Aluminum Perchlorate. In addition to the anhydrous
salt, five different hydrates have been described with 3, 6, 9,
12 and 15 moles of water. 1 All except the 15-hydrate are
hygroscopic. The 9-hydrate has a very high water solubility,
462.75 g per 100 g of water at 0°. Existence of the 12-hydrate
has yet to be confirmed. Anhydrous aluminum perchlorate
can be prepared by drying its hydrates over phosphorous
pentoxide in vacuum at 150°, m by treatment of anhydrous
aluminum chloride with anhydrous perchloric acid,117 or by
reaction of anhydrous aluminum chloride with silver per-
chlorate in methanol or benzene. 118
Viscosities, densities, electrical conductances, molar re-
fractions, molar volumes, and compressibilities of various
aqueous solutions of aluminum perchlorate have been re-
ported.119"122 Conductivities have also been determined in
some nonaqueous solvents.117'123 Thermal decomposition

41
studies have been made identifying aluminum oxide as the
final product.20,124 Vibrational spectra indicate that the anhy-
drous compound is covalent with bidentate perchlorate
groups.125 Raman spectra of aqueous solutions 126 and proton
magnetic resonance spectra of alcoholic solutions127 have
been analyzed.
Preparation of perchloratoaluminates, with infrared
spectra indicating a general structure M[Al(C104)i(H20)2]
(where M is either cesium or rubidium], has been reported.128
Tetraalkylammonium haloperchloratoaluminates have also
been prepared.129
Gallium and Indium Perchlorates. Viscosities and
densities of aqueous solutions of gallium and indium per-
chlorate have been measured over the temperature range
18° to 90°.122 Sound velocities and densities also have been
determined at 25°. m Gallium perchlorate is very deliques-
cent, highly soluble in water, forms a 6- and a 9-hydrate, and
decomposes at 155° in a vacuum or at 175' exposed to air.
Attempts to dehydrate the hexahydrate produced a basic
salt. Isospiestic studies of aqueous solutions of gallium per-
chlorate have been made,130 and activity coefficients have
been measured. 131 Indium perchlorate has been little studied.
Thallium Perchlorate. The crystal structure of thallium
perchlorate is isomorphous with the alkali metal perchlor-
ates, except for the lithium salt. Conductivity measurements
indicate a dissociation constant of 1.00 in water at 25°.133
Activity coefficients132 and sedimentation equilibria134 have
been determined. Density and refractive indices of the solid
have been measurd.135 The solid is volatile at low pressures
at 200°,13" undergoes an orthorhombic-cubic transition at
285°, and decomposes above 430° (with an activation
energy of 54 kcal/mole) to yield TI2O3 and T1C1 (relative
amounts depend on both temperature and atmospheric
composition).130'13'
Group IVA Metal Perchlorates. Preparation of either
germanium or tin perchlorate has not been reported. Lead
perchlorate has been prepared in anhydrous form and as
the mono- and tri-hydrate. 138 It is extremely soluble in water,
forming a saturated solution at 27° with a density of 2.7753.
Use of aqueous solutions of lead perchlorate has been sug-

42
gested for determining densities of insoluble solids by the
suspension method.13" An anhydrous solution of lead per-
chlorate in methanol is explosive. Basic salts have been pre-
pared by reaction of lead oxide with perchloric acid.140 Per-
chlorato complexes of lead have been prepared also,141 pro-
viding the earliest evidence that perchlorate ions can co-
ordinate to metal ions.
Group VA Metal Perchlorates. Antimony oxyperchlor-
ate pentahydrate has been obtained in crystalline form by
cooling a solution of freshly precipitated SbfOHj.s dissolved
in warm 70% perchloric acid.142 The solid compound de-
composes above 60^ to antimony oxides, and it is dissolved
but hydrolyzed by water.
Bismuth perchlorate is prepared as the pentahydrate by
reaction of bismuth oxide and perchloric acid.142 Prepara-
tion from metallic bismuth and perchloric acid generally
leads to an explosion. m,u * Bismuth oxyperchlorate can be
obtained by hydrolysis of bismuth perchlorate in 40% per-
chloric acid. Raman spectra have been determined for both
Bi(Cl0^.j and BiOC10-t.12C The crystal and molecular structure
of u-oxo-bis(perchloratotriphenylbismuth) have been inves-
tigated, indicating that the perchlorate groups are weakly
coordinated to bismuth(V).144

TRANSITION METAL PERCHLORATES

Much of the interest surrounding the preparation and
study of transition metal perchlorates seems to have been
focused on three questions: (1) how does the nature of the
metal cation influence thermal decomposition of perchlor-
ates?, (2) can advantage be taken of the solubility of metal
perchlorates in various solvents to effect separation or other
applications?, and (3) to what extent will perchlorate ions
coordinate to metal ions? Other than this, the transition
metal perchlorates have been explored relatively little. More
attention has been given to the solution chemistry of the
transition metal ions in the presence of perchloric acid or
perchlorates than to the properties of the metal perchlorates.
The crystal structures of the hexahydrates of a number
of bivalent metal perchlorates have been determined. 10.

43
 Hexagonal structures, closely related to that of lithium per-
chlorate trihydrate, are exhibited by the manganese, iron,
cobalt, nickel, zinc, and cadmium compounds. Mercuric
perchlorate hexahydrate is trigonal, and cupric perchlorate
hexahydrate is monoclinic. Infrared spectra of some fully
hydrated transition metal perchlorates indicate that the per-
chlorate group exists in the solids as simple perchlorate ions.
For certain lower hydrates, the spectra suggest the presence
of coordinated perchlorate groups,145
Solubilities in water and other polar solvents are gen-
erally appreciable for most metal perchlorates, as exemplified
by the data compiled in Table 11. Conductivities and other
electrochemical properties of some of the perchlorates in
furfural and in Cellosolve have also been reported.114
Viscosities, densities ,and electrical conductances of 0.001-2
M aqueous solutions of chromium(III), iron(II), thorium(IV),
and silver(I) perchlorates at 20°-35° have been determined,
also molar refractions and apparent molar volumes of the
salts.120
A general method of preparing transition metal perchlor-
ates in solution has been described which consists of three
main steps: (1) dissolution of a weighed amount of the metal
in azeotropic hydroiodic acid, (2) addition of the requisite
weighed amount of standard perchloric acid, and (3) re-
moval of the hydroiodic acid by oxidation with ozone fol-
lowed by volatilization of the iodine formed.146 In some cases,
the simple expendient of dissolving the metal or its hydro-
xide, oxide, or carbonate is satisfactory. Preparation of the
pure solids is frequently complicated by the tendency of the
transition metal ions to undergo hydrolysis, yielding hydro-
xyperchlorates.
Pyridine complexes of a number of bivalent metal per-
chlorates and silver perchlorate have been prepared and
characterized. 85 Hydroxyperchlorates of zinc, cadmium, mer-
cury, and copper have been prepared by dissolving the oxides
in solutions of the corresponding perchlorates. 147
Thermal decomposition studies of some first-row tran-
sition metal perchlorates have been carried out by DTA and
TGA techniques.14!'149 Stabilities decrease with increasing
effective electrical field strength of the metal ion. Activation
energies found for the decomposition of Mn(C104)g • 6 H2O,

44
 TABLE 11. SOLUBILITY OF METAL
PERCHLORATES114
Solubility (g solute/100 ml solvent) in Solvent
Stilt Water Furfural Celiosolve
Ba(C104)2 — 50 100+

Cd(C104)2 • 6 H 2 O 478 80 145

CU(C104)2 • 6 H 2 O — 70 100+

CU(C104)2 • 2 H 2 O 259 20+ —

C0(C104)2 '' 6 H 2 O 292 60 110

Mn(C104)2 • 6 H 2 O 268 90 130

Ni(C104)2 • 6H2O 267 60 100+
Ni|C104)2 • 2H2O — 20 35
AgC104 540 40 125
Pb(C104)a — 25 105
Z n ( C 1 0 4 ) 2 ' •6H2O — 85 130

45
 Co(Cl04)3 • 3HsO, Ni[Cl04)2 " 6H2O, and Cu(Cl04)2 * 6H2O are
respectively 21.4, 24.8, 49.0-52.0, and 43.7-45.9 kcal/mole.149
The hexahydrate of chromic perchlorate undergoes simulta-
neous dehydration and decomposition to chromic oxide with-
out forming anhydrous chromic perchlorate, and the di-
hydrate decomposes at 135°-160° with an energy of activa-
tion of 17.5 kcal/mole. 150 Bel'kova and coworkers have in-
vestigated the thermal behavior of many of the rare earth
perchlorates, including determination of their dehydration
and thermal decomposition temperatures.151"153 At 500° all
the perchlorates decomposed similarly to give both MOCl
and MCls in the solid residue. Thermograms show endo-
thermic effects due to loss of water of crystallization fol-
lowed, by exothermic effects accompanying decomposition
of the perchlorate. Most of the hydrated perchlorates lose
water gradually up to 200° and begin decomposition in the
temperature range 250°-270\ Thermal decomposition is
facilitated with increasing polarization effects of the cation.
For the trivalent rare earth perchlorates, in the series lan-
thanum through holmium, increasing atomic number (de-
creasing cation radius) is accompanied by a decrease in the
decomposition temperature. 151 Thermogravimetric studies of
the lanthanide perchlorates indicate that dehydration is af-
fected more than the decomposition process by change in ap-
plied pressure. 154 Decomposition under non-isothermal con-
ditions yields chlorides. To obtain anhydrous perchlorates it
is best to operate at very low pressure and slow heating
rates to minimize formation of oxychlorides.
The perchlorate ion was long considered to be a non-
coordinating anion and thus ideally suited for avoiding com-
petitive complexation and for maintaining a constant ionic
medium in complex equilibria studies. However, findings re-
ported in the early i960's145'155'156 and the many investigations
that followed have proved this conception false. Vibrational
spectroscopy and single crystal X-ray diffraction techniques
provided the most convincing evidence of perchlorate co-
ordination. In the absence of strongly coordinating solvent
molecules or other ligands, transition metal ions are par-
ticularly receptive to perchlorate ion association and com-
plexation. Examples include the ions of chromium, 15 ' man-
ganese,158 iron,153 cobalt,158'181 nickel,158"161 copper,158"163

46
zinc,158,13" mercury,,r,i' cerium,151' and thallium.159 Although
quantitative data are lacking in most cases, the perchlorato
complexes of metal ions are invariably weak. Labile aquated
metal ions are much less prone than inert complex metal ions
to form perchlorato complexes of measureable stability. The
role of the perchlorate ion as a ligand in solution and the
myth of non-coordinating anions have been the subject of
recent review articles.159'"14

Group IB and MB Metal Perchlorates

Copper and silver perchlorates have been investigated
rather extensively, but no studies have been reported of gold
perchlorate. Heats of formation of cupric and silver per-
chlorates in aqueous solutions are —19.0 and —7.75 kcal/
mole, respectively. Cupric perchlorate is prepared by dis-
solving basic copper carbonate in perchloric acid. Its most
stable hydrated form is the hexahydrate, which has a m.p.
of 82.3'', a water solubility at 23" of 54.3% (density 2.225
g/ml), and a molar heat of solution of —4.6 kcal. Hepta-,
tetra-, and di-hydrates have also been described.1'1' Silver
perchlorate is obtained as a monohydrate on crystallization
from aqueous solution and easily converted to the anhydrous
form on drying at or above 43''. It is deliquescent, explosive,
light sensitive, and very soluble in water. At 0° a saturated
aqueous solution of silver perchlorate contains 82.07%
AgClOi and has a density of 2.7251; at 35° it contains 86.21%
AgClCM and has a density of 2.9173.
Silver perchlorate has been reported to explode on pul-
verizing it in a mortar. In spite of its hazards, silver per-
chlorate has received considerable attention because of its
unusual solubility in organic solvents and ability to form cer-
tain addition compounds. It can be used to generate anhy-
drous perchloric acid in various organic solvents by bubbling
dry hydrogen chloride through its solution in the respective
solvent. Its solubility in grams per 100 grams of solvent at
25 is 5.28 in benzene, 101 in toluene, 5.28 in aniline, and
26.4 in pyridine. It is also soluble in nitrobenzene, chloro-
benzene, glycerine, glacial acetic acid, acetonitrile, nitro-
methane, and ethyl acetate. It is insoluble in chloroform, car-
bon tetrachloride, and ligroin. Silver perchlorate is appre-
ciably associated in. many solvents.li2'(i3'I14'"i7 In aprotic sol-

47
 vents of dielectric constant greater than 30 it behaves as a
medium-strength electrolyte. 168 In anhydrous hydrofluoric
acid, hydrocyanic acid, or polar solvents of high dielectric
constant it is extensively ionized. Silver perchlorate has
proven useful as a reagent in organic synthesis, as described
in Chapter V.
Studies on zinc perchlorate include determination of
water vapor pressure of aqueous solutions (0.5 M to satu-
rated] at 0° and 50°,169 measurement of conductances in the
mixed solvent methanol-acetone, 105 determination of osmotic
and activity coefficients,106'170 and measurement of heats of
dilution with water and with perchloric acid.171
Cell potential measurements indicate that zinc and
cadmium perchlorates are completely dissociated in concen-
trations up to 0.1 M in aqueous solution (in contrast to the
non-dissociated halides) and that the mean transport num-
ber of the perchlorate ion is 0.595.172'173 Both zinc and cad-
mium perchlorates form tetra- and hexa-ammoniates. 174 The
dissociation pressure and composition of the zinc compound
have been determined as a function of temperature and its
heat of formation calculated.82'175
Mercurous perchlorate forms a tetrahydrate which can
be readily converted to the dihydrate on heating above 36Q.176
In water it can undergo three successive stages of hydrolysis,
eventually yielding mercurous oxide. Conductance and po-
tentiometric measurements of highly concentrated solutions
indicate abnormal dissociation. Mercury (I) is reported to
form a stronger perchlorato complex than mercury(II).177
Group 1MB and Rare Earth Metal Perchlorates
Aqueous solutions of scandium, yttrium, and lanthanum
perchlorates have been investigated by viscosity and den-
sity measurements over the temperature range 25° to 90°.
Results indicate that hydration of the cation is weaker the
larger its radius.178 Apparent molar volumes, specific adi-
abatic and molar compressibilities, and isobaric and molal ex-
pansions have also been determined as a function of con-
centration.179
Spedding and coworkers have conducted numerous, ex-
tensive investigations of aqueous solutions of rare earth
perchlorates. Results reported include densities, equiva-

48
lent conductances, transference numbers, 180 activity coeffi-
ents,180'181 heats of dilution at 25 ,1S2 and partial molal heat
capacities at 25 V 8 3
The solubilities of the lanthanide group metal per-
chlorates have been determined over the temperature range
25" to 50". They show very little temperature dependence.1"4
Heats of solution of the tetrahydrates in water have been de-
termined also.185
Ammoniates18f' and antipyrine complexes 18 ' of some of
the lanthanide group metal perchlorates have been prepared
and described. Thermal decomposition of the ammoniates
proceeds through several steps, yielding the metal chlorides
at 500 V 8 6
Neodymium perchlorate hexahydrate has been prepared
and dried at 170° to yield the anhydrous solid.187 Its molar
refraction and absorption spectrum in aqueous solution have
been determined. 188 Gadolinium perchlorate octahydrate also
has been prepared. It is deliquescent and very soluble in
water and in alcohol.18"'18"
Several rare earth perchlorates of the actinide group
have been prepared: solid uranyl perchlorate, 191 solutions of
neptunium(IV], (V), and [VI] perchlorates, 192 and pluto-
nium(III) perchlorate. mi
Group IVB-VIIB Metal Perchlorates
Preparations of perchlorates of only the first member
of each group have been reported. Titanium tetraperchlorate
sublimes at 70°, decomposes on aging at 70° in a vacuum,
and explodes when heated at atmospheric pressure to about
130°.194 Anhydrous titanium tetraperchlorate and chromyl
perchlorate reportedly contain bidentate perchlorato li-
gands.19'"' Vanadyl perchlorate pentahydrate, VO(C104]a*5H:iO,
has been prepared by dissolving vanadium(IV) hydroxide in
perchloric acid followed by precipitation from 3.5 M per-
chloric acid.196 On heating to 140° it loses one water of hy-
dration and is partially oxidized to vanadium(V], at 200°
loss of HCIO4 occurs, and at 300° formation of vanadium
pentoxide takes place exothermically. Absorption spectra of
solutions of vanadium(III) and vanadium(IV] perchlorates
have been determined.197 Chromic perchlorate has been pre-
pared in hydrate form containing 3, 5, 6, 9, and 10 moles of

49
water." 8 Manganese perchlorate hexahydrate is deliquescent,
melts at 155°, begins thermal decomposition at 165°, and de-
composes rapidly at 230° to manganese dioxide. Its anhy-
drous salt could not be prepared in pure state.199 Results of
an electron paramagnetic resonance study of manganese(II)
in aqueous perchloric acid solutions at 25° and 55° have
been interpreted to indicate the presence of an internal
monoperchlorato complex, with a stability constant of
(8.0 ± 1.5) x 10-'.200
Group VIIIB Metal Perchlorates
Ferrous perchlorate hexahydrate forms long green crys-
tals, is stable in air, loses four molecules of water over 96%
sulfuric acid in a vacuum, and is susceptible to air oxidation
in aqueous solution.201 Its solubility in water is 978 g/1 at
0° and 1161 g/1 at 60°. In ethanol at 20° its solubility is 865.4
g/1. Ferric perchlorate decahydrate loses four molecules of
water over concentrated sulfric or phosphorous pentoxide. It
is soluble in water to the extent of 1198 g/1 at 0° and 1517 g/1
at 60°. The existence of perchlorato complexes of iron[III)
has received considerable attention by many investigators.
Johansson, after a critical review of the literature, concluded
that more experimental work of high quality is necessary to
settle the issue.159
Cobalt(II) and cobalt(III) perchlorates have been pre-
pared, the latter by electrolytic or fluorine oxidation of the
former.202 Anhydrous cobalt(II) and nickel(II) perchlorates
reportedly can be prepared either by heating the correspond-
ing hexahydrates below their melting points in a vacuum or
by first replacing their water of hydration with dimethoxy-
propane and then vacuum heating.203 Hexapyridine com-
plexes of cobalt(II) and nickel(II) perchlorates can be
prepared by reaction of the respective perchlorate with
pyridine.85
Rhodium(III) perchlorate hexahydrate forms light yel-
low, long needles, is hygroscopic, and has a face-centered
cubic structure.204 Rhodium(IV) perchlorate is dark red and
has been prepared only in solution, by dissolving RhC)2" 2H2O
[obtained as a precipitate by oxidizing RhCU in alkaline so-
lution with sodium hypobromite) in 1 M perchloric acid.295
Palladium(II) perchlorate tetrahydrate has been prepared by

50
dissolving palladium sponge in concentrated nitric acid fol-
lowed by heating to fumes with 72% perchloric acid.20,i The
brown crystalline needles deliquesce in moist air but can
be dried over phosphorous pentoxide in a vacuum without
loss of the hydrate water.

MISCELLANEOUS PERCHLORATES
Inorganic Perchlorates
Certain non-metallic elements and compounds, notably
those of nitrogen, phosphorous and the halogens, enter into
chemical combination with perchloric acid and its deriva-
tives to provide a number of interesting inorganic com-
pounds. These are briefly reviewed below.
Nitronium Perchlorate. This compound, also called
nitryl or nitroxyl perchlorate, was first prepared by Gordon
and Spinks207 inl940 by mixing chlorine dioxide with a mix-
ture of ozone and nitrogen oxides generated by passage of
dry air through an ozonizer. It has also been prepared by
reaction of dinitrogen pentoxide with anhydrous perchloric
acid.208 Nitronium perchlorate is composed of N02 + and
ClO-T ions,209 has a very low vapor pressure, is soluble in and
recrystallizable from nitric acid, and reacts very rapidly
with water to give nitric and perchloric acids with moderate
evolution of heat. It reacts vigorously with many organic
compounds, explosively with some. Solutions of NOaClCU
in nitromethane or chloroform have been used to nitrate aro-
matic compounds. Raman spectra of crystalline nitronium
perchlorate 210 and of its nitric acid solution211 have been re-
ported. Thermal decomposition at 100-127° yields, among
other products nitrosyl perchlorate. 212 Reaction with metal
oxide yields the metal perchlorate.

Nitrosyl Perchlorate. Raman spectroscopy indicates

that the compound NOClO-i is composed of NO* and ClO-r
ions.213 It has an orthorhombic crystal structure, 214 a density
of 2.169 g/cc,215 and a heat of formation of —41.79 ± 0.08
kcal/mole.216,217 Reaction with water produces nitrogen
oxides. With methanol, nitromethane is formed. Nitrosyl
perchlorate reacts violently with many organic compounds.
It was first prepared by passing a mixture of nitric oxide and

51
nitrogen dioxide into 72% perchloric acid followed by fuming
to 140° and cooling to crystallize out the monohydrate. 218
The anhydrous salt was obtained on drying over phos-
phorous pentoxide, first in an atmosphere of nitrogen oxides
and then in a vacuum. Its thermal decomposition has been
studied extensively, and nitronium perchlorate is one of the
products.219'220
Hydrazine Perchlorate. The *hemihydrate, N2H3CI4
V2H2O, is obtained on neutralization of an aqueous solution of
hydrazine with perchloric acid. At or below 60.5° the equilib-
rium dissociation pressure of the hydrate follows the
equation221
log Prnm = — 3047.6/T + 10.98
The free energy of dehydration is 1.456 kcal/mole, and the
heat of dissociation is 13.95 kcal/mole. The anhydrous salt
can be readily recrystallized from ethanol, with a solubility
at 60° of 69 g/lOOg solution and at 0° very slight.222 It melts
at 137-8° and begins to decompose at 145°. Deflagration re-
sults if it is heated rapidly. Violent detonation occurs on
mechanical impact, shock, or friction.223
Phosphonium Perchlorate. A crystalline product, m.p.
46-47°, formulated as P[0H)4C1C>4, has been obtained through
the interaction of phosphoric and perchloric acids. Soluble
in nitromethane, its conductivity is typical of an ionic salt.
Heats of solution and formation in nitromethane are 13.9
and 11.4 kcal/mole, respectively.224'225
Selenious Acid Salt. Addition of selenious acid to ice-
cold perchloric acid, after warming slightly and then cooling
again, yields a crystalline, deliquesent product of formula
Se(OH)4C104. In nitromethane, its heats of solution and for-
mation are 4.8 and 11.4 kcal/mole, respectively, and con-
ductivity measurements are indicative of an ionic salt.225
Perchloryl Fluoride. The compound ClOsF undergoes a
number of interesting and useful reactions: at 150 to 300°
it vigorously oxidizes a variety of reducing agents, with am-
monia it forms ammonium perchlorylamide, with aromatic
compounds it forms perchloryl substituted compounds, and
for certain active hydrogen containing compounds it is an
effective fluorinating reagent.226 Surprisingly stable, per-

52
chloryl fluoride can be heated in glass to the softening point
without etching the surface. Treatment with a concentrated
strong base or heating with water at 250 to 300° in a sealed
tube is necessary for quantitative hydrolysis to fluoride and
perchlorate ions." 7
Perchloryl fluoride has been prepared by the action of
fluorine on potassium chlorate,"* by electrolylsis of sodium
perchlorate in liquid hydrogen fluoride,2"'1 and by the action
of fluorosulfonic acid on perchlorates. 23 " Its physical prop-
erties have been studied extensively; values reported for
some constants include: m.p. —146 \ b.p. —46.8 '"', heat of
vaporization 4.6 kcal'mole, critical temperature 95.13 , criti-
cal pressure 53.00 atm, critical density 0.637 g/cc, critical
molar volume 161 cc, dipole moment 0.023 ± 0.003D, and
A H f 2OH = —5.12 ± 0.68 kcal/mole.227'231'23- Thermodynamic
properties,-2fi'233 vapor pressure, 2:n surface tension,232 and
viscosity232 have been measured over various temperature
ranges.
The infrared spectrum and fundamental vibration fre-
quencies of perchloryl fluoride have been determined.227,234,;!3:'
The fluorine and three oxygen atoms are bonded separately
to the chlorine atom.

Halogen Perchlorates. Fluorine perchlorate, m.p.

—167.5 and b.p, —167.5°, has been prepared by reaction of
elemental fluorine with 60 to 72% perchloric acid.23" It is
reported to explode always on freezing or when contacted
with organic or easily oxidized matter. Chlorine perchlorate,
ClOClOs, has been prepared by the reaction at —196° of
CISO3F with cesium perchlorate or nitrosyl perchlorate. 236
Bromine perchlorate (red liquid, f.p. < —78°) was prepared
from the chlorine perchlorate (pale yellow liquid] by treat-
ment with bromine at —35 to —78 l \ 23fi Both halogen per-
chlorates react at —78° with anhydrous metal chlorides to
form metal perchlorates and with gaseous hydrogen chloride
or bromide to yield anhydrous perchloric acid and the re-
spective halogen. They also react with fluorocarbon halides
to yield novel fluorocarbon perchlorates 237 and with perhalo-
olefins to form perhaloalkyl perchlorates. 238
Iodine trisperchlorate has been obtained as a white solid

53
in excellent yield by treating iodine with chlorine perchlorate
at—150° for 70 hours. 2M - 240

Organic Perchlorates
Innumerable perchlorate salts of organic bases and
chelated metal cations have been reported. Most were pre-
pared as a convenient means of isolating the desired base or
cation in a crystalline form of definite chemical composition.
Few have been prepared for the purpose of investigating
perchlorate chemistry. All show typical perchlorate proper-
ties: moderate to good solubility in organic solvents, rela-
tively high conductivity, tendency to decompose at or above
their melting point, and the capability of explosive behavior
if overheated or detonated by mechanical shock. Representa-
tives of this group are the perchlorates of methylamine,"'1
pyridine,242 benzenediazonium ion,743 and bis(l,10-phenan-
throline)copper(l) ion.244 Except for the diazonium perchlo-
rates, which are extremely explosive and easily detonated,
most perchlorate salts can be handled safely with suitable
care.
Aromatic aldehydes, ketones, ethers, and various pyran
compounds combine with perchloric acid to form crystalline
salts that can be isolated, although they often decompose on
storing.24,ri Some controversy exists as to whether certain of
these should be considered as oxonium or carbonium salts.
True carbonium salts, identifiable by their conductivity in
nonaqueous solvents and intense color, are formed by the
action of perchloric acid on triarylmethylcarbinols or by
metathesis of the cloride with silver perchlorate in a suitable
solvent.240'247
Preparation of the gem-diperchlorate [CH3]2C(OC10a)a
from acetone has been reported, also the gem-diperchlorate
of 2-butanone.248 Alkyl diperchlorates have been prepared
from dienes and diols.24!)
Relatively few perchloric esters have been isolated or
their physical properties determined because of the severe
explosion hazard involved. Meyer and Spormann 250 reported
that the methyl ester (b.p. 52°) and the ethyl ester (b.p. 89°)
are extremely powerful explosives and that they were unable
to avoid explosions in spite of great foresight taken in hand-
ling them. Ethyl perchlorate, first prepared by Hare and

54
Boye in 1841,251 is immiscible with water and slowly hy-
drolyzed by it. It is soluble in ethanol but still dangerous to
handle, because a dilute solution can spontaneously burn
completely away if not explode. Preparations reported for
other perchlorate esters include those for ClCHsCHfOHJCHa-
C1CV53 HOC2H40C2H4C104,2f'2 trichloromethyl perchlorate, 2 '"-
254
acetyl perchlorate, 255 and benzoyl perchlorate. 255
Perchloryl compounds can be prepared by reaction of
perchloryl fluoride with aromatic compounds in the presence
of anhydrous aluminum chloride.a5(i For example, perchloryl
benzene (C6H5CIO3, mp.m.A3°, b.p. 232°) has been obtained
by this Friedel-Crafts type reaction. It is reasonably stable.
Hyrolysis with potassium hydroxide solution produces po-
tassium chlorate and phenol, and nitration with a mixture of
concentrated sulfuric and nitric acid yields 3-nitroperchloryl-
benzene (pale yellow needles, m.p. 49 to 50°]. The relative
stability of the perchloryl group towards reductants is evi-
denced by its ability to remain intact when 3-nitroperchloryl-
benzene is reduced with stannous chloride in hydrochloric
acid to 3-aminoperchlorylbenzene. 256 Although relatively safe
to prepare and use, perchloryl aromatic compounds are sensi-
tive to vigorous shock and high temperature.

55
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63
 CHAPTER IV
APPLICATIONS OF PERCHLORIC ACID AND
PERCHLORATES IN CHEMICAL ANALYSIS
DISSOLUTION AND OXIDATION
OF INORGANIC SAMPLES
Perchloric acid, alone or in combination with other
strong mineral acids, is frequently employed to dissolve or
decompose inorganic substances as an integral step in their
analysis. In most instances the dissolution process is more
efficient than if some other acid were to be used. Often more
suitable oxidation states of the sought-for substances are at-
tained which facilitate or make possible their subsequent
separation or determination. Moreover, if silica is present its
determination can be greatly simplified by the strong de-
hydrating action of hot, concentrated perchloric acid.
In the process of decomposing samles with hot acid,
especially when prolonged heating or evaporation to dryness
is involved, it is important to be aware of the possibility of
loss of certain elements by volatilization. Obviously, one
should guard against unintentional loss of sought-for sub-
stances. A second consideration is the possibility that sepa-
rations can be achieved if sufficient differences in volatility
exist. The necessary information for such considerations in
the case of perchloric acid treatment of inorganic samples
is available from comprehensive studies by Hoffman and
Lundell1 and by Chapman and coworkers. 2 Their results are
compiled in Tables 12 and 13, respectively. Among the many
conclusions that can be draw from these results, the follow-
ing are illustrative: (1) chromium can be satsifactorily sepa-
rated from manganese in analyses of steels containing high
percentages of chromium, (2) rhenium can be separated from
molybdenum by distillation from a mixture of perchloric and
phosphoric acids to which hydrobromic acid is slowly added,
and (3) boron can be quantitatively distilled from perchloric
and hydrofluoric acid mixtures.

Soils, Clays, and Silicates

A number of successful procedures have been described
for the analysis of soil, clays, and silicates based upon the
 TABLE 12. ELEMENTS AND APPROXIMATE
PERCENTAGES VOLATILIZED'
llements Acid Mixture Distilled at 200-220°
HC1- HBr-
HC1 HBr H3PO4 H3PO4 HC1 HBr
HCIO4 HCIO, HCIO4 HCIO4 H 2 S0 4 H 2 S0 4

Alkali metals 0 0 0 0 0 0
Alkaline earth metals 0 0 0 0 0 0
Rare earth 0 0 0 0 0 0
Group III elements 0 0 0 0 0 0
Cu, Ag,, Zn , Cd 0 0 0 0 0 0
Al, Ga, In 0 0 0 0 0 0
Ti, Zr, Hf, Si, Pb 0 0 0 0 0 0
Nb, Ta , W 0 0 0 0 0 0
Fe, Co, Ni 0 0 O 0 0 0
Rh, Pd, Ir, Pt 0 0 0 0 0 0
As(III) 30 100 30 100 100 100
As(V) 5 100 5 100 5 100
Au 1 0.5 0.5 0.5 0.5 0.5
B 20 20 10 10 50 10
Bi 0.1 1 0 1 0 1
Cr(III) 99.7 40 99.8 40 0 0
Get 50 70 10 90 90 95
Hg 75 75 75 75 75 90
Mn 0.1 0.02 0.02 0.02 0.02 0.02
Mo 3 12 0 0 5 4
Os" 100 100 100 100 0 0
P 1 1 1 1 1 1
Re 100 100 80 100 90 100
Ru 99.5 100 100 100 0 0
Sb 2 99.8 2 99.8 33 99.8
Se 4 5 5 5 20 100
Sn(II) 99.8 100 0 99.8 1 100
Sn(IV) 100 100 0 100 30 100
Te 0.1 0.5 0.1 1 0.1 10
Tl 1 1 1 1 0.1 1
V 0.5 2 0 0 0 0

"Results of Hoffman and Lundell1 obtained by treatment of the ele-

ments or their salts with 15 ml of 60% perchloric acid (or sulfuric
acid), maintained at 200-220" and purged moderately with a stream of
dry carbon dioxide, while slowly adding 15 ml of hydrochloric or
hydrobromic acid. The same procedure was used for distillations from
phosphoric acid mixtures except 5 ml of sirupy phosphoric acid was
added to the 15 ml of 60% perchloric acid prior to heating and
further treatment.
t>To prevent precipitation of Ge0 2 , the hydrochloric or hydrobromic
acid was added before heating the final solution.
«At 200 to 220" no osmium was volatilized from the sulfuric acid solu-
tions; however, at 270 to 300° the osmium was completely volatilized.

65
 TABLE 13. EFFECT OF TREATMENT WITH
PERCHLORIC AND HYDROFLUORIC ACIDS'
Elements Retained Elements Lost
Na, K B, 100%
Cu, Ag, Au Si, 100%
Be, Mg, Ca, Sr, Ba Ge, up to 10%
Zn, Cd, Hg As, 100%
La, Ce Sb, up to 10%
Ti, Th Cr, varies greatly
Sn, Pb Se, varies greatly
V, Bi Mn, up to 3%
Mo, W, U Re, varies greatly
Fe, Co, Ni

"Results of Chapman and coworkers 2 obtained by fuming a solution

of the element or its compound with a mixture of 10 to 15 ml of
70% perchloric acid and 8 to 10 ml of hydrofluoric acid in a platinum
dish at approximately 200°.

66
use of perchloric acid to decompose the sample. Among the
first to advocate the use of perchloric acid in cehmical analy-
sis, Willard and Cake 3 found it to be ideally suited for the
determination of silica in silicates due to the strong dehydrat-
ing action provided by boiling the concentrated perchloric
acid solution. Cadariu recommended perchloric acid for the
analysis of silicon in silicates 4 and for use in decomposing
slags, dried cement slurries, Portland cement, bauxite, and
clay.5 Marczenko and Stepien0 found perchloric acid to be
superior to hydrochloric or sulfuric acid for the determination
of silica in aluminosilicates. Turek 7 employed mixed hydro-
fluoric and perchloric acids for the decomposition of clay
samples. Pratt 8 described a similar treatment for soils to be
analyzed for total potassium and sodium.
The determination of nitrogen and phosphorous in soils
can be effected rapidly through the use of decomposition with
perchloric and sulfuric acid according to Meshcheryakov. !l
Sommers and Nelson10 employed perchloric acid digestion in
a sealed tube to achieve rapid and precise determination of
phosphorous in a wide range of soils. Metson and Collie"
have proposed the use of nitric, perchloric, and phosphoric
acids for the determination of total sulfur in soils. Potassium
content of soils have been estimated by a perchloric acid
extraction method.12
Metals, Alloys and Ores
Chemical analysis of steels has been greatly facilitated
by use of perchloric acid to dissolve the samples, dehydrate
the silica, and oxidize chromium and certain other consti-
tuents to oxidation states that enable their simple and direct
determination. Willard and Cake:! were the first to describe
the use of perchloric acid for the determination of silica in
steel. They were also first in determining chromium in steel
employing hot, concentrated perchloric acid to oxidize the
chromium to dichromate.' ;i Another first by Willard in
analytical applications of perchloric acid was its use for the
determination of vandium as well as chromium in steel,
chromite, and ferrochromium/' 4 Perchloric acid soon became
the acid of choice for steel analysis, with many advocating
its use and reporting extensions and refinements of Willard's
methods. Smith and Smith1'"' demonstrated that a mixture
of perchloric and phosphoric acids rapidly and completely

67
 dissolves chromium steel, stainless steel, tungsten steels,
and metallic tungsten. Seuthe and Schaefer16 employed
a m i x t u r e of p e r c h l o r i c and n i t r i c a c i d s to d i s s o l v e
steel in analysis for chromium, vanadium, tungsten, and
phosphorous. Rapid determination of silicon, chromium,
nickel, and molybdenum in steel and copper alloys was re-
ported by Birckel." Use of perchloric acid for analysis of
steels, alloyed cast iron, and other materials was described
by Raab,18 Croall,18 and Bertiaux and coworkers. 20 The de-
termination of chromium in steel by oxidation with perchloric
acid has been investigated by various workers to find
optimum conditions. 2l,a2 More recently, spectrophotometry
methods for the determination of manganese, 23 chromium
and manganese, 24 and zirconium 25 have been reported for the
analysis of iron and steels following treatment with per-
chloric acid.
Methods utilizing mixtures of perchloric, phosphoric,
and sulfuric acids for the determination of manganese in
tungsten and ferrotungsten, 26 for the determination of
chromium in chromite ores,27 and for the analysis of ferro-
chrome28 have been published by Smith and coworkers,
Iron ore samples are rapidly and completely dissolved
by an equal volume mixture of 72% perchloric acid and
85% phosphoric acid. Goetz and Wadsworth 29 reported that
nearly all iron ores are dissolved within 10 minutes and the
determination of iron can be accomplished without diffi-
culty by passing the perchloric acid solution through a Jones
reductor and titrating with standard cerium(IV) sulfate or
potassium permanganate.
Analysis of chromium metal for lead following pre-
liminary dissolution and oxidation with perchloric acid has
been reported.30 A method for the determination of iron and
cobalt in stellite (a nonferrous alloy of Cr, W, Co, and C) has
been described31 based upon the use of perchloric acid oxi-
dation and distillation of chromium as chromyl chloride to
eliminate interference by chromium.

Miscellaneous Samples
Perchloric acid is reported to be more effective as a de-
composition reagent than either sulfuric or hydrochloric
acid in the microdiffusion method for the determination of

88
carbonate in dolomite, siderite, calcite, and magnesite. 32
To prevent their interference in the electrolytic deter-
mination of lead as lead dioxide, the ions and solute species
of chloride, bromide, arsenic, tin, and antimony can be re-
moved as volatile species by evaporating to fumes with
perchloric acid.:i:1 In this application the use of a nitric and
perchloric acid mixture would be advisable to guard against
explosion if appreciable organic matter is present.
Lichtin34 determined chromium in chrome alums by
iodometric titration after fuming with perchloric acid to
oxidize chromium(III] to chromium[VI). Chitnis and co-
workers 35 report that plutonium can be determined coulo-
metrically after oxidation to plutonium(VI) by fuming with
perchloric acid. The presence of iron(III] is necessary for
quantitative oxidation.
According to Smith;i(! the mineral alunite, KaAlefOH]^-
(SO-i)4, dissolves readily in hot 72% perchloric acid if finely
ground. Thus the determination of silica and the R2O3 metals
can be carried out without recourse to sodium carbonate
fusion.

WET OXIDATION OF ORGANIC MATTER

Determination of trace metals in organic matter, es-
pecially that of biological origin, has grown increasingly more
common and widespread with the growing appreciation of
the important influences exerted by trace metals in chemical
and biological systems. Nondestructive methods of analysis,
such as activation analysis and X-ray fluorescence, are pre-
ferred but not always practical for this purpose. Chemical
methods require that the trace metals be released into solu-
tion prior to measurement. Sometimes simple treatment
with a solvent suffices, but very often total destruction of
the organic matter is necessary. Two general methods for
this are dry and wet oxidations. Although each has its ad-
vantages, wet oxidation methods have proven more widely
applicable and acceptable. Wet oxidations are versatile in
that a variety of oxidants and solution combinations are
available to provide a range of decomposition strengths and
special advantages. Risk of loss by volatilization of sought-
for elements or by their retention in vessels is much less due
to the lower temperatures required. Moreover, mineral resi-

69
dues remain dissolved during wet oxidations, precluding the
need for the sometimes difficult step of dissolving strongly
ignited residues.
An excellent source of information on the destruction
of organic matter has been compiled by Gorsuch.37 Articles
by Smith,38"*0 Middleton and Stuckey,41 and Gorsuch42 also
provide critical reviews and useful general information.
Total decomposition of organic matter by wet oxidation
is most efficient when perchloric acid is employed, either in
combination with nitric acid or with nitric and sulfuric acids.
Numerous studies support this conclusion. Mixtures con-
sisting of only nitric and sulfuric acids are also satisfactory
but less commonly used, because their action is slower with-
out perchloric acid and special care must be taken to avoid
charring. Once formed, char is extremely slow to be oxi-
dized by nitric and sulfuric acids without perchloric acid.
The use of nitric acid with perchloric acid provides an im-
portant safety margin against possible violent or explosive
reactions. Easily oxidized matter is destroyed by the nitric
acid before the perchloric acid becomes sufficiently concen-
trated to exert its strength on the remaining more difficulty
oxidized matter.
After an extensive study on recovery of trace elements
from organic materials, Gorsuch42 concluded that the use of
nitric and perchloric acids was not only very effective but
trouble-free. The only significant loss was that of mercury.
For the destruction of very obdurate materials, he found a
mixture of sulfuric, nitric, and perchloric acids to be es-
pecially effective. Variable amounts of lead were lost, how-
ever, presumably due to coprecipitation of lead sulfate with
calcium and other insoluble sulfates. Loss of lead was great-
est from samples high in calcium. Gorsuch evaluated eight
different oxidation methods; three involved wet oxidation,
four were dry ashing procedures, and one was a hybrid. He
concluded that wet oxidation with nitric and perchloric
acids was the most satisfactory for recovery of all the trace
metals investigated, with the single exception of mercury.
Trace element losses during mineralization of biological
material by dry ashing and by wet oxidation with a mixture
of nitric, sulfuric, and perchloric acids were investigated by
Pijck, Hoste, and Gillis43"48 using radiochemical methods. Ele-

70
ments investigated included Ag, As, Au, Co, Cr, Cu, Fe, Hg,
Mn, Mo, Pbt Sb, V and Zn. Biological materials included
blood, urine, powdered vegetable and muscular tissues. Dry
ashing resulted in losses of the following elements, even at
temperatures of 500-550 : Ag, As, Au, Fe, Hg, and Sb. Above
700° some loss of all of the elements studied occurred. Re-
coveries from wet oxidations were quantitative for all ele-
ments except As, Au, Fe, Hg, and Sb. Use of a reflux con-
denser in wet oxidations prevented loss of any of the trace
metals.
Several other comparative studies have been conducted
to determine which of various decomposition methods leads
to the best recoveries of trace elements in organic matter.
Allcroft and Green4G found that wet oxidation with a mix-
ture of sulfuric and perchloric acids gave the best recoveries
of arsenic from animal tissues, in comparison with three
other wet oxidation mixtures and a dry ashing procedure
which involved addition of magnesium nitrate. According to
Jackson,47 recovery of iron from biological material is quan-
titative by wet oxidation with a mixture of nitric, sulfuric,
and perchloric acids but incomplete by dry ashing, even
though dry ashing be carried out with added sodium car-
bonate, calcium carbonate, or sulfuric acid treatment.
Hiscox,48 in comparing six different methods for destruction
of plant materials in determination of cobalt, concluded that
wet oxidation with nitric and perchloric acids gave the
best recoveries.

Efficiency and Products of Wet Oxidations

Although a number of studies have been devoted to

evaluating recoveries of trace elements from various com-
positions, relatively little attention has been paid to de-
termining the completeness of destruction of organic sub-
stances by wet oxidation with perchloric acid and its mix-
tures with other acids. If clear and colorless solutions are
obtained on wet oxidation, it is tacitly assumed that com-
plete destruction of organic matter has been accomplished,
particularly if quantitative recovery of trace elements is
achieved. Such an assumption, however, can lead to serious
error in some applications if organic matter still persists in
solution. A knowledge of the identities and concentrations of

71
 residual organic matter, as well as inorganic products, is im-
portant in avoiding interference in subsequent measurements.
Martinie and Schilt40 investigated the efficiencies of per-
chloric acid mixtures for the wet oxidation of various organic
substances and attempted to identify any residual matter.
Eighty-five different model compounds and common sub-
stances of a representative nature were subjected to pro-
longed wet oxidation with a mixture of nitric and perchloric
acids. The final solution from each sample was then screened
for organic residue by proton magnetic resonance spectro-
metry, ultraviolet spectrophotometry, and carbon microanaly-
sis. Most solutions exhibited some ultraviolet absorption, and
approximately one-half retained measurable carbonaceous
matter. In general, compounds with N-methyl, S-methyl, C-
methyl, and pyridyl moieties proved the most resistive to-
wards wet oxidation. Glycine, alanine, proline, methionine,
histidine, glutamic acid, and lysine were among the amino
acids incompletely destroyed. Certain N-eontaining hetero-
cycles, purines, and pyrimidines also resisted total destruc-
tion. Residual products were identified in a number of cases.
Ammonium perchlorate, for example, was commonly found
as a product of the wet oxidation of any N-containing com-
pound. Most of those substances that proved resistive to-
wards oxidation by the nitric and perchloric acid mixture
yielded to total destruction in a reasonable period of time on
treatment with a nitric, sulfuric, and perchloric acid com-
bination. Complete oxidation of pyridine and 2,4,6-trimethyl-
pyridine, however, proved especially slow. Vanadium(V],
cerium(ni), and copper(II] with selenized Hengar granules
were found to exert catalytic influences on the wet oxidation
of these and certain others.

The chemical compositions of solid residues obtained by

heating to dryness various inorganic substances with a mix-
ture of nitric and perchloric acids have been examined by
Mansell, Tessner, and Hunemorder 50 using X-ray diffraction
and infrared spectroscopy. Unfortunately, the complete iden-
tification of the residues proved difficult, even for relatively
simple systems. Lack of crystallinity and complexity of mix-
tures, presumably of varying degrees of hydration and
crystallinity, precluded reliable identifications by X-ray dif-
fraction in many instances. Temperature control during heat-

72
ing to dryness was employed in relatively few cases, hence
many of the results are difficult to interpret or reproduce.
Since evaporation to dryness is not a common practice in wet
oxidation procedures, the results obtained by Mansell and
coworkers are of limited value. Their observation, for ex-
ample, that iron and magnesium compounds are converted
to the respective oxides could not apply to residues in solu-
tion. Further investigation of the fate of inorganic constitu-
ents and their identities after perchloric acid oxidations
would be helpful in establishing the effectiveness and suit-
ability of wet oxidation procedures for various analytical
purposes. Few surprises are anticipated, however, based upon
current knowledge of the chemical properties of perchloric
acid and commonly encountered inorganic species.
The oxidation strength of perchloric acid can be regu-
lated by control of temperature and concentration. Studies
by Smith40 indicate that hypophosphorous acid is not oxi-
dized by boiling 30% perchloric acid, iron(II) is oxidized
slowly to iron(III) by boiling 50% perchloric acid, and vana-
dium(IV) is readily oxidized to vanadium(V) with boiling hot
60% perchloric acid. Chromium(III) is easily oxidized to di-
chromate by boiling 72% perchloric acid. The redox potential
is estimated to increase to approximately 2.0 volts on in-
creasing the concentration of the boiling acid to 72% (boiling
point 203°). Practical applications of graded oxidation po-
tentials obtained by controlled concentrations have been de-
scribed by Smith40 and by Diehl and Smith.51
The major products of thermal decomposition of per-
chloric acid vapor at 279-471° in Pyrex vessels are chlorine,
oxygen, and water, according to a study by Gilbert and
Jacobs.52 Small amounts of hydrogen chloride are also pro-
duced. Decomposition is first-order at all temperatures, with
surface effects important at low temperatures. The CIO
radical is possibly a chain carrier at low temperature. In
another study51' of the decomposition of perchloric acid, the
products Oa, Ck, HC1, CIO2, and CkO, were found by chroma-
tographic analysis. Chlorine produced during wet oxidations
with perchloric acid can be readily removed after completion
of the oxidation by dilution with water followed by boiling
to drive out the chlorine." Dilute perchloric acid does not
generate more chlorine on boiling.

73
 Oxidative degradation of glucose with 40% perchloric
acid at 40^ yields a mixture of products which includes
gluconic, glucaric, glucuronic, tartaric, oxalic, citric, succinic,
and levulinic acids.''1 The degradation is gradual and can be
monitored by specific rotation, reducing power, ultraviolet
spectrophotomery, and paper chromatography. Oxidative de-
gradation of starch by perchloric acid, studied by the same
methods, revealed the production of glucose and two ketoses
and saccharic, glucuronic, gluconic and oxalic acids.33,:)S
Total degradation of glucose and starch to carbon dioxide
and water occurs if concentrated perchloric acid is employed
at reflux temperature.411
Oxidation of hydrazine in perchloric acid in the presence
of molybdenum{VI) as catalyst produces ammonium per-
chlorate, nitrogen, hydrogen, chloride, and water.57 The re-
action is zero-order in hydrazine and first-order in perchloric
acid, with an activation energy of 19 kcal/mole.
Dichlorodinitromethane has been obtained as an oxida-
tion product in the digestion of tobacco with mixed perchloric
and nitric acids at 200°.5S Collected by vacuum distillation, it
exploded during distillation at atmospheric pressure. It was
obtained also from the digestion of fir saw dust but not from
nicotine, petroleum, charcoal, or graphite. The danger of ex-
plosion arises only when the dichlorodinitromethane is col-
lected in concentrated form by distillation. No explosions
have been reported by others with experience in wet oxida-
tion of tobacco by nitric and perchloric acid mixtures.:jiM!I

Procedures and Special Techniques

A great number of procedures, many differing only
slightly, have been described for perchloric acid oxidations
of organic and biological matter. Choice of which literature
procedure to use is not generally critical because most will
accomplish the intended purpose. There are, however, cer-
tain aspects that any given procedure should possess if it is
to be both safe and effective. Easily oxidizable material
should first be boiled with an excess of nitric acid present.
Large samples (greater than a few grams) of organic matter
should not be taken until small samples {0.1 g) of the same
material have proven safe to oxidize. Any fat, oil, or volatile
organic substance that is not miscible with perchloric acid

74
should be decomposed with a mixture of sulfuric and nitric
acid prior to boiling with perchloric acid, otherwise a violent
reaction may occur in the hot vapor phase between the per-
chloric acid and organic vapors. Contrary to some published
procedures, care should be taken not to evaporate to dry-
ness solutions that contain heavy metal perchlorates or in-
completely decomposed organic matter. The risk of explosion
is too great. Wet oxidations should be carried out using fume
eradicators*ts or special safety hoods that can be readily
washed down periodically to remove build-up of acid con-
densates. Pyrex or Vycor reaction flasks should be employed.
Also safety shields and safety glasses should be routinely
used.
Most workers may prefer to follow the wet oxidation
procedure decribed in the literature method specifically di-
rected at their particular problem. For those undertaking
possible revisions or new applications one of the methods
described in the following paragraphs may prove helpful.
Plant material, certain animal products, foodstuff, and
many different natural and synthetic products 49 can be wet
oxidized effectively by the following procedure.

Transfer an accurately weighed 1-g sample into a 125-ml

conical flask and add 10 ml of 68% nitric acid and 5 ml of
70% perchloric acid. Piace flask on a hot plate at low heat
(under a fume eradicator iM or in a hood, behind a safety
screen) and heat slowly to boiling. Interrupt the heating when
necessary to avoid excessive foaming, Increase the heating rate
gradually so as to boil away the nitric acid in a period of
approximately 15 min. During this period the temperature of
the mixture will increase from about 120° to 140° and then
rapidly to 203°, the boiling point of 72% perchloric acid.
Interrupt the heating as needed to prevent foaming. Continue
heating at the fuming point or gentle boil for 15 min. or
until the solution is nearly colorless or only a faint yellow
color persists. Allow the solution to cool, add 20 ml of dis-
tilled water, and heat to boiling to expel chlorine. After
cooling and dilution to volume, aliquots can be taken for
analysis of trace metals, phosphorus, and sulfur.

Animal tissues, proteins, heterocyclic compounds, pu-

rines, pyrimidines, certain polymers and other obdurate ma-
terials40 can be wet oxidized effectively by the following
procedure.

75
 Fig. 4.1. Apparatus for performing wet oxidations in
the open laboratory consisting of a Vycor or Pyrex flask
fitted with a refluxing still head and fume eradicator. Gaseous
products are removed by aspiration, and loss of spray is
prevented by the reflux still head.

Transfer an accurately weighed 1-g sample into a 125-ml

conical flask, add 5 ml of concentrated sulfuric acid, and
heat at boiling for 15 min. After allowing the solution to
cool, add 10 ml of 68% nitric acid and heat at a rate to
cause the nitric acid to distill out over at least a 15-min.
period. Cool the solution, add 5 ml of 70% perchloric acid,
place the flask on a hot plate under a fume eradicator 38 or
hood behind a safety screen, and heat at a rate to cause
perchloric acid to reflux gently (about half-way up the sides
of the flask) for 15 min. or until the solution is colorless or
only faintly yellow. Allow to cool, diute carefully with 20 ml
of distilled water, and heat to boiling to expel chlorine. If
chlorine does not interfere in subsequent analysis, this final
boiling step can be omitted. After dilution to a known volume,
aliquots can be taken for analysis of trace metals.

A practical innovation has been devised by Monk59 to

minimize risk of violent reactions in perchloric acid oxida-
tions. The sample of organic matter is first treated with
fuming nitric acid and then added in small portions to boil-
ing perchloric acid, maintained at its azeotropic concentra-
tion (72%), and oxidation is completed before addition of
the next portion. Another approach by Monk,60 suitable for

76
controlled oxidation of cellulose, involves using a mixture
of 7 ml of concentrated nitric acid and 3 ml of 72% per-
chloric acid for each gram of cellulose taken for analysis.
The cellulose dissolves readily at 60-70", and the resulting
solution is then added in 5-ml portions to boiling 72% per-
chloric acid, completing the oxidation before adding the
next portion.
An automated digestion system has been described by
John"1 for use in nitric and perchloric acid oxidations of plant
tissues. Samples are allowed to stand overnight in the acid
mixture, then digested in a system enclosed by metal and
heated with precise temperature control to slowly reach 203°
and until wet oxidation is complete.
A rapid and safe method of heating mixtures of nitric
and perchloric acids in a microwave oven has been de-
described. 62 Evacuation and trapping of the acid, fumes, how-
ever, is troublesome.
Retention of elements [sulfur, arsenic, mercury, etc.)
that form volatile products during wet oxidations can be
achieved by use of special apparatus such as that designed
by Bethge63 or by Kahane and Kahane.64 See Figure 4.2.
Smith and coworkers have investigated a number of
special mixtures for wet oxidations. Sulfuric and perchloric
acid mixtures enable control of oxidizing conditions at graded
potentials by control of concentrations. 51 Periodic and per-
chloric acid mixtures serve effectively and safely at lower
temperatures by promoting extensive degradation of high
molecular weight species to smaller, more easily oxidizable
fragments.65 The generation in situ of perchloric acid through
the action of nitric and hydrochloric acids on ammonium per-
chlorate provides slowly advancing oxidation potentials and
minimizes risk of carbonization and uncontrolled reaction
rates. 66

Determination of Nonmetals
Nitrogen determinations by the Kjeldahl method are
greatly facilitated by use of perchloric acid in conjunction
with sulfuric acid to complete the digestion of samples in
minutes rather than hours. Mears and Hussey, 67 the first
to advocate use of perchloric acid as a digestion aid in
Kjeldahl determinations, obtained accurate results under

77
 Figure 4.2. Bethge digestion apparatus for retention
and/or collection of volatile products formed during wet
oxidation of organic matter with perchloric acid mixtures.
By suitable adjustment of 3-way stopcock condensed vapors
can be returned to the digestion flask, collected in the air-
cooled section of the apparatus, or delivered into an external
vessel.

strict but easily controlled conditions. They noted loss of

nitrogen when perchloric acid was used in excess of that
necessary to decompose the organic matter. Best results
were obtained using 1-g samples treated with 25 ml of con-
centrated sulfuric acid plus 1 g of copper sulfate and 2 ml
of 60% perchloric acid, heated at such a rate that the di-
gestion cleared in no less than 3 nor more than 7 min. and
heated further at least 15 min. after clearing. Later investi-
gators reported different degrees of success with the method
when applied to a greater variety of substances.68"71 Con-
siderable controversy arose regarding the question of nitro-
gen loss; however, recent studies have helped to dispel much
of the confusion. Moore and Diehl72 demonstrated that low
results for nitrogen can be caused by the action of chlorine
or hypochlorous acid (produced from decomposition of per-

78
chloric acid) on ammonium salts. Diluting the mixture after
digestion with a solution of sodium sulfite proved effective
in destroying chlorine and hypochlorous acid, resulting in
quantitative recovery of ammonia. Ginsburg and Shche-
glova73 found that no loss of nitrogen occurs in the digestion
of plant material if the following procedure [similar to the
Mears and Hussey procedure"7) is employed.

Place a 0.2-g sample of ground plant material into a

50-mI Kjeldahl flask and add 0.1 g of copper sulfate, 5 ml
of concentrated sulfuric acid and 0.5 ml of 60% perchloric
acid. Let stand for 30-60 min. Heat slowly for 5-7 min. until
a chestnut-brown paste is formed and then more strongly
until a light-blue clear solution forms in about 15-20 min.
After cooling and dilution to a known volume, aliquots can
be taken for determination of nitrogen, phosphorus, and
potassium.

Batey and coworkers 74 observed that nitrogen loss can

be avoided by adding the perchloric acid in small amounts,
diluted with concentrated sulfuric acid. Others have stressed
the importance of avoiding large excesses of perchloric
acid.75"77
According to Sloane-Stanley and Jones78 small scale di-
gestions in test tubes with only perchloric acid are satis-
factory for microdetermination of nitrogen in tissue sections
and certain aldehyde derivatives.
Nitrogen, phosphorous, and potassium can be de-
termined in a single sample following wet oxidation by a
mixture of sulfuric and perchloric acid.77,79'80
Phosphorus and calcium in plant and animal materials
have been determined rapidly and accurately in a single
sample without interference, following wet oxidation with
nitric and perchloric acids.81 Phosphorus content of animal
feeds have been determined rapidly with good results using
sulfuric and perchloric acid digestion catalyzed by sodium
molybdate. 82
Sulfur has been determined in coal,83 wood and paper
pulp,84 and rubber.85-87 The methods involve oxidation with
nitric and perchloric acids to destroy organic matter with
conversion of sulfur to sulfate, which is subsequently pre-
cipitated and determined as the barium salt. A procedure

79
based on the use of a mixture of periodic and perchloric
acids for the determination of sulfur in coal has also been
described.88
Wet oxidation methods have been successfully applied
to the determination of arsenic in biological materials89'"0 and
medicinals, 91 of silica in lungs82 and plant matter,"3 and of
iodine in organic compounds. 94

Determination of Trace Metals

Applications of perchloric acid oxidations for trace
metal determinations are so numerous that only selected
examples are cited below. For a systematic review, cover-
ing most of the elements, the monograph by Gorsuch" should
be consulted.
One of the earliest applications of perchloric acid oxi-
dations for trace metal determinations was described by
Goss95 in 1917 for the determination of tin in canned foods.
By the early 1930's, wet oxidations with perchloric acid mix-
tures had been employed by Fabre and Kahane96 for the de-
termination of As, Hg, Cu, Mn, and Cr in toxicological ma-
terials; by Gieseking, Snider and Getz87 for the determination
of Ca, Mg, K, and P in plant material; and by Gerritz81 for
the determination of Ca and P in biological materials.
Illustrative of the scope of perchloric acid oxidations,
methods have been published for the determination of vari-
ous metals in brain tissues, 98 coal,88 feeds,99 whetlerized
carbon,100 metal chelate compounds, 101 biological material,' 02
plant material,67'103 wines,104'105 beer,106'107 milk,106 and milk
products. 108
Wet Oxidation is the method of choice for the destruc-
tion of organic matter in which lead, mercury, chromium
and other volatile elements are to be determined. Webber,109
for example, found that lead is lost from hay and pasture
samples when dry ashed above 450° but not when oxidized
with nitric and perchloric acids. Kozelka and Kluchesky110
recommended use of sulfuric, nitric, and perchloric acids
for the destruction of blood and soft tissue and the use of
nitric and perchloric for bone in the determination of lead
in biological samples. For the determination of mercury and
other metals in fish and in coal, Feldman 111 avoided volatility
losses in wet oxidation through careful control of refluxing

80
and evaporation using temperature programing and an in-
sulated air condenser. Excellent recoveries of chromium and
of iron from different leathers, both chrome- and vegetable-
tanned, were obtained by Smith and Sullivan112,118 following
wet oxidation with mixed nitric, perchloric and sulfuric acids.
In an investigation of different acid mixtures for wet
oxidation of metal chelate compounds, Tsuchitani and co-
workers 101 concluded that a 3 :1 mixture of 62% nitric acid
and 60% perchloric acid provided superior metal recovery
in most cases. They recommended that metal content of a
chelate be determined by decomposing 10-20 mg of sample
with 10 drops of the acid mixture, heating until dense fumes
of perchloric acid result, dilution with water and appropriate
buffer, and titration with EDTA and appropriate metallo-
chromic indicator.

TITRIMETRIC REAGENTS
Perchloric acid possesses a number of exceptional prop-
erties that greatly enhance its application in certain titri-
metric determinations. It is one of the strongest acids known,
its metal salts are water-soluble with few exceptions, its
volatility is low, it is not oxidizable, it is not readily re-
ducible in dilute solution, and its metal ion complexing
ability is extremely weak.
The exceptional acidic strength of perchloric acid makes
it the titrant of preference for the determination of weak
bases, especially for nonaqueous titrimetry. Employed in con-
junction with aprotic or predominately acidic solvents thai
do not exert an appreciable leveling effect on its strength, per-
chloric acid is superior to hydrochloric, nitric, and sulfuric
acids for the titration of a host of organic and inorganic bases
too weak to titrate successfully in aqueous solutions. Ex-
amples of solvents suited to its use in such applications in-
clude glacial acetic acid, dioxane, acetic anhydride, methyl
isobutyl ketone, sulfolane, acetone, and acetonitrile. The
number of weak bases successfully titrated in such solvents
with perchloric acid is impressive and includes such very
weak bases as the alkali halides, urea, and caffeine. For
further details and literature references the interested reader
is referred to an admirably concise book on the subject by
Fritz.114

81
 An interesting but apparently little used standard in
acidimetry is 73.60% perchloric acid, prepared by vacuum
distillation and described by Smith and Koch.115 Distillation
products obtained at 2 to 7 mm pressure varied only ±0.03%
in acid'content and ±0.0004 in specific gravity. Another per-
chloric acid-type acidimetric standard, pyridinum perchlor-
ate, was proposed by Arndt and Nachtwey1115 in 1926.
For the accurate determination of carbonate in alkali
and alkaline earth carbonates, Norwitz and Galan117 recom-
mend the use of perchloric acid in place of either sulfuric or
hydrochloric acid. Treatment of samples with an excess of
standard acid, followed by boiling to expel carbon dioxide,
and titration of the excess acid with standard sodium hydro-
xide is uncomplicated by volatilization loss of acid or by
precipitation of insoluble salts if perchloric acid is employed.
The inertness of dilute perchloric acid towards oxidants
and most reductants affords an important advantage to its
use as a solvent medium for various redox titrations, particul-
arly those that require the presence of a strong acid to pre-
vent metal ion hydrolysis or precipitation. Another significant
advantage of perchloric acid in such applications is that its
very weak complexing tendencies do not give rise to adverse
alterations in formal potentials of the redox species involved
in the titration. For such reasons, perchloric acid is com-
monly selected whenever an acid is required in adjusting
solution conditions for redox titrimetry.
Formal potentials of the cerium(IV) — cerium(III) redox
couple depend upon both the nature and the concentration of
acid used in preparing the system. Highest formal potentials
are attained using perchloric acid, providing thereby an
extremely strong cerate oxidant suitable for the determi-
nation of a large variety of reducing substances. For example,
Smith and Duke118,119 employed cerium(IV) perchlorate to
determine polyhydric alcohols, sugars, hydroxy-acids, and
certain ketones. Similarly, Ignaczak and Dziegiec120 deter-
mined p-quinone, p-aminophenol, p-phenylenediamine, p-
aminobenzoic acid, and sulfanilic acid.
Metal perchlorates are occasionally selected as titrants
for the precipitimetric or compleximetric determination of
certain anions and organic complexing agents. Choice of the
perchlorate rather than some other salt of the metal ion of

82
interest can sometimes minimize interferences, solubility
problems, competitive complexing, or coprecipitation error,
depending on the nature of the titration reaction. An example
of such an application is the use of mercury(I) perchiorate
for the precipitation titration of halides and pseudo-
halides, 121,1 " Mercurous perchiorate has also been recom-
mended as a titrant for the reduction of iron(III) thiocyanate
to the iron(II) state123 and for the biamperometric determi-
nation of molybdate and gold.124 Sharma and Gupta12'1 em-
ployed thallic perchiorate for the oxidimetric determination
of thiourea, thiosulfate, and sulfite.

PRECIPITATION AND EXTRACTION

REAGENTS
The earliest analytical application of perchloric acid was
for the precipitation and detection of potassium, as described
by Serullas12fi in 1831. Through the use of 95% ethanol to
further decrease the solubility of potassium perchiorate with-
out adversely decreasing the solubilities of lithium, sodium,
and the alkaline earth perchlorates, the method was de-
veloped into a quantitative procedure. Proving much simpler
and more precise than the Fresenius method, which involves
precipitation of potassium with chloroplatinate, the perchior-
ate method was widely employed1-7"120 and extensively in-
vestigated130"1*' for possible further improvement. By the
1950's, however, the more rapid flame photometric method
completely supplanted the perchiorate gravimetric pro-
cedure for the determination of potassium.
Use of perchloric acid to precipitate and recover po-
tassium from sea salt brine or bittern has been studied by
Bakr and Zatout.134 Recovery proved more quantitative at low
temperature, for high perchloric acid ratio, and with increas-
ing potassium chloride concentration.
Sodium perchiorate has been recommended by Deniges13'
as a microchemical reagent for the precipitation of alkaloids,
as well as for potassium, rubidium, and cesium. Ammonium
perchiorate can be employed to precipitate cobalt, nickel.
manganese, and cadmium salts from aqueous ammonia.13"
The detection and isolation of many different organic bases,
alkaloids, carbonium, oxonium and thionium compounds can
be achieved by precipitation with perchloric acid.1,1'"140

83
 The extraction of acid-soluble phosphorus compounds
from plants is more reproducible with dilute perchloric acid
solutions (0.2-0.5 M) than with either hydrochloric or tri-
chloroacetic acid, according to Sokolov141 When the extracts
are neutralized for phosphate determination, interference
from bivalent cations can avoided by addition of EDTA prior
to neutralization. 142
Perchloric acid has been found to be an efficient ex-
traction reagent for the recovery of total endogeneous platelet
serotonin as well as for the analysis of other platelet con-
situents such as nucleotides. 143
Nucleic acids can be quantitatively extracted from
animal tissues using 2% perchloric acid at 90° with 30-min.
incubation, according to Webb and Lindstrom.144 Conditions
for the extraction of DNA and RNA from tobacco pollen with
perchloric acid and sodium chloride have been investigated
by Suss.145 Use of perchloric acid extraction as an histo-
chemical technique has been evaluated by several groups
of investigators.146'148 The following order of effectiveness in
solubilizing membrane-bound proteins and non-electrolytes
was observed by Hatefi and Hanstein: 148 SCN"> C104">
guanadine> u r e a > Cl"> F".
The pronounced tendency for perchlorate ions to form
stable ion pairs with large, symmetrical cations provides
the basis for liquid-liquid extraction and subsequent determi-
nation of a great variety of substances. Examples include the
spectrophotometric determination of cobalt with 8-quino-
linol,149 flame photometric determination of trace metals with
1,10-phenanthroline,150 and the extraction-infrared spectro-
p h o t o m e t r i c determination of cetyltrimethylammonium
ions.151

DEPROTEINIZATION AGENT
Perchloric acid serves as an effective precipitant for pro-
tein removal prior to determination of other constituents in
biological fluids such as blood, milk, and urine.152,153 It is also
useful for the isolation of protein-free metabolites, peptides,
amines and amino acids.153
The behavior of serum albumin in acidic perchlorate
solutions has been investigated by Cann.154 Precipitation of

84
albumin occurs first on lowering the pH to 3.7, but the
precipitate redissolves progressively as the pH is lowered
further and is redissolved completely at pH 3.2-2.3. Precipi-
tation occurs again on lowering the pH below 2.3. Irreversible
denaturation results on prolonged exposure to acidic per-
chlorate solutions. Tamura and coworkers 155 also have
studied the precipitation of protein from cow serum, report-
ing that approximately 2-3% protein remained in the super-
natant after treatment with 4 vols, of 3% perchloric acid.
Isolation of tissue mucoids (glucoproteins) and sero-
mucoids from human and buffalo serums using perchloric
acid followed by precipitation with ethanol has been reported
by Kumar and coworkers. 156
Comparison studies have demonstrated that perchloric
acid is very satisfactory in place of trichloroacetic acid for
use in deproteinization of biological materials prior to the
determination of formaldehyde 157 and prior to citrate determi-
nations.158

DRYING AGENTS
Anhydrous magnesium perchlorate affords a number of
advantages as a drying agent in comparison with phosphorus
pentoxide, its nearest competitor in dehydrating power.153 Its
capacity for water absorption is several times greater, it
does not become sticky upon handling nor form channels
through use, and it contracts in volume on absorbing mois-
ture. Being neutral it is useful under conditions where the
acidic action of phosphorus pentoxide interferes, and it can
be recovered and reactivated repeatedly. Willard and Smith 1 ' 9
found that anhydrous magnesium perchlorate and phos-
phorus pentoxide are equally efficient in drying moist air
flowing at a rate not over 5 liters per hour. More precise in-
vestigations to determine the relative efficiencies of various
chemical desiccants have since been performed. Bower,16"
employing phosphorus pentoxide as the ultimate standard
of dryness, found a value of 0.002 mg of residual water per
liter of air dried over anhydrous magnesium perchlorate.
Diehl and Trusell,161 using a cold trap cooled in liquid nitro-
gen to determine residual water directly by weight con-
densed, found that the residual content of moist nitrogen gas

85
passed over anhydrous magnesium perchlorate was 0.2/^g/l
while that passed over phosphorus pentoxide was 3.6ju-g/l.
A commercial form of magnesium perchlorate, Anhydrone,
with hydrate water content corresponding to Mg[C104)2 *
I.48H2O gave a value of 1.5/ig/l. Thus, although some ques-
tion may remain as to its relative efficiency compared to
P2O5, the effectiveness of anhydrous magnesium perchlorate
as a drying agent has been convincingly demonstrated. It
combines high capacity with extraordinary affinity for mois-
ture without loss of porosity or ease of handling.
Widely employed as a standard desiccant, anhydrous
magnesium perchlorate is marketed under the trade names
Dehydrite and Anhydrone as well as its generic name. Its
use in laboratory desiccators was described by Smith, Bern-
hart and Wiederkehr.1"- Other applications have been re-
viewed by Smith11" and by Druce.164 It absorbs not only water
but also ammonia, alcohols, and other highly polar vapors.
Considerable risk of explosion arises if much alcohol or other
easily oxidizable substance has been absorbed and the spent
magnesium perchlorate is to be regenerated by heating. This
should not be attempted; instead the magnesium perchlorate
should be dissolved in water and recrystallized before de-
hydration. Methods for regeneration 165 and of determination
of water content166 of magnesium perchlorate desiccant have
been described by Smith. A method for preparing the desic-
cant with added indicator to give visual evidence of its con-
dition when spent has also been described.167
The trihydrate of magnesium perchlorate also serves as
an effective desiccant.188 Although less effective at ordinary
and higher temperatures than the anhydrous salt, it can
be prepared in a more porous form and thus is ideally suited
for use in combustion train procedures for steel and organic
analyses. At 0° it compares favorably with the anhydrous
form in drying gases at equal flow rates.159
Barium perchlorate has been extensively studied as a
dehydrating agent.160"171 Although its efficiency is consider-
ably less than that of anhydrous magnesium perchlorate, it
is easier to prepare and regenerate in anhydrous form. Heat-
ing between 140 and 400° drives out absorbed water without
fusion or physical disruption of the solid. In drying effi-
fiency it is comparable to anhydrous calcium chloride.

86
Bower160 found 0.82 and 0.36 mg residual water per liter of
air dried over anhydrous barium perchlorate and calcium
chloride, respectively. Diehl and Trusell161 observed 0.60
and 0.99 mg residual water per liter of nitrogen dried over
barium perchlorate and calcium chloride, respectively.
Barium perchlorate also is an efficient absorbent for am-
monia and other small polar molecules.
Although considerably limited in scope of application,
an azeotropic solution of perchloric acid (72.5% HClO-t, b.p.
203° at 760mm) exerts a strong dehydrating action at its
boiling point. Willard and Cakea greatly improved and sim-
plified gravimetric silica determinations by taking advantage
of this property. Silica in metals, silicates, limestones, etc. is
rendered more insoluble, more quantitatively recoverable
without necessity of evaporation to dryness and baking, and
less contaminated when concentrated perchloric acid is em-
ployed in place of hydrochloric acid. Presumably other hy-
drous oxides could be similarly dehydrated if inert to boil-
ing 72% perchloric acid. To be practical the dehydration
should be essentially irreversible so that perchloric acid can
be washed from the dried solid.

MISCELLANEOUS
An important innovation in the widely used Babcock
method172 for the determination of butterfat in milk and milk
products was introduced by Smith, Fritz, and Pyenson17'' in
1948. By replacing concentrated sulfuric acid with a mixture
of equal parts by volume of 72% perchloric acid and glacial
acetic acid, they demonstrated that only one centrifugation
was necessary. No charring of sugar occurs, and the presence
of various flavor additives and egg products do not inter-
fere. Their method is especially advantageous in application
to ice cream.
In a study of methods for determining fat in meat and
fish products, Rudischer174 reported that the fat content is
released from 5-g samples by digestion with either 5 ml of a
1:1 mixture of 70% HC104 and 100% HaPC>4, 5 ml of a 1:4
mixture of 70% HClCU and 85% H3PO4, or 20 ml of 30%
potassium hydroxide.
Perchloric acid has found use as a chromogenic reagent

87
for the detection and determination of steroids. Extraction
into chloroform followed by treatment with perchloric acid
and heating at 56° yields different colors for various
steroids.175 Color reactions of steroids and some thirty dif-
ferent aromatic aldehydes with perchloric acid have been
described by Few.176 Kimura and Harita177 studied the re-
action of testosterone with perchloric acid, concluding that
the chromophore is produced from isomeric olefins which
are intermediates in the chromogenic reaction.
Detection of sugars and sugar derivatives on paper
chromatograms 178 and cellulose thin layers179 can be achieved
by spraying with perchloric acid solutions. Godin180 em-
ployed an aqueous solution of perchloric acid and vanillin
for the detection of sugar alcohols and ketoses. Others181"183
found that this same spray mixture also detected deoxysugars
sensitively. Nagaswa and coworkers 179 observed that the use
of aqueous perchloric acid alone greatly reduced the sensi-
tivity of detection of sugar alcohols thereby increasing the
selectivity of the test for deoxysugars and ketoses. They per-
formed a comprehensive and systematic study of color re-
actions of sugars, sugar derivatives and related compounds
after chromatographic separation on cellulose thin-layers
syrayed with aqueous perchloric acid (5%), dried at room
temperature, and heated for 10 min. at 80°. Their paper179
should be consulted for details and applications. The use of
perchloric acid as a spjay reagent for the detection of amino
acids in paper chromatography has been described by Giri.184
More recently, a simple and specific determination of keto-
hexoses with urea and perchloric acid has been reported. 18j
Use of a hydrogen-perchloryl fluoride flame for flame
spectrophotometry has been investigated by Schmauch and
Serfass.186 The flame is easily controlled, has a very low
background, produces primarily atomic lines and metal
fluoride and chloride band radiation, and the spectra pro-
duced for many metals are suitable for analytical purposes.
Oxide bands of refractory metal oxides were only a minor
interference, despite high oxygen content in the flame.
Addition of ammonium perchlorate to the sample solu-
tion in atomic absorption spectrophotometry is recommended
by Oguro for the elimination of iron interference in de-

BB
termination of chromium187 and for enhancement in the
atomic absorption of europium.188
Use of mixed hydrochloric and perchloric acids in
forced-flow anion exchange chromatography, investigated
by Seymore and Fritz,' 89 permits several separations that are
otherwise impractical due to the inability of hydrochloric
alone to elute certain metal ions. Examples include the sep-
aration of arsenic(III), antimony(III), and bismuth(III] and
the separation of nickel(II), palladium(II), and platinum(IV).
Other applications of perchlorates in analysis include
the following: silver perchlorate solution as an absorbent
for acetylene in the determination of lithium carbide in
metallic lithium,190 lithium perchlorate trihydrate as a source
of water for coulometric generation of hydrogen ions in non-
aqueous titrimetry of weak bases,1511 perchlorate ion as a
probe in NMR studies of protein anion binding,132 perchloric
acid as an oxidant for dithionite-citrate in the determination
of reductant soluble iron phosphate in soil,1"3 and magnesium
perchlorate as an oxidant in the determination of biochemical
oxygen demand of waste waters. 194

89
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95
 CHAPTER V
APPLICATIONS IN ORGANIC SYNTHESIS,
INDUSTRY, AND COMMERCE
Catalysts
Esterification and Acetylation. Perchloric acid serves
as a highly effective catalyst in the acetylation of cellulose
and has been extensively studied in this regard due to the
considerable commercial importance of cellulose esters.1'2
The catalytic activities of various acids in acetylation of
cellulose have been found to parallel their Hammett acidity
functions in the acetylating medium, 3 with perchloric acid
exhibiting the greatest activity. Studies suggest that the
catalytic activity is associated with acetyl perchlorate,
formed through reaction of perchloric acid with acetic anhy-
dride. The rate of formation of acetyl perchlorate is con-
siderably faster than the rate of cellulose acetylation, 4 and
the equilibrium constant for the formation reaction is 33
l/mole. s Kinetics of the reaction between perchloric acid and
acetic anhydride also have been studied.6,7 Patent claims by
Lamborn8 indicate that the catalytic action of perchloric acid
can be easily controlled by addition of metal chlorides, bro-
mides, phosphates, or sulfonates. Some degradation occurs
during acetylation of cellulose catalyzed by sulfuric and
perchloric acid mixtures. 9 If sulfuric acid is added after com-
pletion of acetylation with perchloric acid, formation of
sulfo-esters is prevented. 10
Processes have been patented for the preparation of
acetic acid esters of hydroxyethyl cellulose using perchloric
acid and zinc chloride as catalyst 11 and for the preparation of
mixed esters of cellulose using 0.2 to 0.5 per cent per-
chloric acid.12 A process for treating wood by impregnation
with a solution of magnesium perchlorate in acetic anhy-
dride followed by heating has been reported to increase the
weight of wood by formation of cellulose acetate. 13
Esterification of acetic, stearic, or benzoic acid in ethanol
is catalyzed by alkaline earth chlorides, bromides, nitrates,
and perchlorates. 14 Stearic and acetic acids behave similarly,
both with and without catalysts; their rates of esterification
at 80° increase 500-fold with 0.5 N calcium perchlorate. Cal-

96
cium bromide provides less catalytic activity, calcium
chloride even less, and calcium nitrate the least. Benzoic
acid completely resists esterification in ethanol at 80° but
reacts readily in the presence of calcium perchlorate.
Ring acetylations with acetic anhydride are catalyzed as
effectively by sodium perchlorate and acetyl perchlorate in
glacial acetic as by acetylium perchlorate or perchloric acid,
according to Mathur and coworkers. 15 A 30 to 50 per cent
conversion of monohydric or polyhydric phenol ethers into
corresponding acetophenones can be generally achieved
within 4 to 6 hours at 40 to 60°. Perchloric acid is a much
more effective catalyst than zinc chloride for acetylation of
quinones with acetic anhydride. 16 Acetylation of hydroxy-
deoxybenzoins with acetic anhydride and perchloric acid
has been reported. 17 The catalytic acetylation of aromatic
and heterocyclic compounds in the presence of perchloric
acid has been described by Dorofeenko.18
Polymerization. Perchloric acid, acetyl perchlorate, and
various metal perchlorates find application as catalysts in a
variety of polymerization reactions. The polymerization of
styrene and its catalysis by perchloric acid has been studied
most extensively because of its considerable commercial
importance.
Styrene undergoes polymerization in the presence of
phenol, using acetic acid as solvent and perchloric acid as a
catalyst, at a rate proportional to the square of the per-
chloric acid concentration, according to Lilley.10 Pepper and
Reilly20 found first order kinetics for the polymerization in
ethyl chloride, dichloroethane, and carbon tetrachloride-
dichloroethane, catalyzed by perchloric acid. From a study
of the mechanism of polymerization of styrene in carbon
tetrachloride initiated by anhydrous perchloric acid, Hamann
and coworkers 21 proposed that the rate determining step may
be formation of an ion-pair between perchlorate and a car-
bonium ion in association with neutral molecules of styrene.
The kinetics of polymerization in methylene chloride,
measured by stopped-flow methods at 0 to —80°, indicate
the presence of a transient intermediate. 22
Use of acetyl perchlorate as a catalyst for the selective
dimerization of styrene to 1,3-diphenyl-l-butene 23 and for
the dimerization of methylstyrenes 24 has been described by

97
oxymercuration of D-glucal triacetate with mercuric per-
chlorate has been described by Honda, Dulenko, and
Zhdanov. 55
Solvents
In addition to its use as a solvent for various inorganic
reactions and substances, perchloric acid has proven useful
as a solvent for certain organic materials, notably cellulose
and cellulose derivatives, 56,57 acrylonitrile polymers,58'59 and
crystalline methacrylonitrile polymers. 60
Perchloric acid can be used alone or in conjunction with
other solvent media to selectively dissolve alcohols, ketones,
and amines from resinous materials. 61
The dissolution of fir cellulose by 45-65% perchloric
acid has been found to be preceded by formation of addition
compounds of the oxonium type.62 Optimal conditions for
formation of addition compounds at 20° are 62 to 65 per
cent perchloric acid and 15-min. treatment time. Extensive
dissolution of the cellulose occurs with longer treatment. No
ester formation was observed.
Explosives, Propellants, and Pyrotechnics
The use of ammonium perchlorate as an ingredient in
various explosive mixtures was first described by Oscar
Carlson in British and Swedish patents granted in 1897.
Since that time various perchlorates have been extensively
studied and implemented in a great variety of explosive and
combustive devices. Much of the information available on
the subject is in the patent literature. Undoubtedly a great
deal more is known but restricted to the realm of classified
and trade secret information. A brief survey of the uses of
perchlorates is provided in the following paragraphs. For
further details and literature citations the reader is referred
to reviews by Gale and Weber,63 Kast,64 Medard,65 and
Girard.88
Inorganic perchlorate explosives generally consist of
mixtures of either ammonium or potassium perchlorate as
oxidant with sulfur and/or various organic materials as fuels.
Special additives are commonly present also to modify shock
sensitivity, caking qualities, products of combustion, ex-
plosive characteristics, bulk density, and water-repellancy.
Explosive mixtures that contain potassium perchlorate are

100
generally less powerful than those with ammonium per-
chlorate and are suitable for special purposes, such as blast-
ing soft rock and coal deposits. Numerous composite ex-
plosives have been investigated, including ammonium per-
chlorate in combination with other oxidants such as nitrates,
nitrocompounds, and organic perchlorates. With a large
variety of choices available among oxidants, fuels, and ad-
ditives, the number of different combinations and compo-
sitions that can be formulated is enormous. An almost end-
less array of composite explosives can be formulated, each
with its special advantages and limitations.
Perchlorate explosives afford certain advantages over
dynamite and other nitroglycerine explosives. They are safer
to handle, less sensitive to shock, considerably less affected
by freezing, free from exudation in warm climates, and rela-
tively nontoxic. Their explosive action is somewhat slower
and extends laterally more than dynamite, but they are
capable of producing relatively greater destructive effects.
Perchlorate explosives can be formulated in a great variety
of compositions, permitting their adaptation to a wide range
of purposes. For example, a suitable choice between ex-
plosive compositions enables fragmentation of hard rock into
either minute pieces or large blocks.
Organic perchlorates are unique as explosives in as
much as the oxidant and combustible material are both
present in the same molecule. Consequently, thej tend to
undergo very rapid reaction and violent explosion when
detonated. The perchlorates of guanidine, dicyanodiamidine.
aniline, pyridine, methylamine, hydrazine, and metal hydra-
zines are typical examples.
Ammonium perchlorate is manufactured on a large scale
for use in preparing propellant mixtures formulated to un-
dergo relatively slow burning so as to produce nearly uni-
form acceleration in propulsion of roGkets, missiles, air
planes, and other projectiles. Both liquid and solid propellant
mixtures are prepared. An example of the liquid kind con-
sists of a suspension of ammonium perchlorate in nitro-
methane. Solid propellants are basically of two types: com-
posite and homogeneous. The former consist of solid oxidant
particles dispersed in a matrix of the fuel; the latter are
colloidal mixtures of the oxidizer and fuel in which the

101
separate phases are not readily distinguishable. Commonly
used solid fuels include natural and synthetic polymers, such
as polysulfide rubber, hydrocarbon rubbers, epoxy resins,
and polyester resins. Composite propellants that contain
potassium perchlorate as the primary oxidant generally have
higher burning rates, higher flame temperatures, and denser
smoke production compared to those with ammonium per-
chlorate as principal oxidant. Homogeneous or colloidal pro-
pellants generally contain a nitrate or nitro compound as the
major oxidant and a perchlorate as a supplemental oxidant.
Numerous recent patents describe various organic per-
chlorates as oxidants for use in rocket propellant com-
positions."7"75
Relatively slow burning compositions are required for
time fuses, signal flares, and pyrotechnics. Many of the
formulations described to date contain one or more metal or
organic perchlorates to serve as auxiliary oxidants. Some
recently described examples include the use of zirconium
perchlorate in pyrotechnical lacquer for primers,' 6 am-
monium and potassium perchlorates in signal flares,77
tungsten perchlorate in time delay pyrotechnic composi-
tions,78 and potassium and titanium perchlorates in fire-
crackers. 70
Mixtures of sulfamic acid and ammonium perchlorate
have been described in a patent 80 for use in producing a dense
smoke or fog. Ignition of an optimum mixture of the two
components results in a rapid, self-sustaining reaction yield-
ing hydrogen chloride and sulfur trioxide as combustion
products. In the presence of moist air these products absorb
water and give rise to a dense cloud or fog-like mist.

Electrolytes

Electropolishing. The technique of electrolytic oxida-

tion is extensively employed to bright polish metal surfaces
and to remove surface irregularities in high precision ma-
chining of metal parts. In this electrolytic process the metal-
lic sample is immersed in a suitable electrolytic solution and
a controlled electric current is passed between the sample as
the anode and some suitable metal serving as the cathode.
Two important requirements of the electrolyte are that (1)
it is not readily oxidized, and thus does not compete with

102
 the oxidation of the anodic sample, and (2) it forms readily
soluble salts with the metal ions produced by anodic re-
action of the sample to be bright polished. Perchloric acid
or its salts possess both of these desirable attributes and thus
find frequent application as electrolytes in electro-plating
baths. Nonaqueous solutions are most commonly employed,
because their use facilitates control of current densities,
over-voltage effects, and solubilities.
Among the first to report studies on electropolishing,
Jacquet and Rocquet81 found that iron and steels can be
electrolytically polished in a bath consisting of acetic anhy-
dride and perchloric acid, kept below 30°. A d.c. voltage of
50 v., a current density of 4 to 6 amps/dm 2 , and an aluminum
cathode were employed. Their procedure was evaluated fur-
ther by Pellissier, Markus, and Mehl82 and applied to the
electropolishing of tin, aluminum, lead, and alloys of lead
and tin. Adjustment of the current density is important in
electropolishing in order to prevent evolution of gas at the
anode at too high densities or objectionable etching of the
sample surface at too low a current density.
Numerous electrolytic solutions have been investigated
and applied in electropolishing of metals. Those which con-
tain either perchloric acid or a metal perchlorate in their
formulation include the following: (1) acetic anhydride and
perchloric acid,81"84 (2) acetic acid and perchloric acid,85"87 (3)
acetic acid and sodium perchlorate, 88 [4] ethanol and per-
chloric acid,89""1 (5) methanol and perchloric acid,92'93 (6)
ethanol, ethylene glycol monobutyl ether, and perchloric
acid,94'"5 and (6) dimethylsulfoxide and perchloric acid.96
Procedures for bright polishing the following samples
have been published or described in patents: aluminum and
its alloys,82"84'88"92 iron and steel,81'85"88'90'93 n i c k e l and its al-
loys,87,95 razor blades, 97 tin and lead alloys,82 and zirconium
and its alloys.94
The hazards associated with the use of perchloric acid
in electropolishing baths have been emphasized and reviewed
by various writers.98"100
Voltaic Cells and Batteries. Use of perchlorate salts in
electrochemical cells and batteries have received consider-
able attention in recent years owing to the improved per-
formance, lower concentration polarization, and longer shelf -

103
life afforded by such electrolytes. Considering the com-
mercial importance of electrochemical devices it is not sur-
prising that much of the information is of a proprietary
nature. A brief summary of the literature and patents on the
subject is provided in the following paragraphs. The cells
are classified according to anode material employed.
Aluminum anode cells, with aluminum perchlorate elec-
trolyte and manganese dioxide cathodes have been described
by Schumm.1"1 These dry cells are similar to Leclanche dry
cells, but they possess relatively low concentration polari-
zation and low activation polarization which compensate for
their relatively high internal resistance. Corrosion resistance
of aluminum is satisfactory in the presence of the electrolyte
buffered at pH 0 to 4; however, gas production and swelling
during discharge cause problems in sealing the cells against
leakage.
Lead storage cells, employing perchloric acid in place
of aluminum is satisfactory in the presence of the electrolyte
Cadariu and Schonberger.103 Although better performance is
claimed, some difficulty arises in fabricating lead dioxide
adherent electrodes.
Lithium batteries, attractive for their low weight, have
been explored extensively. Gaines and Jasinki104 described
lithium - nickel sulfide batteries and their design to improve
performance at high discharge rates and at low temperatures.
A solution of lithium perchlorate in tetrahydrofuran is em-
ployed as the electrolyte in a patented lithium - metallic
chromate battery. 105 According to the patent claims, it is
light-weight, high in energy density, and both chemically and
dimensionally stable. Another battery of high energy density,
also patented, is the lithium - copper sulfide battery de-
scribed by Garth.1015 Its electrolyte consists of a solution of
lithium perchlorate in 1,3-dioxolane. A self-sealing battery
with a nonaqueous electrolyte was patented by Alder,107 con-
sisting of a lithium anode, a copper sulfide cathode, and an
electrolyte containing by weight 10% lithium perchlorate,
23% 1,2-dimethoxyethane, and 67% tetrahydrofuran. A Ger-
man patent10" describes electrolyte solutions of high electric
conductance for use in lithium batteries consisting of 1 - 2.5 M
lithium perchlorate in mixtures of dioxolane or propylene
oxide with 10 to 50% propylene carbonate or ethylene car-

104
bonate. A lithium-copper sulfide battery, in which the anode
chamber containing lithium perchlorate in dimethylforma-
mide is separated by means of a membrane from the cathode
compartment containing sodium polysulfide, has been de-
scribed in a recent German patent.1""
Magnesium batteries are relatively light in weight and
possess high energy densities, especially those that incor-
porate perchlorate salts as electrolytes. It is reported that a
magnesium - silver(II) oxide cell with an electrolyte solution
2.4 M in sodium perchlorate, 1.6 M in lithium perchlorate and
0.1 M in sodium borate can be operated over several weeks at
an average discharge voltage of 1.6 volts.11" A German patent
has been granted to Bauer and Winkler111 for a magnesium -
manganese dioxide battery employing a depolarizer contain-
ing 7.4 to 24% magnesium perchlorate, 6% carbon black,
2% barium chromate, 0.5% magnesium oxide, and 57%
manganese dioxide. Garbacher112 received a British patent for
a high energy density electrochemical cell comprised of a
magnesium anode, a cathode of perforated stainless steel
coated with sintered nickelous ammonium sulfate and an
organic-bound mercuric oxide paste containing carbon black,
and an electrolyte of 5 N magnesium perchlorate containing
5% lithium perchlorate incorporated into fibrous carriers.
Various investigators have reported achieving improved
performances for zinc - manganese dioxide dry cells
[Leclanche cells] by using zinc perchlorate as the electrolyte
in place of zinc chloride. For example, significant improve-
ment in retention of capacity at high temperature is
claimed.113 Machat and Sohm114 found that a Leclanche-type
cell with an electrolyte containing 1.5 M zinc perchlorate
and saturated with zinc hydroxide provided 4-times the
capacity of a conventional Leclanche cell after aging 2
months at 60°. These same authors 115 also found that (1) cor-
rosion of the zinc anode in a Leclanche-type cell is greatly
reduced by replacing zinc chloride with zinc perchlorate
and [2] the capacity on both continuous and intermittant dis-
charge is improved without loss of other favorable proper-
ties. According to Watanabe and coworkers, 116 Leclanche-
type dry cells with zinc perchlorate and ammonium chloride
as electrolyte give good discharge characteristics under heavy
loads.

105
 A zinc - lead dioxide storage cell has been described117
that contains 40% perchloric acid and resists freezing even
at —56°.
Machat and Sohm118 obtained a patent for a zinc - mercuric
oxide battery containing zinc perchlorate electrolyte which
is suitable for use in electrical measuring instruments and
watches because of its freedom from gas formation.
Perchlorate salts are effective in alkaline electrolytes
for preventing passivation of primary and rechargeable
cells.119 For example, calcium perchlorate and sodium hy-
droxide are employed in an alkaline battery patented by
Berger and Dietlin.120 Tetrabutylammonium perchlorate is
employed in nonaqueous electrolyte solutions patented by
Katv and Saito.121
MisceJJaneous. Concentrated perchloric acid is em-
ployed in a carbon fuel cell patented by Becker.122 The cell
consists of a carbon cathode in contact with sulfuric acid, a
porous clay diaphragm, and a copper or carbon anode in
perchloric acid. It operates by oxidation of carbon at 20 to
90°.
A carbon fuel cell utilizing lithium perchlorate and ni-
trogen dioxide dissolved in a mixture of propylene carbonate
and acetonitrile as a nonaqueous electrolyte has been
patented by Schlaikjer.123
Lead perchlorate has been recommended as a highly
soluble salt for use in electroplating and refining of lead.124'125
Ozone can be electrogenerated from 40% perchloric
acid using platinum anodes refrigerated at —60 to —65°.
Yields and current efficiencies improve with increasing cur-
rent densities and decreasing temperatures. 126
Animal Feed Additives
Considerable attention has been devoted by Russian in-
vestigators to the use of ammonium, sodium, and potassium
perchlorates as stimulants for increasing the weight of farm
animals and poultry. According to Solun and coworkers, who
first reported on the subject,127128 the perchlorates exert a
significant thyrostatic effect when added to feed rations of
cattle and broilers that leads to increased weight gains of
up to 20%. The perchlorates are metabolized and completely
eliminated, mainly via urine, within 24 to 48 hours.

106
 Young bulls fed at a daily rate of 2,5 mg of ammonium
perchlorate were found to have gained 19 kg or 17.8% more
than controls after 90 days.128 Their thyroid weights did not
change significantly in this time, but some marked dif-
ferences in microstructure were observed. Similar studies re-
vealed that weight gains depend on feed type.130"132
Ammonium perchlorate administered to oxen in their
feed increased the deposition of internal fat by 39% and
caused a slight but noticeable effect on the aroma of the
meat.133 Ammonium perchlorate administered daily to bulls
in their feed over a 7-month period (2.5 mg NHUClCVkg body
weight) resulted in increases in body weights of 17.3 to
20.7%, meat yields of 9.4%, and fat yields of 22.4%.134 Am-
monium perchlorate in the diets of young rams and ewes led
to increased weight gains without affecting meat quality.135
The effect of ammonium perchlorate on nitrogen meta-
bolism in young chickens has been investigated by
Kurilova.138
Explosion hazards associated with storage and handling
of weight stimulators containing perchlorate salts can be de-
creased by addition of other salts, according to Yakimenko
and coworkers. 137
Miscellaneous
Anhydrous magnesium perchlorate and barium per-
chlorate serve as highly efficient drying agents for gases.
They are also useful for removing small amounts of polar
compounds from inert gases. Further details are described
in Chapter IV.
Potassium perchlorate is employed together with stron-
tium azide and boron in a gas-forming composition which
is suitable for the rapid inflation of automobile safety bags.
Ignition of the gas-forming composition can be rapidly
achieved in the event of an automobile collision by use of
ignition mixtures of boron, zirconium, aluminum, and/or
magnesium with lithium perchlorate, sodium perchlorate,
potassium perchlorate, ammonium perchlorate, and/or po-
tassium nitrate. 138
A patented composition to ignite charcoal briquets con-
tains 15 to 16% potassium perchlorate, 4 to 6% metal nitrates,
and 78 to 8 1 % carbonized matter.139
Use of iron and potassium perchlorate mixtures in the

107
 form of pellets to serve as heat generating material or
"thermal reservoir pellets" for initiating action of thermal
batteries has been described by Bush.140'141
Lead borate glass solder, used for connecting the front
and funnel parts of color television tubes, can be selectively
dissolved in 0.4 to 1.6 M perchloric acid at or below 85°.142
It can also be dissolved in 0.5 to 6 M perchloric acid at 21°
in the presence of ultrasound. 143
Perchloric acid has been employed as a drilling agent
and disinfectant in the treatment of dental canals,144 as a
component in etching solutions for production of semi-
conductor devices,145 and in solutions to passify or prevent
corrosion of iron14G and steel surfaces.147
Addition of small amounts of sodium perchlorate to
cooling-lubricants employed in machining of metals is re-
ported to reduce friction and increase the wear life of high-
speed cutters. 148
An antipollution system for internal combustion engines
has been patented based upon use of oxygen generated by
heating lithium perchlorate in the presence of manganese
dioxide catalyst.149 By elimination of air, oxides of nitrogen
are eliminated.
Use of aluminum perchlorate in styptic preparations 150
and of diazonium perchlorate in diazo copying composi-
tions151 have been described in recent patents.
Quaternary ammonium perchlorates have been found
useful as sensitizing agents for photographic emulsions.152'153

108
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143. H. Merk and E. Rischauer, German Patent 2,218,103 (1973).
144. J. M. Caillot and C. L. Knapp. French Patent 2,131,816 (1972).
145. R. Schlmmer, ]. Nicolussi, and H. Gesing, German Patent 2,430,560
(1976).
146. L. Kiss, L. DoNgoc, and M. L. Varsanyi, Acte Chim. (Bmhipest).
70, 313 (1971).
147. P. F. Drozhzin and L. I. Osipova. U. S. S. R. Patent 134,673 (1961).
148. A. A. Kut'kov, I. P. Golyubin, and A. S. Serebrenikov, Tr. Novo-
cherkassk. Politekh. Inst.. 191, 110 (1969); Chem. Abstr., 72,
46566 (1970).
149. E. J. Cettin, ]. P. Pappas, and S. E. Lager, II. S. Patent 3,709,203
(1973).
150. ). M. Blass, British Patent 1,416,296 (1975).
151. K. Ikari, H. Kato, and H. Tomii, Japanese Patent 76,135,527 (1976).
152. C. F. H. Allen and C. V. Wilson, U. S. Patent 2,299,782 (1942).
153. B. H. Carroll, U. S. Patent 2,271,623 (1942).

113
 CHAPTER VI
CHEMICAL ANALYSIS OF PERCHLORATES
Detection Methods
Methods for the detection of perchlorate ions generally
involve precipitation, extraction, or decomposition of per-
chlorate to give characteristic, identifiable products. Invari-
ably, the methods lack specificity unless either prior sepa-
rations are performed or reliable confirmatory tests are
conducted.
Milligram amounts of perchlorate (5 to 50 mgj can be de-
tected and identified with considerable confidence by precipi-
tation with tetraphenylarsonium chloride, followed by in-
frared examination of the precipitate in a potassium bromide
disc.1 Although perrhenate, permanganate, periodate, per-
technate, molybdate, chromate, and tungstate also form
insoluble tetraphenylarsonium salts, none exhibit an intense
absorption band at 9.12 fim (1096 cm'1) as does the perchlo-
rate. For greater sensitivity, particularly when the perchlorate
concentration is too low to provide sufficient precipitate,
perrhenate can be added to the sample so that tetraphenylar-
sonium perchlorate will coprecipitate and thus be collected
by the precipitate of tetraphenylarsonium perrhenate. 2 This
method is reportedly 3 capable of a sensitivity of 0.03 fj.g of
perchlorate per ml, with a coefficient of variation of approxi-
mately 10%.
Diphenyliodonium chloride can be used in place of tetra-
phenylarsonium chloride to precipitate perchlorate and other
polyatomic anions for identification by infrared spectropho-
tometry. 4 In a KBr disc, the perchlorate salt is identified by
strong absorption bands in the 1080 to 1140 and 620 to 634
cm"1 regions.
For the detection of microgram amounts of perchlorate
the method of Feigl and Goldstein 5 is suitable, provided that
nitrates and halates are absent or first removed by fuming
with concentrated hydrochloric acid. The test sample is fused
with cadmium chloride while testing the gas phase above the
melt for the appearance of chlorine, which will cause a test
paper impregnated with thio-Michler's ketone to turn blue or
a paper with a fluorescein-bromide mixture to turn red. De-
tection of chlorine by either test paper indicates the presence

114
of at least 1 or 5/*.g, respectively, of perchlorate in the test
sample. A similar method, involving fusion with zinc chlor-
ide, was described by Gooch and Kreider6 in which 50 ju.g of
perchlorate was detected.
To distinguish between perchlorate and some 26 other
anions the systematic qualitative analysis scheme of Belcher
and Weisz7 as extended by Hayes 8 can be followed. The
final supernatant, remaining after the sequence of separations
described by Belcher and Weisz, is treated with a slight
excess of 4 N sulfuric acid to remove barium ions. Addition of
4 drops of saturated zinc sulfate and 1 drop of 0.03 per cent
methylene blue to 1 drop of the barium ion-free supernatant
test solution will result in a color change to purple-red if
perchlorate is present.
Simple precipitation tests with various cations, although
subject to many interferences, are useful for indicating the
absence or possible presence of perchlorates. These include
tests with methylene blue {free of zinc salts] to give a violet
precipitate, 9 ' 10 potassium chloride to precipitate perchlorate
from ethanol-water solution,11 rubidium chloride and potas-
sium permanganate to yield a colored precipitate of rubidium
perchlorate, 12 strychnine sulfate,10 tetraphenylarsonium chlo-
ride, 13 and copper(II) sulfate-pyridine solution.14
Perchlorates may also be detected by fusion with sodium
carbonate to yield chlorides which yield silver chloride on
treatment with silver nitrate. 10 Chlorates and chlorides inter-
fere.
A colorimetric detection procedure 15 is suitable for the
detection of 3 jug or more of perchlorate but suffers from
interferences by trace amounts of iodide, thiocyanate, perio-
date, perrhenate, and cetain metal ions as well. The test is
based upon the extraction of perchlorate into methyl isobutyl
ketone in the form of an ion-association complex with azide
and a copper(II] complex of 2,2'-bipyridine. Extractability of
the yellow copper(II) complex is dependent on the presence
of both azide and perchlorate anions. The limit of detection
can be lowered to 0.15 jug of perchlorate by applying the
Weisz ring-oven technique to the extract. 15
Separation Methods
Because so many possibilities exist for interference, the
most reliable methods for the identification or determination

115
of perchlorates in complex mixtures are those that incorpo-
rate some separation step to either remove suspected inter-
ferences or isolate the perchlorate before measurement. Chro-
matographic separations have proven to be highly effective
in this regard. Solvent extraction is often useful but not
always entirely effective, Separation by systematic precipi-
tation procedures, although feasible,7'8 are rarely employed.
Partition chromatography has proven to be one of the
most effective chromatographic methods for resolving mix-
tures of chloride, chlorite, chlorate, and perchlorate. For
example, Harrison and Rosenblatt 16 employed paper partition
chromatography to successfuly resolve mixtures of these
anions on Whatman No. 1 filter paper. Starting with as little
as 10 /xg of each ion, they obtained well-defined chromato-
grams in about four hours using a mobile phase consisting of
a mixture of isopropyl alcohol, water, pyridine, and concen-
trated ammonium hydroxide [15:2:2:2]. The following Rf
values were observed for the anions: CI", 0.25; CIO2", 0.34;
CIO3', 0.50; and CICV, 0.65. The perchlorate band was located
by spraying with 0.2 per cent aqueous methylene blue solu-
tion.
Thin-layer partition chromatography was used by Pes-
chke17 to separate halites, halates, and perhalates on silica gel
and modified silica-alumina layers. A mixture of n-butanol,
acetone, concentrated ammonium hydroxide, and water
{40:50:10:5} served as eluent. Chromatograms with a specific
color for each ion were developed by spraying with a solution
of 0.93 g of aniline and 1.66 g of o-phthalic acid in 100 ml of
water-saturated n-butanol followed by heating at 130° for
20 min. Thielemann 13 employed a similar chromatographic
procedure to separate and identify the oxy-anions of chlorine
produced from chlorine dioxide treatment of drinking water.
Using specially activated silica gel plates (impregnated with
2:1 acetone-formamide and dried at 120°), an elution mixture
of 60:25:5:10 n-butanol-acetone-ammonium hydroxide-pyri-
dine, and spraying with bromocresol purple to detect the
spots under ultraviolet light, he obtained the following Rf
values: CIO2", 0.36; CIO", 0.38; ClCV, 0.68, and ClOt", 0.64.
Chloride interfered in the detection of hypochlorite.
Starobinets and Mechkovskii 19 found that perchlorate
could be separated from chloride and chlorate ions by par-

116
tition chromatography on a cation-exchange resin using
aqueous acetone as eluent.
The separation of chloride, chlorite, chlorate, and per-
chlorate by ion exchange chromatography has been described
by Boyd and Larson.20 The ions were eluted from Dowex-1,
in the order named, using potassium bicarbonate solution
for the first three and sodium fluoroborate solution for per-
chlorate, the most strongly held ion of the four.
Loach21 took advantage of the strong affinity of per-
chlorate ion for ion exchange resins of the quaternary am-
monium type to isolate and concentrate perchlorate ions from
plant and animal extracts. Perchlorate was recovered from
the columns, after washing with water, by elution with 1 M
ammonium trichloroacetate.
The chromatographic behaviors of halate and perhalate
ions on strong-base anion exchange resin paper (nitrate ion
form) have been described by Lederer and Sinibaldi.22 Their
study also included evaluation of the properties of the same
anions in paper electrophoresis and in thin-layer partition
chromatography. They concluded that many different sepa-
rations of one ion from others were practical by the dif-
ferent methods.
Extration of perchlorate ions from aqueous solution into
a liquid anion exchanger or into an immiscible solvent in the
form of an ion- association complex can serve as a means of
minimizing interferences as well as for converting the per-
chlorate into a measurable form. Although numerous extrac-
tion systems have been investigated, none provide as high a
degree of separation specificity as that afforded by the
chromatographic methods just described. The following
liquid anion exchangers are typical of those that have been
evaluated for perchlorate extractions: Amberlite LA-2 liquid
resin (a long-chain aliphatic amine), 23 tetrahexylammonium
picrate in methylene chloride,24 and trilaurylamine in various
water-immiscible solvents. 25 Of the many different extract-
able ion-association complexes that have been utilized in per-
chlorate separations, the following are typical: tetrabutyl-
phosphonium perchlorate, 26 tris(l,l n -phenanthroline]zinc(II)
perchlorate, 27 and dimethyldiaminophenazine perchlorate. 28
The extraction of perchloric acid and other strong
monobasic mineral acids by organic bases dissolved in

117
water-immiscible solvents has been extensively investigated.
Amines studied as extractants include tri-n-octylamine,2!'"'!l
hexadecylamine,11- and di-2-ethylhexylamine.,,,:i Extraction by
trioctylamine increases in the order HCl<HBr<HNOs,<
HCIO4. Extractants studied of the organophosphorus type
include tributyl phosphate, 84,35 trioctylphosphine oxide,3(i"aH
and diethylfN.N-dibutylcarbamoylJphosphonate. 3 " Other ex-
tractants investigated include di(n-octyl)arsenic acid,40 tris(n-
octyl)arsine oxide,41 di-n-heptyl sulfoxide,42 and nitroben-
zene,43 For acid concentrations less than 4 N, the order of
extractability from aqueous solution by 1 M di-n-heptyl sulf-
oxide in trichloroethane solution was HNO:t> HC104>>
HCI =- HsSO-i.

Reduction of Perchlorates
Various methods have been described for the reduction
of perchlorates as a means for either their direct determi-
nation by titrimetric redox methods or their indirect determi-
nation by gravimetric or titrimetric determination of the
liberated chloride. The reduction methods are of three types:
thermal decomposition, wet chemical, and electrochemical.
With few exceptions, they require the use of appropriate
catalyst or high temperature in order to be practical for
analytical purposes.
Although perchlorates can be decomposed simply by
heating at sufficiently high temperatures, it is more common
and prudent to decompose them by fusion with other sub-
stances or by cautiously heating their mixtures with suitable
reductants. Lamb and Marden 44 employed fusion with sodium
carbonate in glass test tubes to reduce perchlorates to chlor-
ide prior to gravimetric determination of the resulting chlor-
ide. Loss of chloride was prevented by the use of two plugs
of asbestos wool inserted in the test tubes, one near the top
and the other in the middle. Lenher and Tosterud 45 found
that they could reduce perchlorates rapidly by heating the
samples in a porcelain crucible in the presence of manganese
dioxide at 600 to 700°; however, they recommended that the
method of Lamb and Marden be used for greater accuracy.
Dobroserdov and Erdmann4(i reported obtaining high ac-
curacy with the Lamb and Marden method, provided that
the heating was carried out in hard glass tubes and not too

118
intensely nor prolonged. Reduction of perchlorates to chlor-
ides can be accomplished rapidly by fusion with either
sodium nitrate in a nickel crucible47 or with sodium peroxide
in a Parr bomb. 48 Meldrum and coworkers 49 employed a stain-
less steel bomb to explosively decompose perchlorate in
smokeless powder prior to determining the resulting chloride
by the Volhard method. Scharrer 50 reduced perchlorate to
chloride by heating with copper powder in a covered crucible.
Sodium oxalate has also been employed as a reductant in
thermal decomposition of perchlorate to chloride.51
Numerous wet chemical methods for the reduction of
perchlorate to chlorine or to chloride have been described.
Reduction with titanium[II] has been most extensively
studied, first by Rothmund 52 in 1909, then by Meldrum and
coworkers,4i> followed by Burns and Muraca/' 3 The latter
employed osmium tetroxide as a catalyst for the reduction
of perchlorate to chloride by titanous chloride. Back-titration
of excess titanium(II) with standard iron(III) served to de-
termine the perchlorate. Use of oxygen-free or inert atmos-
pheres is necessary to avoid air oxidation of titanium. Others
who have employed titanium(II) as a reductant for per-
chlorate include Vil'yamovich,*4 Peebles, 55 Schnell,50 Eagles,57
and Huskens and Gaty/' 8 Reduction of perchlorate with fer-
rous oxide in sodium hydroxide was studied by Sjollema,511
with starch in concentrated sulfuric acid by Willard and
Thompson,"" and with sulfur in sulfuric acid by Durand. 61
Molybdenum catalysis of the reduction of perchlorate to
chloride ions by stannous salts or zinc metal has been studied
extensively by Haight,,i2 who concluded that molybdenum(IV)
was the only catalytic species, Laitinen and Rechnitz 03 quan-
titatively reduced perchlorate to chloride with liquid cad-
mium amalgam in the presence of sodium molybdate catalyst.
Crowell and coworkers"4'05 found that osmium(IV) and ru-
thenium[III) ions exert catalytic effects on the reduction of
perchloric acid by hydrobromic acid in dilute solutions at
room temperature and at 100". Tbe redox reaction was
stoichiometric, yielding chloride ions and bromine. Tungsten
carbide was found to catalyze the reduction of perchlorate
ions by molecular hydrogen.™
Electrochemical reduction of perchlorate to chloride ions
can be accomplished in concentrated sulfuric acid and in

119
mixtures of sulfuric and nitric acid; however, considerable
time is required for quantitative reduction.'" At a mercury
cathode the reduction is much more rapid when catalyzed by
the presence of molybdenum(IV). Rechnitz and Laitmen""
found that molybdenumflV] was reduced by a molyb-
denumfV) dimer which was then reoxidized by perchlorate.
Polarographic studies by Haight"2 indicated that molyb-
denum(IV) was the active catalytic species. The reduction
of perchlorate to chloride ions at tungsten carbide electrodes
by electrochemically generated molecular hydrogen was re-
ported by Vertes and Horanyi.™

Determination of Perchlorate
To assure reliability of results the determination of per-
chlorate in complex mixtures should involve some prior sep-
aration procedure. None of the chemical reactions upon which
the measurements are based possess sufficient selectivity to
be free of interferences. The most direct, and perhaps the
most reliable, determinations are those that are based upon
precipitation of perchlorate as an insoluble salt. Precipitation
can be monitored by radiometric or spectrophotometric
means or measured gravimetrically or titrimetrically. Meth-
ods based upon the determination of either the reduction
products of perchlorate (chloride or chlorine) or the excess
reducing agent after reduction of perchlorate are indirect
and more susceptible to error and interferences.
Photometric methods for perchlorate are commonly
based upon precipitation or extraction of perchlorate ions in
association with a colored cationic species, followed by
measurement of the absorbance due to either the cationic
species combined with the perchlorate ion in the new phase
or that remaining in excess in the original phase. These in-
direct methods, although very sensitive, are highly subject
to interferences. Under optimum conditions, the relative
random error involved in photometric determination of per-
chlorate may be 1 to 2%.
In the absence of interferences, perchlorate ion concen-
tration can be estimated most rapidly by potentiometry with
appropriate ion-selective electrodes. Relative errors, depend-
ing on concentration levels, are commonly in the 1 to 10%
range.

120
 Gravimetric methods. At one time perchlorate was
commonly precipitated as potassium perchlorate for gravi-
metric purposes, even though addition of a large amount of
ethanol was necessary to minimize solubility losses.70"73
Quantitative precipitation of perchlorate now can be achieved
most effectively through use of certain organic precipitants,
notably those which possess or give rise to large bulky
cations. Nitron, as described by Storm,74 was one of the first
of such reagents to be employed. It has been used by vari-
ous investigators as a gravimetric reagent for perchlorate,75"'11
however its solutions are relatively unstable and nitrate
interferes. 80
The tetraphenyl derivatives of trivalent phosphorus, ar-
senic, and antimony salts have been found to be practical
gravimetric reagents for the determination of perchlorate.
Investigations by Willard and coworkers81-82 have demon-
strated that a large number of different anions are quanti-
tatively precipitated by these reagents. Of the three, tetra-
phenylarsonium chloride forms a more insoluble perchlorate
than either tetraphenylphosphonium 83 or tetraphenylstibo-
nium chloride, 84 and thus is the most commonly used. Ac-
cording to Loach85 a 1 M sodium chloride medium provides
desireable influences on the solubility and filterability of
tetraphenylarsonium perchlorate precipitates. Gravimetric
procedures for perchlorate using tetraphenylarsonium chlo-
ride have been described by various workers.81'83'86"88 Use of
tetraphenylphosphonium chloride for the gravimetric de-
termination of perchlorate in the presence of nitrate 80 and in
the presence of chlorate ions 90 has also been described.
Tetra-n-pentylammonium bromide has proven to be a
suitable gravimetric reagent for the determination of per-
chlorate in amounts ranging from 2 to 50 mg.fll Equivalent
amounts of chlorate do not interfere. Interference from larger
amounts of chlorate can be eliminated by prior reduction with
sodium bisulfite, an excess of which is without interference.
Various other organic compounds have been proposed
or utilized as perchlorate anion precipitants for gravimetry.
These include methylene blue,92,93 a-phenyl~/3-diethylamino-
ethyl-p-nitrobenzoate, 94 1,2,4,6-tetraphenylpyridinium ace-
tate,"1" and 3,5,6,8-tetramethyl-l,10-phenanthroline. !,,i
Gravimetric determination of perchlorate can also be

121
effected indirectly after reduction of chloride by appropriate
methods. Lamb and Marden 44 employed decomposition by
sodium carbonate fusion followed by precipitation of chlor-
ide as silver chloride for the gravimetric determination of
various perchlorate compounds. Dobroserdov and Erdman46
reported obtaining satisfactory results by the same procedure.
Willard and Thompson 60 employed starch and a concentrated
solution of sulfuric acid for the reduction, distillation for
the separation of the resulting hydrochloric acid, and pre-
cipitation with silver nitrate to determine perchlorate in-
directly. Various other methods for the reduction of per-
chlorates prior to gravimetric determination of the resulting
chloride are cited in the previous section.
Titrimetric methods. Potentiometric,97'99 thermometric, 100
and conductometric 101 methods have been described for the
precipitation titration of millimolar amounts of perchlorate
with tetraphenylarsonium chloride. Various interferences are
commonly encountered, especially by large and symmetrical
anions.
An amperometric titration procedure based on tetra-
phenylstibonium sulfate as the precipitation titrant for per-
chlorate has been described by Morris.84 Contrary to many
classical titrimetric procedures for perchlorate, the method
is claimed to be free of interferences from chloride, chlorate,
nitrate, phosphate, and sulfate.
Bolliger102 reported determination of as little as 1 mg of
perchlorate by treatment with a measured amount of methy-
lene blue, removal of the precipitate by filtration, and titra-
tion of the excess methylene blue with picric acid.
Willard and Smith81 determined perchlorate iodometric-
ally by precipitation as tetraphenylarsonium perchlorate fol-
lowed by titration with standard iodine.
Several back-titration methods based upon the reduction
of perchlorate followed by titrimetric determination of the
excess reductant have proven satisfactory for the determina-
tion of perchlorates. Burns and Muraca53 employed titanous
chloride as reductant, osmium tetroxide as catalyst, and
standard ferric ammonium sulfate for back titrating the ex-
cess titanium(III) reductant. A similar procedure was de-
scribed by Meldrum and coworkers. 4 " Aravamudan and
Krishnan'":i found that iron(II) can quantitatively reduce per-

122
 chlorate to chloride if sufficiently high temperature and con-
centrated sulfuric acid are employed. They determined per-
chlorate by reduction of the perchlorate with ferrous sulfate
in 11 to 12 M sulfuric acid at 150 to 155° for about 15 min.,
followed by back-titration of the excess iron(II) with stand-
ard permanganate after cooling and dilution with water.
Nitrate interference could be corrected for by a separate de-
termination under milder conditions such that only nitrate
was reduced by iron(II).
Numerous methods have been described for the determi-
nation of perchlorate based upon the titration of chloride ob-
tained on reduction of the perchlorate. They differ in choice
of reductant or decomposition method and in chloride ion
precipitation technique. Sjollema5y employed reduction by
iron(II) in sodium hydroxide followed by the Volhard method
for chloride. Kurz, Kober, and Berl104 fused the perchlorate
with an excess of sodium nitrite at 500° for 1.5 hr. to obtain
chloride which was subsequently determined by the Volhard
or the potentiometric method. A similar reduction method
was employed by Yamasaki and coworkers 47 to determine
perchlorate in a mixture of chloride, hypochlorite, chlorite,
chlorate, and perchlorate. Chloride before and after reduc-
tion was determiend by amperometric titration with standard
silver nitrate. Fusion with sodium peroxide and potassium
superoxide followed by the Volhard method was used by
Riolo and Occhipinti 105 to analyze mixtures of perchlorates
and chlorates. Reduction of perchlorate was complete only
in the presence of organic matter. Simonyi and Tokar10G em-
ployed reduction with Raney nickel alloy in boiling, dilute
sodium hydroxide to reduce chlorate without interference
from perchlorate. For the reduction of both chlorate and per-
chlorate they fused the sample with potassium hydroxide in
the presence of ethyl alcohol. Chloride was determined by
the Volhard method, and the difference in amounts of chlo-
ride produced by the two reduction methods enabled calcu-
lation of the perchlorate content of the sample. Alley and
Dykes107 determined perchlorates in pyrotechnics by reduc-
tion with titanium hydride in sulfuric acid followed by poten-
tiometric titration of the resulting chloride with silver nitrate.
Matrixes impenetrable by 1:3 sulfuric acid and the presence
of cellulose nitrate interfered. The oxygen flask method was

123
 used by Secor, Ricci, and White 108 to decompose organic and
inorganic samples for the microdetermination of perchlorate,
yielding chloride which was titrated coulometrically.
A simple and accurate method for the determination of
perchlorate in mixtures of other chlorine-containing com-
ponents consists of the total decomposition of one portion of
sample to chloride by ignition with ammonium chloride and
the selective reduction of another portion such that all chlo-
rine-containing components except perchlorate are converted
to chlorides. The difference in chloride content obtained by
the two methods represents chloride obtained from perchlo-
rates, from which the perchlorate concentration is deter-
mined. Procedures based on this method have been described
by Scott100 and others. 4 ''• ItHMl0,m
Titration of perchloric acid. Determination of per-
chloric acid with standard base requires special techniques if
the sample contains other acids or metal ions that interfere
through formation of stable hydroxy species. Mixtures of
perchloric and nitric acids have been differentiated by po-
tentiometric titration with potassium hydroxide using ethyl
methyl ketone as solvent.112 Perchloric acid has also been de-
termined without interference from ni'ric acid by a high
frequency titration method using a solvent medium of acetic
acid and a titrant of diphenylguanadine or pyridine in acetic
acid. m Sansoni" 4 reported obtaining 6 distinct and points in
the conductometric titration of a mixture of perchloric, p-
methylsulfonic, sulfuric, nitric, and trichloroacetic acids in a
solution of acetic acid containing 1% water using 0.5 M
sodium acetate as titrant.
Potentiometric titration of aqueous mixtures of per-
chloric and acetic acids with sodium hydroxide yields two
breaks; however, the first is rather poorly defined and limits
the precision of the perchloric acid determination. If the solu-
tion is diluted with an equal volume of dioxane prior to the
titration, the sharpness of the first inflection or break is
considerably enhanced, permitting much improved accuracy
in the differentiation and determination of the individual
acids.115
The determination of perchloric acid in cotton acetyl-
ation bath solutions has been described by Buras, Cooper, and
Cruz.116 Containing either acetic acid - acetic anhydride -

124
perchloric acid or acetic acid - water - perchloric acid, the
three-component baths were analyzed for perchloric acid
content by titration with standard potassium biphthalate after
dilution with a mixture of acetic anhydride and acetic acid to
assure that any water present was consumed by the acetic
anhydride. p-Naphtholbezein was used as indicator.
Certain metal perchlorates have been analyzed by elec-
trodeposition of their aqueous solutions followed by titration
with standard base of the metal-free perchloric acid solution
that results from the electrolysis. This unique method yields
the metal content by electrogravimetry and the perchlorate
content by acidimetry. It has been applied with varying de-
grees of success to the analysis of the perchlorates of
cobalt,117 nickel,117 iron,118 and lead.75 Interference from anions
other than perchlorate and incompleteness of electrodeposi-
tion constitute rather serious limitations for the method.
Photometric methods. Most of the photometric methods
for perchlorate are indirect, involving spectral measurements
either in the visible or ultraviolet region of the absorbance
due to cationic species in association with the sought-for
perchlorate after some suitable ion-association reaction or
separation. The only direct photometric methods are those
based upon measurement of the characteristic infrared
absorption of perchlorate ions.
For the determination of perchlorate by infrared spectro-
photometry it is necessary to separate the perchlorate anion
from other strongly absorbing components in the sample.
Aqueous samples can be evaporated to dryness or extracted
with a suitable non-polar solvent, one which is relatively
non-absorbing in the infrared region of interest. Such pro-
cedures, although time consuming, also serve to concentrate
the perchlorate and thus enhance the sensitivity of the de-
termination. Solid samples are prepared for spectral measure-
ment in the form of pressed disks, either as obtained or
mixed with potassium bromide. Absorbance measurements
are made at a wavelength corresponding to the maximum of
any one of the three intense, broad bands due to perchlorate
in the 8.5 to 10 jam (1000 to 1150 cm1} region.
Loach21 devised a sensitive and highly selective method
for the determination of trace amounts of perchlorate in
plant and animal tissues utilizing infrared spectrophotometry

125
for the final measurement step. Organic samples were
chopped into small pieces and boiled with distilled water to
extract the perchlorate content, After filtration, clarification,
and concentration of the filtrate by evaporation, the perchlo-
rate was taken up on a column of Dowex 2-X8 ion exchange
resin in the trichloroacetate form. After washing, the per-
chlorate was eluted from the column with 1 M ammonium
trichloroacetate and precipitated by addition of tetraphenyl-
arsonium chloride solution and potassium perrhenate solu-
tion. Complete recovery of the perchlorate was achieved by
its coprecipitation with the tetraphenylarsonium perrhenate.
The precipitate was collected, dried in vacuo, ground with
potassium bromide, pressed into a disk, and its infrared
spectrum recorded. The following results for 2-kg samples
were typical, with concentrations expressed as ju-g KClCu per
kg of sample: sea water, 0; silver beet, 9-12; urine, 21-28; and
cabbage, 9.
Briggs and coworkers" 51 employed infrared spectrophoto-
metry to determine potassium perchlorate in potassium chlo-
rate. To avoid the risk of explosion when compressing
samples into transparent disks, the samples were first boiled
with hydrochloric acid to decompose the chlorate and then
evaporated to dryness. A relative standard deviation of 5%
for 0.3% KCICM was found for the method.
Dolinski and Wilson" determined perchloric acid in the
presence of nitric, sulfuric, and chloric acids by extracting
the aqueous solution with a carbon disulfide solution of
Amberlite LA-2 liquid resin followed by infrared spectro-
photometric analysis of the extract. The absorbance of per-
chloric acid in the extracts, followed Beer's law up to a
concentration of 3 mg/ml.
Many indirect photometric methods are based upon
liquid-liquid extraction of perchlorate in the form of an
ion-association complex with a suitably absorbing cationic
species or dye. Measurement of the absorbance of the ex-
tracted dye-perchlorate complex thus provides an empirical
basis for determining the perchlorate. In general, these
methods provide high sensitivity but low selectivity. One of
the more commonly employed methods 120 of this type uses
methylene blue as the cationic species and chloroform as the
extracting solvent. The methylene blue-perchlorate complex

126
in chloroform obeys Beer's law at 655 nm for perchlorate con-
centrations up to 0.5 ppm. A similar method, using 1,2-
dichloroethane and extraction from an acidified solution
rather than from pH 5-7, was described by Iwasaki and co-
workers.1"1 Interferences from many different anions and
certain cations can be minimized by masking with mercuric
ions or by an aqueous sulfuric acid backwash, Applied to
the determination of perchlorate in sea water samples, the
method indicated that perchlorate was either absent or less
than 0.05 ppm.
Various workers122"126 have investigated brillant green for
the extraction-photometric determination of trace amounts of
perchlorate. The ion-association complex is extracted from
aqueous solution of pH 4-7 into benzene where its molar
absorptivity is 9.4 x 104 at 640 nm.123'124 Nile blue127 and
neutral red28 have also been studied as extraction-photometric
reagents for perchlorate. The molar absorptivity of the
neutral red-perchlorate in nitrobenzene is 9.39 x 104, and
that of the nile blue complex in 1,2-dichlorobenzene is 8.1xl04
at 650 nm.
Certain complex metal cations have proven to be sensi-
tive extraction-photometric reagents for the indirect determi-
nation of microamounts of perchlorate. The ion association
complex of ferroin perchlorate has been extracted into n-
propylnitrile 128 and into nitrobenzene' 29 to determine perchlo-
rate at the ppm-level. Similarly, the extraction of neocuproine
copper{I)-perchlorate into ethyl acetate has served as the
basis for the indirect determination of perchlorate concentra-
tions from 0.5 to 5.0 ppm.130,131 The perchlorate salt of the
copper(I) complex of 6-methylpicolinaldehyde azine is ex-
tractable into either chloroform or isobutyl methyl ketone,
enabling determination of perchlorate indirectly by measure-
ment of the absorbance of 480 nm, with a relative error of
± 1 . 3 % at a level of 12 ppm.132
For the determination of perchlorate in potassium chlo-
rate Fogg and coworkers 26 extracted the perchlorate into
o-dichlorobenzene in the form of its ion association complex
with tetrabutylphosphonium ion. On washing the organic
extract with an aqueous solution of ferric thiocyanate, the
perchlorate ion was replaced yielding a colored iron{IIi}-
thiocyanate, tetrabutylphosphonium complex for which an

127
absorbance measurement at 510 nm proved proportional to
the original perchlorate ion concentration. Beer's law was
followed up to 500 /xg of perchlorate. Interferences from
iodide, nitrate, chlorate, and bromide were eliminated by
evaporating the sample to dryness, first from hydrochloric
acid and then water solution before analysis.
Irving and Damodaran 133 determined perchlorate in the
presence of other halogen species in aqueous solution by
means of a colored liquid anion exchanger, a solution of tetra-
hexylammonium tetranitritodiamminocobaltate(III) in xylene
and iso-butyl methyl ketone. The cobaltate complex anion
exchanged quantitatively with perchlorate and was de-
termined photometrically by the increase in absorbance in
the aqueous solution or by the decrease in the absorbance in
the organic phase at 353 nm. Interferences by chlorate, hypo-
chlorite, chlorite, nitrite, and nitrate were eliminated by
fuming the samples with hydrochloric acid. A similar indirect
photometric method for perchlorate, using tetrahexylam-
monium picrate as the liquid anion exchanger, was described
by Gustavi and Kylberg.24
Precipitation of perchlorate with a measured quantity of
a colored precipitant followed by filtration and photometric
determination of the excess is a general indirect method for
perchlorate that has been employed by a number of investi-
gators. Nabor and Rammachandran 134 used methylene blue
as the precipitant. A correction was necessary for the solubil-
ity of the precipitate. To determine perchlorate in blood Kurz
and Renner135 precipitated the perchlorate with nitron and
determined the excess nitron by its absorbance at 490 nm in
alcoholic sodium hydroxide solution. Shahine and Khamis13"
also used nitron for precipitation of perchlorate; however,
they determined the excess by extraction with a cobalt thio-
cyanate complex anion into cyclohexane-carbon tetrachloride
and photometric determination of the absorbance of the
complex at 625 nm. Excess nitron, after precipitation of per-
chlorate, has also been determined by an iodometric pro-
cedure.137
\ simple indirect method which involves neither precipi-
tation nor extraction was devised by Trautwein and Guyon.138
Based upon the quantitative interference of perchlorate in
the spectrophotometric determination of rhenium with a-

128
 furildrioxime, their method proved capable of a sensitivity
comparable to that of the ferroin method. m Interference,
however, was quantitative for the following ions: S2O3'2,
Cu +2 , V +5 , NO"*, NO'2, and U02+.
Ion-selective electrodes. In recent years a variety of
membrane type electrodes have been developed that exhibit
highly selective, Nernstian response toward perchlorate ion
activities. Most perchlorate ion selective electrodes are of
the liquid-membrane type, although a few precipitate-type
(heterogeneous) and at least one solid state type have been
devised. Details of the theory and technology of ion selective
electrodes in general have been extensively reviewed.139*142
The most significant advantages afforded by ion selective
electrodes are simplicity and speed of application for mea-
surement or monitoring of individual ion activities. However,
a number of limitations and precautions should be recognized
in their use: (1) they are not specific, only selective towards
a particular ion or group of ions, so interferences are com-
mon; (2) they measure activities not concentration of ionized
species, thus complexation and ionic strength influence the
results; and (3) their response frequently depends upon pH,
age, condition, temperature, and time. For direct concentra-
tion determination their use requires empirical calibration
curves, frequent standardization, and reliable standard solu-
tions. However, for use as indicating electrodes in poten-
tiometric precipitation or complexation titrations these fac-
tors are no longer important. Ion selective electrodes are
ideally suited for following concentration changes in chem-
ical processes, reations, kinetics, and titrations.
One of the most extensively studied and widely used per-
chlorate ion electrode is the commercially available Orion
Model 92-81 electrode. This liquid-membrane type electrode
has a useful linear response to perchlorate ion over the molar
concentration range of 10"1 to 10"35 in aqueous systems of pH
to 11. It exhibits reasonable selectivity for perchlorate over
common anions such as nitrate, bicarbonate, and the halides.
143.144 j t s r e s p 0 n s e time increases with decreasing perchlorate
ion concentration, requiring at least 15 sec. to attain a con-
stant reading for 10"5 molar solutions.145 In addition to direct
determination of perchlorate ion concentration, the Orion
perchlorate electrode has found use in determination of solu-

129
 bility products of slightly soluble perchlorates 144 and in po-
tentiometric titrations' 4 ' 1 It has also proven useful for the
determination of vicinal glycols by an automatic reaction
rate method14" and for catalytic titrations involving periodate
indicator reactions. 147
Various liquid membrane electrodes have been studied
that are similar to the Orion perchlorate electrode in con-
struction and composition. Ishibashi and Kohara148 reported
that a nitrobenzene membrane containing bathoferroin per-
chlorate showed excellent selectivity for perchlorate over
nitrate or iodide and gave linear Nernstian response up to
about 10"" M perchlorate. Reinsfelder and Schultz149 de-
termined selectivity coefficients for common inorganic ions
of electrodes of the liquid membrane type containing trisfl,
10-phenanthroline)iron[II) in nitrobenzene and tris(4,7-di-
phenyl-l-,10-phenanthroline)iron(II) in nitrobenzene, chloro-
form, or n-amyl alcohol. Selectivity coefficients were rela-
tively independent of membrane composition with a sequence
of selectivity as follows: PF<r> C104'> SCN'~ I"~ BF4">
N03"> Br"> Cl", Rohm and Guilbault150 prepared a mixture of
polyvinylchloride and the commercially available liquid ex-
changer from the Orion perchlorate electrode to construct
disks and coated-wire electrodes which proved similar in
response to the Orion electrode. The construction and
characterization of liquid-membrane electrodes containing
the perchlorate salt of tris(l,10-phenanthroline)iron(II) and
that of the analogous nickel(II] chelate in nitrobenzene have
been reviewed by Hopirtean et al.iS1
A liquid membrane type electrode based on the per-
chlorate salt of tetrakistriphenylphosphinesilver(I) has been
claimed to be superior to the Orion perchlorate electrode
with respect to applicability to basic solutions.152
Coetzee and Freiserir);Ma4 found that methyl tricapryl am-
monium salts [Aliquat 336S) in 1-decanol function effectively
as organic phase components in liquid-liquid membrane elec-
trodes for the determination of a variety of anions. The per-
chlorate electrode of this type exhibited linear potential re-
sponse over the concentration range 10"1 to 10"3 M perchlorate.
Tateda, Fritz, and Hani155 reported that electrodes prepared
using Aliquat 336 proved less favorable than the commerc-
ially available (Orion) perchlorate electrode. Other quater-

130
nary ammonium perchlorates have been explored extensively
as ion-exchangers for use in liquid membrane electrodes in
attempts to develop improved perchlorate electrodes. These
studies have produced practical but not superior electrodes.
Systems studied, in addition to Aliquat 336, include methyl-
cetylbenzylammonium salts in nitrobenzene fixed on char-
coal or graphite,156 tetraalkylammonium salts in ethyl
bromide,157'158 tetraoctylammonium perchlorate in toluene,159
and octadecyldimethylbenzylammonium perchlorate in nitro-
benzene.160 Quaternary phosphonium salts have also been
investigated for use in perchlorate selective membrane elec-
trodes. 161
A perchlorate ion-selective electrode based upon a liquid
membrane consisting of methylene blue perchlorate in nitro-
benzene has been reported to exhibit Nernstian response to-
wards perchlorate over the range 10"(i to 1 M perchlorate, in-
dependent of pH over the range 2 to 12.5.162 Similar results
were found for the electrode based on the tetrafluoroborate
salt of methylene blue in tetrachloroethane. 163 Brilliant green
perchlorate, another basic dye salt, has been examined as an
ion-exchanger for liquid membrane electrodes.164 A chloro-
benzene solution of this salt in natural rubber served as the
membrane for an electrode that gave a useful linear response
over the 0.1 to 0.001 M perchlorate range for solutions of pH
4 to B.165
An exceptionally sensitive perchlorate electrode has
been described by Sharp.186 Containing a liquid membrane
prepared from a solution of N-ethylbenzothiazole-2,2'-azavio-
lene perchlorate in 1,2-dichlorobenzene, the electrode ex-
hibited a response range of 1 to 10"*-5 M perchlorate, a useful
pH range from 1 to 12, and high selectivity for perchlorate
over all of the 10 anions tested. A solid-state electrode made
from the same radical ion salt has also been described by
Sharp,107 but its selectivity and response characteristics are
less favorable.
Several moderately sensitive perchlorate electrodes of
the heterogeneous (precipitate) type have been described
that possess the advantage of simplicity of fabrication. James
and Freiser163 prepared such an electrode by coating platinum
wire with a mixture of polyvinylchloride and the ion-associ-
ation complex between Aliquat 336S (a quaternary ammon-

131
ium cation] and perchlorate. Hiro, Tenaka, and Kawahara 196
coated graphite with a suspension of cadmium perchlorate in
Urushi (a natural lacquer). Ishibashi and coworkers 170 impreg-
nated tetradecyldimethylammonium perchlorate in a plastic
matrix to construct an electrode that exhibited linear re-
sponse down to 10"' — 10"4 M perchlorate.
Miscellaneous methods. The radiochemical method of
isotopic dilution, using chlorine-36 labeled potassium per-
chlorate, was employed by Johannesson171 to investigate the
question of the occurence of perchlorate in natural sea water.
None was found within the limits of sensitivity and accuracy
of the method [± 4.2 mg KCIO4 per 250ml, for the 95% con-
fidence limits).
An indirect radiometric procedure for the determination
of perchlorate after extraction as an ion association complex
with radioactive zinc complexed with 1,10-phenanthroline
has been described by Shigematsu et air7 Extracted into a
measured volume of nitrobenzene from pH 7 aqueous solu-
tion, the activity of zinc-65 is measured by a scintillation
counter and is proportional to the amount of perchlorate ion
in the original sample. For 10-ml samples a concentration
range for perchlorate of 2 x 10"7 to 10" M is suitable for analy-
tical purposes. Many different anions interfere but not chlo-
ride, sulfate, or phosphate.
Bishop and Evans 173 devised a rate-measurement method
for the determination of perchlorate at low concentrations
based on the homogeneous reaction kinetics of titanium(III)
with perchlorate.
Mixtures of chloride, chlorate, and perchlorate were
analyzed by Gnanasekeran et al.m by ion-exchange chrom-
atography. The anions were collected on a column of a special
resin (developed from cashew nut shell liquor, treated with
hydrochloric acid, and washed with water), eluted in their
acid form with water, and titrated with standard sodium
hydroxide. Replicate determinations agreed to within ± 1%.
Analysis of Perchlorate Compounds and Mixtures
Specifications and procedures for testing the quality of
sodium, potassium, and magnesium perchlorates and of 70%
perchloric acid as reagents have been compiled by Rosin.174
Requirements for ACS reagent grade specification and test
procedures have been established for lithium perchlorate 175
and for potassium and magnesium perchlorates as well as for

132
 60% and 70% perchloric acid solutions.176 Procedural details
include assay of the perchlorate, determination of residue on
ignition, insoluble matter, chloride, chlorate, nitrate, sulfate,
calcium, heavy metals (as lead), and iron.
A rapid method for checking the concentration of per-
chloric acid solutions (from 0 to 72% by weight) has been
described by McLean and Pearson177 based on measurement
of refractive index. Smith and Lamplough178 proposed the
same procedure, however the accuracy of their data was
questioned by McLean and Pearson. Refractive indices at 20 :
and 30° are compiled in Table 8 (in Chapter II] as a function
of concentration. Density determinations have also been
advocated for checking perchloric acid concentrations.17'1
Elemental analysis of organometallic perchlorates for
carbon, hydrogen, and nitrogen involves considerable risk
of explosion when the samples are ignited. It has been re-
ported that explosive compounds can be safely analyzed by
first mixing the perchlorate sample with cobalt oxide and/or
copper oxide prior to heating in the combustion furnace. A
disadvantage of the method, however, is that nitrogen values
tend to be about 0.2 to 0.3% low.180
Trace metal determinations by atomic absorption spec-
trophotometry are subject to interference from perchloric
acid. Oguro181 found that perchloric acid (0.001-0.10 M) en-
hanced the absorption of Mg, Ca, Mo, Cr, and V and
decreased the absorption of Fe, Ni, and Co in an air-acetylene
flame. Others have reported perchloric acid interference in
flameless atomic absorption spectroscopy also.182'183 Thus
in the analysis of perchlorates and perchloric acid solutions
for trace metal content, it is imperative that standard solu-
tions used in preparing calibration curves be of the same per-
chlorate composition as the unknown if reliable results are
to be obtained.
Other than the above, relatively few procedures have
been published for analysis of impurities in perchlorates. A
method for the determination of free acid in aqueous solu-
tions of uranium(IV) perchlorate has been described by
Schmid and Junger.184 Several papers by Russian chemists on
the determination of titanium in perchloric acid and per-
chlorate salts have been cited by title in Chemical Ab-
stracts. 1B5-187

133
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[19771-

139
 CHAPTER VII
BIOLOGICAL EFFECTS OF PERCHLORATES
Although ammonium, potassium, and sodium perchlo-
ates exert a number of interesting biological effects in both
plants and animals, their toxicities are relatively slight.
Nothing in the long history of their manufacture points to
any appreciable biological hazard for a worker exposed to
them for prolonged periods. According to Levens,1 who crit-
ically review the literature prior to 1960 on the biological
action of perchlorates, sodium and potassium perchlorates
would be considered to be "slightly toxic" to animals under
the classification scheme of Hodge and Sterner. 2 Several
other perchlorates, however, are known to be more than
slightly toxic. Still others, the great majority, remain to be
investigated. In fact, much remains to be learned about the
biological modes of action of even the sodium and potassium
perchlorates.

ANIMAL STUDIES
Physiological Effects
The effects of perchlorate ions on the thyroid gland and
its function in various animals have been extensively investi-
gated. Wyngaarden et al? found that perchlorate produced
quantitative discharge of iodide from rat thyroids within 15
min. It proved to be about 10 times more effective than thio-
cyanate and approximately 300 times more than nitrate. The
blocking effect of perchlorate on the collection and retention
of iodide by the thyroid may also extend to other iodide-
concentrating mechanisms of the animal body.4 Rats treated
with perchlorate for 17 days developed goiters which were
hyperplastic and low in iodine. 3 Wyngaarden and coworkers 5
also observed that increasing amounts of iodide, thiocyanate,
perchlorate, and nitrate produced marked reductions in the
ability of the thyroid glands of rats to concentrate radioiodide.
Perchlorate proved to be the most potent anion and nitrate
the least. Halmi and Stuelke 8 found that subcutaneous in-
jection of 100 mg or more of sodium perchlorate prevented
active uptake of iodine-131 by thyroid glands of rats treated
with propythiouracil, a substance known to prevent hor-

140
mone synthesis involving the combination of tyrosyl groups
and iodide ions. According to Kruskemper and Kleinsorg,7'8
potassium perchlorate interferes with the synthesis of thy-
roxine in rats and mice. Results obtained by Breslavskii and
Simon" appear to support this opinion, although contradictory
to the conclusions of Wyngaarden and coworkers/'
Kleinsorg and Kruskemper 10 studied the effects of thy-
roxine and antithyroid substances (potassium perchlorate
and methylthiouracil) on the serum proteins in rats. Potas-
sium perchlorate fed simultaneously with thyroxine pre-
vented decreases in serum protein. When fed alone at 100
or 250 mg/kg/day, potassium perchlorate raised the (3-
globulin content by 68 percent in 14 days and 94 percent in
28 days, while the albumin/globulin ratio was lowered from
1.8 to 1.17.
Selivanova11 found that ammonium perchlorate is as
potent as either sodium or potassium perchlorate for in-
hibiting thyroid accumulation of iodine-131 in rats, mice,
and rabbits. Tests with NH43fiC104 revealed no cumulative
toxicity for dogs. Mice gained weight on an ammonium
perchlorate treatment of less than 260 mg/kg and lost weight
at greater doses.
Kinetics of the distribution of radioactive perchlorate
in rat and guinea pig thyroid glands have been studied by
Chow and Woodbury. n A kinetic analysis of iodide transport
in dog thyroid slices by Rochman et a]. 13 indicated that 1 m M
sodium perchlorate solution inhibited the influx of iodide
into the follicles and discharged the trapped iodine-131 into
the media with an increased efflux rate.
Mikhailov14 found that sheep which received 100 mg of
ammonium perchlorate per kg weight or 2 mg/kg/day for 6
months with their feed showed decreased protein-bound
iodide in the blood, decreased inorganic blood iodide, and
increased urinary and fecal iodide excretion. The compound
acted directly on the thyroid gland to alter its function.
Yamada and Nakamura 15 observed that perchlorate exerts a
contracting effect on thyroid glands.
Rousset, Orgiazzi, and Mornexlfi reported that per-
chlorate ions enhanced thyroid responsiveness in mice to
thyrotropin, human chorionic gonadtropin, and long active
thyroid stimulators.

141
 Harden and coworkers 17,18 found that perchlorate ions
exert an inhibitory effect on the parotid saliva gland. After
administration of perchlorate the labial saliva/plasma ratio
for iodine-131 decreased 47 per cent from the control value.
Nevelichuk1" observed antiarrhythmic properties for po-
tassium perchlorate. Given orally to rats each day for 3 days,
potassium perchlorate {20 to 910 mg/kg) prevented ventric-
ular fibrillation induced by calcium chloride [i.v.] and de-
creased the cardiotoxic action of calcium chloride. No re-
lationship between the antiarrhythmic effect and antithy-
roidal activity was found.
The use of perchlorate as weight stimulators for poultry
and farm animals is discussed in Chapter V.

Perchlorate Metabolism
Various studies indicate that perchlorate salts are elimi-
nated in the urine by most animals without chemical reduc-
tion of the perchlorate ion. Kerry and Rost20 recovered per-
chlorate unchanged in the urine of rabbits administered
sodium perchlorate intravenously. Rabuteau 21 recovered po-
tassium perchlorate unchanged in urine of patients treated
for malaria. Quantitative recovery of sodium perchlorate
from a human subject was reported by Durand.22 Appearance
of sodium perchlorate in urine occured 10 min. after injec-
tion. Maximum concentration was observed at about the
third hour, and 95 per cent elimination had occured after
48 hr. Eichler" observed similar results for the elimination
of potassium perchlorate taken orally by man.
Durand22 discovered sodium perchlorate distributed
throughout the entire rabbit body soon after intravenous or
intramuscular injection or oral administration. Determined
over the time interval of 20 to 130 min. after administration,
concentrations of sodium perchlorate were highest in the
ovaries, adrenal glands, and urine. Next highest concentra-
tions were found in the spleen, gall bladder, and intestinal
mucosa. Concentrations were lower in the heart, liver, kid-
neys, lungs, brain, blood, gastric mucosa, muscle, bone and
testes.
Eichler and Hackenthal 24 found that perchlorate elimi-
nation by rats was faster following high than smaller doses,
especially during the early hours after administration. Up

142
to 95 to 97 per cent of the perchlorate was excreted in the
urine within 60 hr. In a study of potassium perchlorate
metabolism in the rat, Goldman and Stanbury 25 found no
significant concentration of perchlorate in the kidney, liver,
brain, or spleen. They found that the rates of disappearance
of the 3BC10"4 radionuclide from the plasma and thyroid and
the rate of appearance in the urine were exponential, with
half-times of approximately 20 hr. These rates are similar
to the disposal rates of 131I found by others.
A comparative study of iodine and potassium perchlorate
metabolism in the laying hen revealed that the concentration
and distribution properties of the two were very similar.2*' It
was concluded that perchlorate may be of value in blocking
uptake of iodide into eggs.

Pharmacological Studies
The earliest pharmacological study of perchlorates was
reported in 1868 by Rabuteau 21 concerning the use of po-
tassium perchlorate as a therapeutic agent against malaria.
Some years later Sabbatini 27 completed a comparative study
of the pharmacology of oxygenated chlorine compounds in
which he found that toxicity generally decreases with de-
creasing oxidizing properties. The toxicity of perchlorate,
however, proved to be an exception to the rule, being less
than that of chlorite but more than that of chlorate. Sabbatini
attributed the distinctiveness of perchlorate toxicity to an
immobilizaion or diminution effect on potassium ions.
Pharmacodynamic effects of sodium perchlorate on
excised muscles and muscular tissues were investigated by
Messini.28'2" Its addition to the perfusion fluid produced
muscle contractions that could be relieved by small amounts
of potassium chloride, smaller than that required to inhibit
muscle excitability. Washing with sodium chloride enabled
the sequence of contraction and relaxation to be repeated.
Reversing the order of treatment produced the reverse
change, i.e. increased muscle tone from high doses of potas-
sium chloride was diminished by sodium perchlorate. Messini
attributed these phenomena to a disturbance of potassium
ion equilibrium within the muscle.
From investigations on the reactions of transversely
striated muscle (frog sartorius) to sodium and potassium per-

143
 chlorates and other salts, Boehm30,31 concluded that the con-
tractions produced by perchlorate and fluoborate anions were
both related to precipitation of albumin within the muscle.
Adding to the controversy regarding the cause of mus-
cular contractions induced by perchlorates, Eichler23 re-
ported that perchlorate-poisoned frog hearts were cured by
addition of calcium ion but not by potassium, indicating that
the poisoning was more akin to an excess rather than a de-
ficiency of potassium. On the other hand, both Cartolari3*
and Spagnol3'1 reported pharmacodynamic studies leading to
the conclusion that the contraction effect of perchlorate on
muscles is produced by decreasing the potassium ion balance.
Hypothyroidism has been successfully controlled in
human patients by treatment with oral doses of potassium
perchlorate.34"17 In general, symptoms improved, basal me-
tabolisms fell to nearly normal, body weights increased, and
blood cholesterol levels increased to normal values. No sig-
nificant changes in the formed elements of the blood nor
evidence of liver damage were observed for patients treated
for as long as 52 weeks. 35
Bru'gel38 found that perchlorate inhibited accumulation
of iodine-131 in human subjects to the extent that less than
1 per cent of the normal amount was taken up. According
to Lewitus,:!" perchlorate ions compete with inorganic iodide
in the trapping mechanism of the thyroid gland, leading to
decerased throxine synthesis.

Toxicology and Pathology

Symptoms of perchlorate poisoning in various animals
have been described by Kerry and Rost.20 Frogs injected with
0.015 to 0.030 g of sodium perchlorate exhibited fibrillation,
twitching, and strong contraction of transversely striated
muscle radiating from the injection site. Reflex excitability
was greatly heightened, and heart action was gradually
paralyzed. Symptoms were less severe when sodium per-
chlorate was administered orally, and the frogs recovered
completely from oral doses of less than 0.15 g. Similar symp-
toms (convulsions, paralysis, and rigidity) were observed for
sodium perchlorate poisoning in mice, rats, guinea pigs,
pigeons, cats, and dogs. Subcutaneous doses of 0.025 g of
sodium perchlorate were lethal to mice, 0.22 g to rats, and

144
 1.35 g to guinea pigs. Relatively large oral doses (1 to 2 g)
proved fatal to a guinea pig and a rabbit but not to a cat or
a dog.
Kahane40 found that goldfish were unaffected after 3
days in 0.1 per cent sodium perchlorate solution, but that ex-
posure to higher concentrations [0.2 to 2 per cent) resulted
in their death. No evidence of reduction of perchlorate to
either chlorate or chloride was found.
Leeches immersed in 0.5 per cent sodium perchlorate
solution were unharmed after 5 days; two-thirds died after
2 days in 1 per cent solution; and all expired in a 4 per cent
solution within 1 hr." Tadpoles were affected after 48 hr
in 0.1 per cent sodium perchlorate solution, and in 0.2 per
cent solution all died within 36 hr.'22
The toxicity of sodium perchlorate in rabbits has been
investigated extensively by Kahane 40 and by Durand.22 Autop-
sies performed on rabbits subjected to repeated injection of
sodium perchlorate revealed no evidence of changes in the
heart, kidney, or intestines. However, necrosis of the tissue
at injection sites, caseation of the liver, and hepatization of
the lungs were evident.
According to Selivanova, 11 the minimum lethal dose of
ammonium perchlorate is 3.5 g/kg for rats, 2 g/kg for mice
and 0.75 g/kg for rabbits.
The general toxic effects of perchloric acid when ad-
ministered orally or sub-cutaneously to rats, mice, and dogs
have been described in detail by Selivanova and coworkers. 41
In addition to its acidic irritating effect, perchloric acid pro-
duces a specific antithyroid action and induces abnormalities
in hepatic, renal, cardiovascular, and hemapoietic functions
when administered internally.
Perchloric acid is reported to cause dermatitis if per-
mitted to act directly on the skin in sufficient amount or con-
centration for a sufficient length of time.42,43 Although its
corrosive or irritating action is not as immediate nor quite as
damaging as concentrated nitric, sulfuric, or hydrochloric
acid, due caution should be followed in handling it to avoid
contacting it with the eyes and skin. In the event of acci-
dental contact the affected area should be washed imme-
diately with large quantities of water. Prompt medical at-
tention should also be received after any contact with the eye.

145
 The inhalation toxicity in rats and mice of the com-
bustion products of rocket propellants containing various
amounts of perchlorate has been investigated by Feinsilver,
Mac Namee, McGrath, and Oberst.44
Toxicological effects of perchloryl fluoride in rodents
and dogs have been studied by Kunkel and coworkers.45,46
The subject also has been reviewed in detail by Levens.1

BACTERIAL AND MICRO-ORGANISM

STUDIES
Chemical reduction of perchlorate to chloride ions by
several species of heterotrophic bacteria has been demon-
strated with the use of chlorine-36 labelled perchlorate by
Hackenthal and coworkers. 47 The reduction is strictly de-
pendent upon the presence of bacterial nitrate reductase.
Perchlorate reduction by Staphylococcus epidermidis in com-
plex media is inhibited by nitrate, and nitrate reduction of
resting cells of Bacillus cereus is inhibited by perchlorate.
These results suggest that nitrate and perchlorate ions may
be substrates to the same enzyme, nitrate reductase. The
rates of reduction of both anions are equally dependent on
pH, incubation temperature, and several electron donor and
cofactors.48 Moreover, both activities are inhibited by cy-
anide, azide, and 2,4-dinitrophenol, as well as by one an-
other in a competitive manner. Perchlorate is reduced first
to chlorate and then to chloride by cell-free extracts obtained
from nitrate-adapted cells of B. cereus. Hackenthal and Ar-
babzadeh 49 also found that nitrate reduction in cell-free ex-
tracts from B. cereus and Pseudomonas aeruginosa is com-
petitively inhibited by perchlorate, thiocyanate, selenocya-
nate, chlorate, rhenate, and azide.
The toxicity of aqueous solutions of sodium perchlorate
to Escherichia coli and other micro-organisms has been de-
scribed by Durand.22 Development of E. coli and of Staphylo-
coccus pyogenes aureus is prevented by concentrations of
about 2.5 to 3.0 and 7.5 to 10 per cent, respectively. For the
mold Sterigmatocystis nigra development is retarded by a
concentration of 1.3 and arrested by a 10 per cent solution
of sodium perchlorate.
Sodium perchlorate has been found to affect both the

146
respiration and enzymatic activity of soil micro-organisms. 50
In either incubated soil or pure cultures of soil bacteria,
sodium perchlorate inhibited the respiratory activity and de-
creased the number of ammonifying, nitrifying, and deni-
trifying bacteria. According to Karki and Kasier,50 the per-
chlorate acted in 2 ways: by altering metabolites that are
toxic to the cells and by competing with nitrate for the
nitrate-reductase A enzyme.
Ammonium perchlorate solutions of concentrations 0.5
to 2 mg/ml are reported to have cytostatic or lethal effects
on Paramecium caudatum, Saccharomyces cerevisiae, and
Candida tropicaJis.51

PLANT STUDIES
Numerous investigations have been devoted to charac-
terizing the action of perchlorate on plants, prompted in large
measure by findings in 1896 by Sjollema52 that Chile salt-
peter contains potassium perchlorate as a contaminate (in
amounts ranging from 0 to 6.79 per cent) which can cause
damage to plants. The early studies clearly demonstrated
that certain levels of sodium or potassium perchlorate were
harmful to plants; however, some doubt arose as to whether
or not Chile saltpeter normally contains sufficient per-
chlorate to be injurious to plants if used in fertilizers.53'54
Sjollema55'56 pointed out that the sporadic nature and irre-
producible results observed for crop damage attributable to
use of Chile saltpeter could be accounted for by the fact
that potassium perchlorate is not uniformly distributed in
the mineral but occurs in pockets. He also observed that cer-
tain suppliers had begun (in about 1892) to recycle the salt-
peter mother liquor excessively, which could account for
a build-up of perchlorate content and the sudden appearance
of damage to crops. In 1914, Maschhaupt 57 reported finding
a maximum perchlorate content of 1.5 per cent in crude and
about 1 per cent in refined saltpeter; in fertilization ex-
periments these levels inhibited plant growth, confirming
the findings of Sjollema. More recently, Tollenaar and Mar-
tin58 found that potassium perchlorate in Chilean nitrate was
responsible for stunted soybean plants in Chile. The plants
were 50 to 67 per cent reduced in size and bore strongly
crumpled leaves with burned tips. Clearly there is a need

147
for careful quality control in processing and selecting of
Chilean saltpeter for use in fertilizers.
Considering the possible implications to marine vegeta-
tion, a 1958 report'" that perchlorate had been found at 10 to
1000 ppm concentrations in sea water from various localities
attracted considerable attention. Extensive and thorough
testing by others'-""'14 failed to substantiate the claim that
perchlorate is an important constituent of sea water. Green-
halgh and RileyG1 attributed the alleged finding of perchlorate
to chloride ion interference in the analytical method.
The major action of moderate concentrations of per-
chlorates on plants is inhibition of growth. Seed germination
is also retarded or terminated. Various studies substantiate
these general conclusions, at least with regard to plant
species studied thus far. Sjollemar'2 observed that rye plants
grown in pots containing as little as 50 mg of sodium or
potassium perchlorate were stunted and bent and bore yellow
leaves. Exposure to perchlorate delayed germination of rye
seeds and resulted in abnormal embryos. Growth retardation
of rye and oats by perchlorates was reported by Maerckerfi5
and by Caluwe.1'1' Lauffs''7 found small amounts of perchlorate
to be beneficial to growth but large amounts to be toxic to
wheat plants. Root hairs of wheat plants grown for 8 days
in dilute perchlorate solution were deformed.07 Exposure of
young wheat and rye plants to 0.2 to 1.0 per cent potassium
perchlorate for only a few hours, however, had little effect
on root hairs.lia Ekdahl(i!t confirmed these findings by similar
experiments with root hairs of young wheat plants grown in
water culture. He also found that the toxicity of various
chlorine compounds decreased in the following order: hypo-
chlorite, chlorite> perchlorate> chlorate> chloride. Aberg™
observed, however, greater toxicity for chlorate than for
perchlorate in studying the effects of a number of salts on
young wheat plants grown in nutrient solution.
The toxicity of perchlorates vary with the type of plant,
culture or soil in which the plant is grown, and lighting con-
ditions to which the plant is exposed. The relative toxicity of
perchlorate to chlorate apparently may vary also, or even be
reversed, from one set of conditions to another. Weaver 71
found that perchlorate was more toxic to Biloxi soybean
plants grown in water culture than to those grown in sand; he

148
 found the opposite to be true for chlorate. Aberg70 observed
that root injuries to young wheat plants occurred when they
were grown in nutrient solution containing 0.5 mmol/I po-
tassium perchlorate, but only for lightgrown plants. Those
grown in darkness under similar conditions withstood dam-
age. Cook found that perchlorate is more toxic than chlorate
when sprayed on plants grown in soil,72 even though in solu-
tion culture the lethal dose to the plants is about the same
(0.25 per cent) for both anions. 73 Perchlorate and chlorate
both caused mottling of leaf foliage in young tomato plants,
but with different characteristics, in experiments reported
by Owen.74
Alvisi and Orabona 75 observed that solutions of po-
tassium perchlorate were initially harmful but eventually
beneficial to bean plants. With ammonium perchlorate, the
anion appeared to be decomposed by the bean plant, while
potassium perchlorate was unchanged. Similarly, Penicilium
giaucum reduced ammonium perchlorate but not potassium
perchlorate. Treatment of a 0.2 per cent potassium perchlo-
rate solution with pepsin, papain, or diastase for 1 to 6 days
at room temperature failed to produce detectable chloride ion
levels.
Weiske 54 found that germination of grains and vegetables
was retarded by exposure of the seeds to potassium perchlo-
rate. Germination of the seeds of Pisum sativum was only
slightly decreased by soaking 24 hr in 1 per cent potassium
perchlorate, according to Vandevelde. 78 In experiments con-
ducted by Durand,22 seeds of poppy, lentil and flax germi-
nated almost normally in 0.2 per cent sodium perchlorate,
but they were affected appreciably in 0.5 per cent solution.

149
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93 (1949).
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22. |. Durand, Bull. soc. chim. biol., 20, 423 (1938).
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27. L. Sabbatini, Arch. sci. med., 51.144 (1887).
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34. f. B. St anbury and J. B. Wyngaarden, Metabolism., 1, 533 (1952).
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36. H. Kleinsorg and H. L. Kruskemper, Naunyn-Schmiedebergs Arch.
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eases of the Skin., 3rd ed.. Lea & Febiger, Philadelphia, Pa., 1957,
pp. 44-7, 951.
43. F. A. Patty, "Industrial Hygiene and Toxicology," 2nd ed., Vol. 1,
Interscience, New York, N. Y., 1958, pp. 449-51.
44. L. Feinsilver, J. K. MacNamee, F. P. McGrath, and F. W. Oberst,
"Inhalation Toxicity of Combustion Products of Perchlorate Fuel
Propellants," Medical Division Research Report No. 20, Chemical
Corps, U. S, Army Chemical Warfare Laboratories, 1950.
45. A. Kunkel in "Transactions of the Symposium on Health Hazards
of Military Chemicals," K. VV. Jacobson, Ed., C W L Special Publi-
cation 2-10, U. S. Army Chemical Warfare Laboratories, August
1958, pp. 101-8.
46. E. Greene, R. Brougii, A. Kunkel. and W. Rineharl, "Toxicity of
Perchloryl Fluoride," C W L Technical Memorandum 26-5, U. S.
Army Chemical Warfare Laboratories, Jan. 1958.
47. E. Hackenthal, W. Mannheim, R. Hackenthal, and R. Becher, Bio-
chem. Pharmacol., 13, 195 (1964).
48. E. Hackenthal, Biochem. Pharmacol., 14,1313 (1965).
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(1966),
50. A. B. Karki and P, Kaiser, Rocz. Glahozn., 2tt, 213 (1975: Chem.
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55. B. Sjollema, Chem. Weekblad., I. 125 (1904).
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57. J. G. M a s c h h a u p t , Dircklie van den L n n d b o u w , 1914, I.
58. H. T o l l e n a a r a n d C. M a r t i n , P h y t o p a t h o l o g y , 62, 1164 (1972].
59. L. G. M. B a a s - B e c k i n g , A, D. H a l d a n e , a n d D. Izard, N a t u r e , 182,
645 (1958).
60. R. G r e e n h a l g h a n d J. P. Riley, N a t u r e , 187,1107 (1960).
61. R. G r e e n h a l g h a n d J. P. Riley, /. M a r i n e Bioi. A s s o c . U n i t e d King-
d o m . 41, 175 (1961).
62. J. K. ( o h a n n e s s o n , Ancii. C h e m . , 3 4 , 1 1 1 1 (1962).
03. K. W. L o a c h , N a t u r e , 196, 754 (1962).
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65. M a e r c k e r , L a m h v . Vers.-Stu., 52,34 (1899).
66. P. de C a l u w e , O r g a n n V e r e e n . Oudleeri Rijks L u n d b o u w s c h o o i , 12,
103 (1900).
67. A. Lauffs, D i s s e r t a t i o n at K o n i g s b e r g , 1902: cited by E k d a h l , Ref.
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08. G. S t i e h r . D i s s e r t a t i o n at Kiel, 1903; Cited b y E k d a h l . Ref. (69).
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70. B. A b e r g , A n n . Roy. Agr. Co)!. S w e d e n , 15, 37 (1948).
71. R. J. W e a v e r . P l a n t Physiol., 17.123 (1942).
72. W. H. Cook, C a n . /. R e s e a r c h , 15C, 229 (1937).
73. W. H. Cook, Can. /. R e s e a r c h . 15C, 520 (1937).
74. O. O w e n . I Porno/, ( t o r t . Sci.. 7, 270 (1929).
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76. A. J. V a n d e v e l d e , C h e m . W e e k h l a d . , 1, 410 (1904).

152
 CHAPTER VIII

SAFETY AND ENVIRONMENTAL

CONCERNS
Those who would employ perchloric acid or perchlorate
compounds to good advantage in any of the many and valued
applications known for these substances should be mindful
of the hazards involved, There resides in most perchlorates an
awesome oxidation power which, if unleashed in force upon
oxidizable matter, can give rise to conflagrations or fierce
explosions. Although calamity awaits anyone who would use
a source of power carelessly, there are sensible and safe
methods available to harness and control power for useful
purposes. In the case of perchlorates, the essential guidelines
for safeguarding against unleashed power and havoc involve
the following measures: (1} control of both temperature and
concentration, (2) exclusion of certain substances from con-
tact with perchlorates to prevent reactions that produce un-
stable perchlorates, and (3) minimization of amounts of ox-
idizable matter that can come into contact with any perchlo-
rate at any one time. The ferocity of the oxidation power of
perchlorates thus can be controlled as needed in any applica-
tion.
Specific precautions and safety measures in handling and
storing perchloric acid and perchlorates are summarized in
the following sections. For a clearer understanding of the
risks and safety measures, the properties of perchloric acid
and of the perchlorates should be reviewed in Chapters II,
III, and VII. Other sources1"4 may also be consulted.

Some Reported Perchlorate Explosions

Reports of early investigators attest to the fact that
anhydrous perchloric acid is not only exceedingly corrosive
to the skin and mucous membranes but explosively un-
stable.5'6 It should not be prepared except for special needs
nor stored except at low temperature and then only in small
quantities and for short periods.5"7 The monohydrate, al-
though considerably more stable and storable for long periods
in sealed glass ampoules without discloration, 7 can react ex-
plosively with organic matter.5'8

153
 Explosions have been reported for various perchloric
acid mixtures. The most tragic was that which caused the
death of 17 persons in Los Angeles, California, on February
20,1947. It apparently occurred as a result of the introduction
of a plastic holder into an electropolishing mixture consisting
of approximately 150 gal. of 68-72% perchloric acid and 20
gal. of acetic anhydride, contained in a stainless steel tank
(phenolic resin-coated) and fitted with a refrigeration system
that had been shut off.910 The unfortunate combination of
elevated temperature, concentrated perchloric acid, and re-
active organic matter resulted predictably in the inevitable.
Other mixtures for which explosions have been reported on
subjection to high temperature include those with aqueous
perchloric acid and the following: alcohols,11'12 hydrogen, 13
wood,14 acetic anhydride (containing more than 57% by vol-
ume HC104),1S bismuth metal,16 easily oxidizable organic ma-
terials, 1 ' sodium hypophosphite, 18 sulfoxides,19 phenol,19 and
various wet-oxidation mixtures which were not adequately
preoxidized by treatment with nitric acid20"22
It should be emphasized that aqueous perchloric acid
solutions of concentration 72% by weight or less are remark-
ably stable to heat and shock, provided that reactive material
is absent. For example, Dietz13 reported that boiling 72%
perchloric acid could not be detonated when primed with
mercury fulminate or lead azide. Introduction of an electrical
arc into the boiling acid produced only weak, local deflag-
rations.
Some of the most hazardous perchlorate compounds
reported are those that are most sensitive to heat or shock.
They include the highly explosive diazonium perchlorates, 23
silver perchlorate, 24 ' 25 hydrazine perchlorate, 2 ' fluorine per-
chlorate, 27 and perchlorate esters of aliphatic alcohols.28'31
Other perchlorate compounds for which explosions have
been reported on heating or by detonation are the following:
aniline perchlorate, 32 pyridine perchlorate, 33 benzyl perchlo-
rate,34 silver perchlorate-benzene complex,35 dimethylamine
perchlorate, 36 tris(ethylthio)cyclopropenyllum perchlorate, 37
and 1,3-dithiolium perchlorate. 38 The perchlorates cited here
serve only as examples; they certainly do not constitute a
complete list of hazardous perchlorates. It is both reasonable
and wise to assume that all organic perchlorate and most

154
heavy metal perchlorates are capable of explosive decomposi-
tion under certain conditions. There is no substitute for
cautious experimentation guided by diligent attention to
lessons from the past.
Precautions in Use and Handling of Perchloric Acid
Various agencies, organizations, and individuals have
formulated recommendations concerning the handling, stor-
age, and dispensing of aqueous perchloric acid. These include
the Manufacturing Chemists' Association,38 the National
Safety Council,40 the Factory Mutual Engineering Division,41
the Association of Casualty and Surety Companies,42 the As-
sociation of Official Agricultural Chemists, 43 the Analytical
Methods Committee of the Society of Analysts, 44 Harris, 45
Muse, 3 Gawen, 46 the National Fire Protection Association,47
and Everett and Graf.48 Many of these rules and recommenda-
tions are summarized below. However, as emphasized by
Levens,2 the prospective user should not rely unquestioning
only on rules but rather should "first become familiar with
the properties of perchloric acid, and then, for each case,
give the most careful consideration to the proposed experi-
mental conditions and operating procedure." Some of the
most important properties of perchloric acid in this regard are
therefore also cited in the list of precautions that follows:
1. Anhydrous perchloric acid is extremely unstable and
will explode spontaneously. Its preparation should be
avoided.
2. Aqueous perchloric acid in concentrations less than
85% is completely stable under ordinary storage conditions.
The concentrations normally supplied commercially are 60
and 72%
3. Boiling or evaporation of an aqueous solution of per-
chloric acid will not produce a dangerously high concentra-
tion because an azeotropic mixture results which at 760 mm
pressure contains 72.5% HCIO4 and boils at 203°C. If metallic
salts are present, however, the mixture should not be evapo-
rated to dryness over an open flame.
4. The dangerously explosive anhydrous acid may form
accidentally if the aqueous acid is subjected to strong dehy-
drating conditions, such as provided by hot, concentrated sul-
furic acid or phosphorus pentoxide. Special care is required

155
in performing analyses requiring the use of perchloric acid
with such agents,
5. Hot and concentrated perchloric acid {60 to 72%) is
a powerful oxidant, but it loses its oxidation strength when
cooled or diluted with water. Organic and combustible ma-
terials are readily oxidizable by hot, concentrated perchloric
acid, some explosively. The major source of danger in hand-
ling and storage of perchloric acid arises from the risk of
contact with oxidizable material accompanied or followed
anytime thereafter by exposure to high temperature.
6. Perchloric acid reacts with alcohols and certain other
organic compounds to form very unstable (shock or heat)
perchlorate esters. Dilute solutions pf these may be reason-
ably safe but extreme caution should be exercised to avoid
concentrating them by evaporation or other laboratory pro-
cedures.
7. When perchloric acid is used for the destruction of
organic matter, preliminary treatment with nitric acid or mix-
tures of nitric and perchloric acids is recommended. Nitric
acid moderates the reaction by oxidizing the more reactive
components at lower temperatures, before the perchloric acid
becomes sufficiently concentrated to begin its oxidation
action. Wet oxidation of heterocyclic nitrogen compounds,
fats, oils, and certain other reactive substances that are im-
miscible in perchloric acid require special treatment (see
Chapter IV).
8. When filter paper is used to collect precipitates from
solutions that contain perchloric acid [even if dilute), the filter
paper and precipitate should not be dried before thorough
washing with water or suitable aqueous solution to remove
all of the perchloric acid. Glass or porcelain filter crucibles
should be used in place of filter paper, and no precipitate
should be dried unless it is known to be stable at elevated
temperatures.
9. Safety equipment should be employed whenever per-
chloric acid is used in high concentration and/or at elevated
temperatures. Safety glasses, cleanliness, pipet bulbs, and
other routine safety measures are advisable at all times. A
special fume hood or fume eradicator (see below), portable
safety shield, and heat-resistant glassware should be em-

156
ployed for perchloric acid digestions and wet oxidation pro-
cedures. A fire extinguisher and safety shower should also
be at hand.
10. Perchloric acid spills should be diluted immediately
with water, taken up with swabs (preferably wool), and then
washed with generous amounts of water, The swabs should
also be washed with water before discarding. Swabs of cot-
ton or other cellulose material contaminated with perchloric
acid should be regarded as fire or explosion hazards if not
washed throughly with large amounts of water.
11. Perchloric acid should be kept apart from organic
chemicals and reducing substances, especially alcohol,
glycerol, and hypophosphites. Bottles of the acid should
stand in glass or porcelain dishes on ceramic or other non-
absorbing and non-inflammable benches, shelves, or supports.
Flooring and furnishings near the stored acid should be of
noncombustible and nonabsorbent materials also.
12. Organic matter [such as paper, wood, grease, plastic,
or cloth] which has been exposed to perchloric acid fumes
or solutions should be regarded as a fire or explosion hazard
until it has been washed thoroughly with water and tested. A
small sample (0.01 to 0.10g) placed on a hot-plate will indicate
whether or not the material has become unduly inflammable.
Precautions in Handling Perchlorates
No generally applicable set of recommendations have
been promulgated for the safe handling of perchlorates.
Several sources of information regarding hazards and pre-
cautions, however, are available for some. These include
articles and reports by Burton and Praill,1 Levens,2 Wolsey,4H
Davis,23 and Elliott and Brown.17 Safety requirements for per-
chlorate propellant mixtures have been summarized by War-
ren50 together with details of plant layout and design, equip-
ment, operating procedures, storage, and transportation.
Another source of information on handling and storage of
perchlorate propellant and explosive mixtures is the Ord-
nance Safety Manual of the U.S. Army Ordnance Corp.51
Covalent perchlorate compounds, such as organic per-
chlorates and certain heavy metal perchlorates, tend to be
dangerously explosive. The higher the degree of covalent
character, the greater the explosive tendency. Ionic perchlo-

157
rates, such as the alkali and alkaline earth metal perchlorates
and rare earth metal perchlorates, generally have much
higher stabilites. Other than these simple qualitative distinc-
tions, no simple guidelines exist for predicting the degree of
hazard associated with any given perchlorate, Each must be
evaluated separately in the context of its use. Several general
precautions, however, are applicable to all:
1. Exposure of the pure compounds to high temperature,
mechanical shock, or easily oxidized organic matter greatly
increases the risk of fire or explosion.
2. Perchlorate salts that have come into contact with
organic solvents should not be heated, subjected to shock, or
discarded without first adding a large amount of water.
3. When undertaking to prepare or use any perchlorate
for the first time one should review the literature for possible
risks and start with very small quantities to become acquaint-
ed with its properties and any hazards involved.
4. Anhydrous magnesium and barium perchlorates
should not be used as desiccants or dehydrating agents for
samples that contain volatile organic matter, such as ether
or alcohol. Spent desiccant should not be regenerated by
drying until first recrystallizing it from water to counter-
act any possibility that organic vapors have been absorbed.
5. Aqueous perchlorate solutions that contain heavy
metal ions or organic matter should not be evaporated to
dryness by heating. Extreme caution should be exercised in
scraping or agitating unknown solid products obtained from
perchlorate solutions.

Perchloric Acid Fume Hood and Eradicators

Digestions and other routine operations involving pro-
duction of perchloric acid fumes should be carried out in
specially designed hoods reserved solely for that purpose.
If a suitable hood is not available, special apparatus or a glass
fume-eradicator should be employed. No organic materials
should be stored in the hood. An oil bath should never be
used for heating. Electric heating mantles and hot plates of
the vapor proof type to protect against corrosion from acid
fumes should be used instead of open flames. The floor and
furnishings surrounding the hood or fume-eradicator should

158
be clean, non-absorbent, and constructed of inorganic ma-
terial.
The construction of fume hoods for perchloric acid
service has been described by various workers.52"50 A stainless
steel fume hood was described by Dieter and coworkers, 04 and
a polyvinyl chloride system was designed by Bcies.55 Most
incoporate the following recommendations, set forth by the
Factory Mutual Engineering Division:41

Fume hoods should be noncombustible, contructed

of metal or stoneware, and left unpainted or protected
with an inorganic coating such as porcelain. A water
spray is desirable for washing down the hood after per-
chloric acid fuming. Otherwise, a separate hood should
be reserved for perchloric acid use only. Stoneware
hoods equipped with water spray for washing down are
commercially available. Ducts must not be manifolded
and should take the shortest path to outdoors. Use an
electric hot plate or heating jacket, a steam bath, or an
electrically or steam-heated sand bath for heating re-
action vessels. Do not use gas flames or oil bath for
heating. Provide a generous supply of water-type fire
extinguishers.

Some special recommendations that should be consid-

ered in the design and construction of fume hoods for per-
chloric acid are the following:
1. Grouting material and cements for sealing the hood
should not contain organic constituents. For example, litharge
and glycerine cement should not be used. Pastes of Alundum
cement and sodium silicate or similar materials are satisfac-
tory.44
2. Bends, horizontal runs, and porous construction ma-
terials should be avoided in constructing the exhaust system.
A vertical exhaust system which will not collect dust or distil-
late is preferred. 3
3. A washdown system should be incorporated in the
ductwork (as well as the hood proper) to wash out all de-
posits daily.3
Recommendations as to how to dismantle perchlorates
contaminated lab exhaust systems have been given by

159
 Breysse and Lehman,57 together with recommendations for
the installation of safe systems. Wolsey58 has described a
simple test for perchlorates and perchloric acid in hoods and
chemical apparatus, based upon the reaction with methylene
blue to give a violet precipitate with perchlorate ion.
A portable, self-contained laboratory scrubber unit was
described by Silverman and First.59 Designed to be placed in a
standard lab hood, the unit incorporates a wet absorption
stage, a droplet eliminator, and a dry filter. It was reported
that up to 99.9% of the perchloric acid fumes from acid di-
gestions could be eliminated through use of the unit.
Smith designed a perchloric acid fume eradicator for
convenient use in the open laboratory.60'61 The apparatus is
illustrated in Figure 4.1 and is commercially available from
the G. Frederick Smith Chemical Company. It includes a
refluxing still head to prevent loss of spray from the hot acid,
a glass fume collector that surmounts and envelops the top
of the reaction flask and refluxing still head, and an aspirator
flask attached to the fume collector and filled with sodium
hydroxide solution to scrub the aspirated fumes which pass
thru it and out into the water aspirator. The fume eradicator
may be placed in position with ease and readily removed and
set aside, being supported by the weight of the aspirator
flask and its contents.
A modified distillation unit for performing perchloric
acid digestions without a special fume hood was described
by Griffen and Hocking.0" It is constructed from a regular
long-necked borosilicate glass Kjeldahl flask and fitted at
the top with a distilling head connected by a large polished
glass, spherical joint separated by a perchloric acid resistant
Telflon-covered silicone rubber o-ring. The condenser is
fed into a double-sphere anti-suckback drip tube which dis-
charges into 350-400 ml of cold water.

Environmental Pollution
Disposal of waste perchlorates and perchloric acid in
small quantities, such as expended in ordinary laboratory
procedures, is commonly accomplished by flushing the wastes
into sewer drains with generous quantities of water. This
expediency is much safer than accumulating the wastes for
special treatment and separate disposal. No adverse effects

160
 to plants or animals from small concentrations (up to 100 ppm
as KCIO4) of perchlorate have been reported. Biological ef-
fects of perchlorates are summarized in Chapter VII.
To the writer's knowledge, no problem has been en-
countered in sewage treatment plants arising from the dis-
posal of perchlorate wastes from ordinary laboratory opera-
tions. The disposal of large quantities of waste perchlorates,
however, is a different problem because concentrations of
perchlorate at the 0.1 to 1.0% and above levels are known to
affect certain aquatic life and soil micro-organisms. Here, the
safest and most expedient approach appears to be biodegra-
dation of the waste perchlorates before discharging into
streams or landfills.
A method for purification of industrial waste waters
from perchlorates and chlorates has been described in patents
by Koren'kov, Romanenko, Kuznetsov, and Voronov.63"65 The
perchlorates and chlorates are reduced anaerobically by a
micro-organism classified as Vibrio dechloraticans, isolated
from a mixture of industrial and domestic sewage contain-
ing perchlorates. 66 An example described in the patents fol-
lows. Industrial waste water containing 600 mg NH4CIO4/I
and clarified domestic sewage [BOD 180 mg/1) were pumped
into a bioreducer at rates such that the ammonium perchlor-
ate concentration and the biological oxygen demand of the
mixture was 100 mg/1 and 150 mg/1, respectively. An en-
riched culture of V. dechloraticans (Cuznesove B-1168) was
continuously pumped into the bioreducer so that its concen-
tration was 3.5 g/1 based on the biomass solids. During the
one-hour residence time of the culture and mixture in the
bioreducer the perchlorate was completely reduced to chlo-
ride. The rate of reduction of perchlorate was 28 mg/g bio-
mass solids-hour.
The rate of biodegradation and its effect on the environ-
ment has been studied in terms of growth, metabolic rate,
and total biomass of selected animal and plant species.(i7
Short term effects on selected organisms were investigated,
and a long term experiment was designed to assess the
changes incurred by ammonium perchlorate on nitrogen and
chloride contents of soil over a three-year period. 68

161
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163
 APPENDIX
PHYSICAL AND EQUILIBRIUM DATA FOR
BINARY, TERNARY, AND QUATERNARY
SYSTEMS OF PERCHLORATES
The sheer quantity of physical data available for the
many different mixed perchlorate systems precludes detailed
presentation in a monograph of this kind. Therefore, in
place of the data, a reasonably complete bibliography of the
published literature on the subject is presented in systematic
fashion in the tables which follow.
References are cited by number in the tables and iden-
tified in the list of references at the end of this Appendix.
For brevity, several of the more commonly cited journals are
abbreviated as follows:
Abbreviation Name 0/ Journal
UZY Uch. Zap., Yaroslav. Gos. Pedagog. Inst.
SNT Sb. Nauch. Tr. Yaroslov. Gas. Ped. In-t
ZNK Zh. Neorg. Khim.

164
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168. A. S. K a r n a u k h o v a n d A. V. B o g a c h e v , UZY, 1973, No. 120, 36-9.
169. A. K. D e v d a r i a n i , N. P . K o l o b o v , K. N. M a r e n i n a , Zh. Fiz. Kriim., 47,
1152, (1973).
170. V. F. T a r a k a n o v , UZY. 1976. N o . 154. 35-7.
171. A. S. K a r n a u k h o v a n d V. F. t a r a k a n o v . UZY, 1972, No. 103, 3-6.
172. A. D. G u s e v a , UZY. 1976, No. 154, 27-30.
173. V. P. Z a k h a r o v a , UZY, 1971, N o . 95, 98-100.
174. Y. A. G o r y u n o v , SNT, 1975, 46-9.
175. Y. A. G o r y u n o v , SNT, 1975, 8-11.
176. T. Y. A s h i k h m i n a a n d N. N. R u n o v , SNT, 1975, 403.
177. T. Y. A s h i k h m i n a . UZY, 1976, N o . 154. 35.

179
 INDEX
Acetyl perchlorate 55, 96, 97
Acidity functions 21
Actinide metal perchlorates 49, 133
Activity coefficients 12, 15, 38, 40, 42, 4B, 49, 167
Alkaline earth perchlorates
ammines 38
basic salts 38
hydrates 38-40
preparation 38-41
properties 38-41
solubilities 38, 39
structure 40
Alkylation reagents 99
Aluminum perchlorate 41-2, 104, 108
Ammonium and alkali metal perchlorates
animal feed additives 106-7
crystal structure 32
preparation 34-7
properties 32-8
solubilities 33
thermochemical data 33
Analysis of perchlorates 132-3
Animal feed additives 106-7
Antiarrhythmic effect 142
Antimalarial effect 143
Antimony perchlorate 43
Antithyroidal activity 142, 145
Atomic absorption spectrophotometry 133
Babcock method 87
Bacteriological effects 146-7, 161
Barium perchlorate 41
Benzoyl perchlorate 55
Beryllium perchlorate 38-9
Bethge apparatus 76-7
Biodegradation 161
Bismuth perchlorate 43
Boiling points 17, 24, 53, 54, 55
Bromine perchlorate 53
Butterfat determination 87
Cadmium perchlorate 48

181
Calcium perchlorate 40
Catalysts
acetylation 96-7
esterification 96
isomerization 98
miscellaneous 98
polymerization 97-8
rearrangement 98
Cerium perchlorates 48-9, 82
Chlorine heptoxide 27
Chlorine perchlorate 53
Chromic perchlorate 49
Chromyl perchlorate 49
Cobalt perchlorates 50
Conductance 12, 34, 38, 40, 41, 48, 49, 165-73
Crop damage 147
Cupric perchlorate 47
Decomposition of organic matter
animal products and tissues 71, 80, 81
apparatus for 75-6, 77, 78
biological material 80
cellulose 77
efficiencies of 72
foodstuff 75, 80
graded potentials 73
heterocyclics 72, 75-6
loss of volatile elements 69, 70-1
metal chelates 80-1
plant material 71, 75, 80
polymers 75-6
procedures 74-7
products 71-4
proteins 72, 75-6
recovery of elements 70-1
starch 74
synthetic products 75
tobacco 74
Dehydration reagents 87, 99, 107
Densities 12, 13, 14, 33, 35, 36, 38, 39, 41, 42, 48, 51, 133,
166, 168-73
Deproteinization 84-5

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Desiccants
magnesium perchlorate 85-7
magnesium perchlorate trihydrate 86
barium perchlorate 86
perchloric acid azeotrope 87
Detection of
perchlorates 114-5
steroids 88
sugars 88
Determination of
calcium 79, 80
mercury 80
metal chelates 81
nitrogen 77-9
perchlorates 120-32
potassium 80
sulfur 79
trace metals 80-1
Deuterium perchlorate 24
Diffusion coefficients 12
Disinfectant 108
Dissociation pressures 35, 36, 48, 52
Dissolution of
alloys 67-8
carbonates 68-9
clays 64, 67
metals 67-8
ores 67-8
products from 72-3
silicates 64, 67
soils 64, 67
steel 67
volatility losses 64-6
Electrolytes for
batteries and cells 103-6
electropolishing 102-3
Electrosyntheis 9
Elemental analysis 133
EPR spectra 50
Etching solutions 108
Ethyl perchlorate 1, 54, 55
Explosions reported 153-4

183
Explosives 100
Explosive compounds 154-5
Extraction reagents for
metal complexes 84
nucleic acids 84
nucleotides 84
phosphorus compounds 84
serotonin 84
Ferric perchlorate 50, 99, 107, 108
Ferrous perchlorate 50
Fluorine perchlorate 53
Fluorocarbon perchlorates 53
Freezing points 12, 24, 53
Fume eradicators 158-60
Fume hood 158-60
Gallium perchlorate 42
Gravimetric methods 121-2
Growth inhibitor 147-9
Halogen perchlorates 53
Handling and storage 157-8
Heat capacities 25, 33, 49
Heats of hydration 37, 38
Heats of solution 32, 33, 47, 49, 52
History 1
Hydrazine perchlorate 52
Hypothyroidism 144
Igniters for
charcoal briquets 107
heat generating materials 108
inflation of safety bags 107
Indium perchlorate 42
Infrared spectra 5, 23, 25, 34, 38, 40, 41, 42, 114, 126-6
Iodine perchlorate 53-4
Ion exchange chromatography 117, 132
Ion selective electrodes 129-32
Kjeldahl method 77-9
Kreider-Mathers synthesis 10
Lanthanide metal perchlorates 48-9
Lead perchlorate 42-3, 106
Lithium batteries 104-5
Lithium perchlorate 36, 89, 98, 99, 104-5, 106, 108
Magnesium batteries 105

184
Magnesium perchlorate 40, 85-6, 89, 105, 132
Magnetic susceptibilities 38, 39
Manganese perchlorates 50, 99
Manufacture of perchlorates 2
Melting points 22, 23, 40, 47, 50, 52, 53, 55, 166, 171
Mercury perchlorates 48, 83, 100
Metabolism 142-3
Metal perchlorates
group IA-IIA 32-41
group IB-1IB 47-48
group IIIA-VA 41-43
group IIIB and rare earths 48-9
group IVB-VIIB 49-50
group VIIIB 50-1
Methyl perchlorate 54
Molar refractions 12, 33, 41, 44, 49
Molar volumes 12, 38, 39, 44, 48, 166
Natural occurence 3, 126, 132, 148
Nickel perchlorate 50
Nitronium perchlorate 51
Nitrosyl perchlorate 51-2
NMR spectra 5, 6, 21, 34, 36, 37
Nonaqueous titrations 81, 124-5
Organic perchlorates 54-5
Osmotic coefficients 16, 40, 48, 167, 174
Palladium perchlorate 50-1
Partition chromatography 116
Passification of iron and steel 108
Pathology 144-6
Perchlorate ion
complexation 6
contact ion-pair 5, 33
detection 114-5
determination 120-32
effect on water structure 5
exchange rate 6
hydration number 5
properties 4, 33
reduction 118-20
separation 115-8
structure 4, 33
Perchlorate ion electrodes 129-32

185
Perchlorato complexes
lead 43
bismuth 43
transition metals 46-7
Perchloric acid anhydride
preparation 26-7
properties 27
Perchloric acid, anhydrous
preparation 24
properties 24
reactions 26-6
thermal decomposition 25
Perchloric acid, aqueous
acid strength 21
acidity functions 21
activity coefficients 16
analysis 132-3
boiling points 17
chemical properties 21-2
densities 13, 14
electrosynthesis 9-10
high purity 11
manufacture 2
osmotic coefficients 16
physical properties 11-21
preparation 9-11
purification 11
refractive index 20
surface tension 19
titration 124-5
vapor pressure 15
viscosity 18
Perchloric acid, hydrates
composition 22-3
properties 23-4
structure 23
Perchloric esters 53-4, 99
Perchloryl benzene 55
Perchloryl compounds 55
Perchloryl fluoride 52-3, 88
Perhaloalkyl perchlorates 53
Pernet process 10

186
Perylene perchlorate 99
Pharmacology 143-4
Phosphonium perchlorate 52
Photometric methods 125-29
Physiological effects 140-2
Plant studies 147-9
Pollution 160-1
Potassium perchlorate 37-8, 132
Potentiometric titration 124, 130
Precipitation reagents for
alkaloids 83
cesium 83
organic bases 83
potassium 83
rubidium 83
Propellants 101
Pyrotechnics 102
Qualitative analysis 114-5
Quaternary ammonium perchlorate 106, 108, 130, 131
Radiolysis 22
Radiometric methods 132
Raman spectra 5, 21, 25, 34, 40, 42, 43, 51
Reaction rate method 132
Reagents for analysis
acidimetric standard 24, 82
chromatography 88-9
extraction 83-4
precipitation 83-4
steroids 88
sugars 88
titration 81-3
Reduction methods
bacterial 146, 149, 161
chemical 118-9, 122-4
electrolytic 119-20
thermal 118
Refractive indicies 20, 42, 133, 168, 172
Resistivities 12
Rhodium perchlorates 50
Safety
equipment 156-7, 160
precautions 155-8

187
Seed germination 148-9
Selenius acid perchlorate 52
Sensitizing agents 108
Separation of perchlorate
ion exchange chromatography 117
partition chromatography 116
precipitation 115, 128
solvent extraction 117-8
Silica dehydration 87
Silver perchlorate 47-8, 89, 98, 99
Sodium perchlorate 36-7, 132
Solubilities 32, 33, 34, 35, 36, 38, 39, 41, 44, 45, 47, 49, 50,
52, 167-74
Solubility products 32
Solvent extraction 117-8, 126-8
Solvents for
acrylonitrile polymers 100
cellulose 100
methacrylonitrile polymers 100
lead borate glass solder 108
Specific gravities 14
Spectra
EPR 50
infrared 5, 23, 25, 34, 38, 40, 41, 42, 114, 125-6
NMR 5, 6, 21, 34, 36, 37
Raman 5, 21, 25, 34, 40, 42, 43, 51
X-ray 4, 5, 40
Stannous perchlorate 99
Steroids, detection, determination 88
Strontium perchlorate 40
Surface tensions 12, 19, 53
Tetrabutylammonium perchlorate 106
Thallium perchlorate 42, 83, 99
Thermal analysis 34, 35, 44, 46
Thermal decomposition 37, 40, 41, 42, 44, 46, 49, 51, 73
Thermodynamic data 25, 32, 33, 34, 37, 39, 40, 47, 48, 51, 52
Thyroid effects 140, 141
Titanium tetraperchlorate 49
Titrimetric methods 122-5
Toxicities 140, 144-6
Toxicology 144-6
Transition metal perchlorates

188
 crystal structures 43-4
preparation 44-50
properties 44-51
pyridine complexes 44
solubilities 44-5
thermal decomposition 44-6
Transport numbers 12, 48
Trichloromethyl perchlorate 55
Vandayl perchlorate 49
Vapor pressures 12, 15, 40, 48, 53, 167, 169
Vibrio dechloraticans 160
Viscosities 12, 18, 41, 42 48, 53, 166-73
Waste disposal 160-1
Waste water treatment 161
Wet ashing procedures 74-6
Wet oxidations 69-81
Willard synthesis 10
X-ray
diffraction 23, 32, 36, 43, 46
spectra 4, 5, 40
Zinc perchlorate 48, 105, 106

189