• TABLE OF CONTENTS
HIDE
 Title Page
 Dedication
 Acknowledgement
 Table of Contents
 List of Tables
 List of Figures
 Introduction
 Experimental part
 Presentation of data
 Discussion of results
 Bibliography
 Biographical sketch
 Copyright














Title: Solvates of light metal perchlorates with hydrogen peroxide.
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Title: Solvates of light metal perchlorates with hydrogen peroxide.
Series Title: Solvates of light metal perchlorates with hydrogen peroxide.
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Table of Contents
    Title Page
        Page i
    Dedication
        Page ii
    Acknowledgement
        Page iii
    Table of Contents
        Page iv
        Page v
    List of Tables
        Page vi
        Page vii
    List of Figures
        Page viii
        Page ix
    Introduction
        Page 1
        Page 2
        Page 3
        Page 4
        Page 5
        Page 6
        Page 7
    Experimental part
        Page 8
        Page 9
        Page 10
        Page 11
        Page 12
        Page 13
        Page 14
        Page 15
        Page 16
    Presentation of data
        Page 17
        Page 18
        Page 19
        Page 20
        Page 21
        Page 22
        Page 23
        Page 24
        Page 25
        Page 26
        Page 27
        Page 28
        Page 29
        Page 30
        Page 31
        Page 32
        Page 33
        Page 34
        Page 35
        Page 36
        Page 37
        Page 38
        Page 39
        Page 40
        Page 41
        Page 42
        Page 43
        Page 44
        Page 45
        Page 46
        Page 47
        Page 48
        Page 49
        Page 50
        Page 51
        Page 52
        Page 53
        Page 54
        Page 55
        Page 56
        Page 57
        Page 58
        Page 59
        Page 60
        Page 61
        Page 62
        Page 63
        Page 64
    Discussion of results
        Page 65
        Page 66
        Page 67
        Page 68
        Page 69
        Page 70
        Page 71
        Page 72
        Page 73
        Page 74
        Page 75
        Page 76
        Page 77
        Page 78
        Page 79
        Page 80
        Page 81
        Page 82
        Page 83
        Page 84
        Page 85
        Page 86
        Page 87
        Page 88
        Page 89
        Page 90
        Page 91
        Page 92
        Page 93
    Bibliography
        Page 94
        Page 95
        Page 96
    Biographical sketch
        Page 97
        Page 98
    Copyright
        Copyright
Full Text











SOLVATES OF LIGHT METAL PERCHLORATES

WITH HYDROGEN PEROXIDE












By
WALTER DAVID FOUCHER, JR.


A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY









UNIVERSITY OF FLORIDA


August, 1962
















DEDICATION



To the three people whose contributions to this dissertation

and my life are too great for a mere acknowledgement: To my wife,

Linda, whose understanding and patience have known no bounds; to

my parents whose encouragement concerning education has caused me

to strive toward this Ph.D. degree. To these three people I

dedicate this work.















ACKHOWLEDGEIMETS


An acknowledgement does not seem quite enough for the man who

has contributed as much to this work as the author himself.

Dr. George E. Ryschkewitsch has not limited himself to the role

of advisor or chairman of committee, rather he has been more a

partner in this endeavour. His combination of patience, availability

and a sincere desire to teach have made an unmeasurable contribution

to this dissertation and to the author's life.

The author would also like to express his thanks to the other

individuals who have aided him in the accomplishment of his research

and the preparation of this dissertation: His committee as a whole;

Dr. Wallace Brey for aid in interpreting nuclear magnetic resonance

spectra; 1r. Ken Lawson for taking valuable time away from his own

work to run nuclear magnetic resonance spectra; Mr. Newton Levy, Jr.,

the author's laboratory partner, for encouragement, criticism and

advice; Mr. Robert Gould of the Metallurgical Laboratory for x-ray

diffraction work; and finally, a special thanks to Mr. Richard

Logsdon for emergency preparation and repair of glass apparatus,

especially in the final months of this work.

The author also expresses his gratitude to the Office of Naval

Research for the funds necessary to carry out this research.
















TABLE OF CONTENTS


Page

ACKIO-LEDGMIENTS iii

LIST OF TABLES vi

LIST OF FIGURES viii

INTRODUCTION 1

I. Justification of Study 1

II. Historical Background 3

III. Various Physical Chemical IMeasurements Which
May Indicate the Existence of Solvates 5

IV, Scope of the Present Study 7

EXPERIMENTAL PART 8

I. Materials 8

II. Treatment of Glassware 11

III. Apparatus 12

IV. Procedure 13

PRESENTATION OF DATA 17

I. Solubility and Density Data 17

II. X-Ray Data 19

III. Nuclear Magnetic Resonance Data 20







TABLE OF CONTENTS (Continued)


DISCUSSION

I.

II.

III.

IV.

V.


Page

65

65

77


OF RESULTS

Solubility Data

Density Data

Isolation of Solvates With Hydrogen Peroxide

Nuclear Magnetic Resonance Data

Reaction of IMG(CO04)2 With N2H4


BIBLIOGRAPHY

BIOGRAPHICAL SKETCH















LIST OF TABLES


1. Solubility of LCIOCl in 1202-H20 Mixtures at 300C.

2. Solubility of NaC104 in H1202-H20 Mixtures at 30 C.

3. Solubility of KC104 in 1202-H20 Mixtures at 300C.

4. Solubility of RbC104 in 11202-H20 Mixtures at 300C.


5. Solubility of CsC104 in

6. Densities of the Liquid
12024-20 at 300C.

7. Densities of the Liquid
H202-1120 at 300C.

8. Densities of the Liquid
H202-1120 at 300C.

9. Densities of the Liquid
H202-H20 at 300C.

10. Densities of the Liquid
H202-120 at 300C.


1202-H20

Phase in


Mixtures at 300C.

the System LiClO4-


Phase in the System NaC1O4-


Phase in the System KC104-


Phase in the System RbC104-


Phase in the System CsClO4-


11. Solubility of 1~(C104)2 in H202-H20 Mixtures at 300C.

12. Solubility of Ca(C104)2in H202-H20 Mixtures at 300C.

13. Solubility of iH4C104, Sr(C104)2, and Ba(C104)2 in
H202-1120 Mixtures at 309C.

14. X-ray Patterns for Two Wet Residues from 1g(C104)2-
H202-H20 System

15. X-Ray Patterns for 3(C104)2 and its Hydrates

16. Chemical Shift in the Proton Resonance Spectrum of
H202-1120 Mixtures


Page

22

23

24







LIST (O TABLES (Continued)


Page

17. Width of Proton Resonance Line In H202H20 Mixtures 38

18. Proton Magnetic Resonance Data for Dilute Solutions
of Fe(C104)2 H202-H20 Mixtures 39

19, Molality in H202 for the Group aI Nitrates 65

20. Morality in %02 for the Group Ia Perchlorates 66

21. Holality Differences in H202 and 40 of Group
Za Perchlorates 67

22. tj Ft for the Group IS Perchlorates 72

23. Molality of Group ZIt Perchlorates in HO2 and H20U2-20
Mixtures 76

24. Apparent Molar Volume of Group Ia Perchlorates in
Saturated Solutions 78

25. Z*Ray Patterns of Wet Residues and 1f(Cl04)2.2H202 85


vii













LIST OF FIGURES


Page
I, Diagram of Apparatus Used for Dehydration of Rydrazine 10
2. LiC104-H202-B20 System at 30OC 40

3. NaC104-BH02-H20 System at 300. 41
4. 1C104-202-H20 System at 300C. 42
5, abclo4-H2o2-~20 system at 300c, 43

6. CeCL04-20r2-o20 System at 30oC. 44
7, W(0104)2-H202-H20 System at 300C. 45

8. Ca(C104)2-H202-H20 System at 30C. 46
9, Sr(CiO4)2-H202-H20; Ba(C0104)ZM2202-H20; S4C104-Ha02*eH20
System at 300C. 47
10. Plot of Molality of Salt Versus Mfle fraction H202 In
Solvent for the System LIC1O4-H202H20 at 300C. 48
11 Plot of Iblality of Salt Versus Ibles Fraction 4202 In
Solvent for the System NaC104-H202"-20 at 3000. 49

12, Plot of Holality of Salt Versus Mole Fraction 202 In
Solvent for the System KC104-H202-H20 at 3000. 50

13, Plot of Mblality of Salt Versus tole Fraction g202 In
Solvent for the System RbC104-H202-H20 at 300o. 51

14. Plot of Molality of Salt Versus Mote Fraction 8202 In
Solvent for the System CsCC04:0gO%02 0 at 300c 52

15, Plot of Density Versan !'le Fraction U202 In Solvent
for the System LiCla0z4-H202'20 at 300o. 53

16. Plot of Density Versus tblality of Salt for the System
LiC104*H202-H20 at 300C. 54


viitt








LIST OF FIGURES (Continued)


Page

17, Plot of Density Versus Mole Fraction H202 In
Solvent for the System KC104-H202-120 at 300C. 55

18, Plot of Density Versus lolality of Salt for the System
KC104-H202-1120 at 300C, 56

19. Plot of Density Versus -Mole Fraction 1202 In
Solvent for the System PbC104-H202-120 at 300C. 57

20. Plot of Density Versus Molality of the Salt for the System
PlbCl04-H202-1120 at 300C. ,S5

21. Plot of Density Versus ole Fraction H202.In
Solvent for the System CsCO14-H202-H120 at 300C. 59

22. Plot of Density Versus oblality of Salt for the System
CsC104-H202-IO2 at 300C. 60

23. Plot of Density Versus Mole Fraction H202 In
Solvent for the System NaC104-H202-1120 at 300C. 61

24. A Representation of the Diffraction Patterns for
Salt A, Salt B, and If,(C004)2.21120 62

25. Chemical Shift in the Proton Resonance Spectrum of
1202-120 Mixtures 63

26. Width of Proton Resonance Line Versus Mole Fraction H1202 64

27. Plot of AFt Versus the Reciprocal of the Cationic Radius 73

28. Plot of &Ft Versus Apparent ioblar Volume Difference
(vH20- VH20) 79















INTRODUCTION

I. Justification of Study


There are two main questions which it was expected could be

partially answered by this study. One concerns the relationship

between ions in solution and salts which can be crystallized out

of the solution; the other concerns elucidation of the differences

and similarities between hydrogen peroxide and water as solvent

systems.

It has long been assumed that, in aqueous solutions of

ionized substances, ions may be closely linked with one or more

molecules of water (1). This has given rise to the concept of

primary hydration number which is interpreted by some as referring

to the existence of a definite complex ion resulting from hydration,

(e.g. Werner's emphasis upon analogies between hydrates and ammines (2).

In the light of more modern theory, however, there is no need to

suppose that ions dissolved in a solvent combine with the solvent to

form solvated complex ions of definite composition. Random movement

can account for an everchanging atmosphere of solvent molecules (3).

Nevertheless, T. G. Owe Berg (4) has suggested that there exist

certain liquid hydrates in addition to those crystalline hydrates

which have been isolated and studied in the solid state. Although the

water molecules are not always considered bound to the same atom and

may not represent the same kind of structure usually referred to as a









hydrate in the cold state, still some of these substances are

highly analogous, or perhaps it should be said, similar to those
crystalline hydrates,

Other investigations have been conducted to provide a clearer

understanding of the nature of liquid water, especially the manner

in which ions may modify this structure by their action on molecules

in their immediate neighborhood. In addition to this study on

solutions the study of crystalline hydrates hds led to recognition

of several modes of hydration, e.g. water of coordination, zeolyte

water, etc. Hydrogen peroxide systems, on the other hand, have not

been studied sufficiently to establish such categories.

It has been stated that hydrogen peroxide and its water solutions

possess, in general the solvent or solute properties that water alone

possesses (5). This observation is supported by the fact that such

physical properties as dipole moment, dielectric constant, magnetic
susceptibility, and others, including amount of hydrogen bonding, are
very similar. However, there are notable differences between the two

substances, and even between mixtures of the two solvents. For
*6 *1 *1
example, the specific conductivities of water (0.5 X 10 ohm cm. )
*6 *1 -1
and 98 per cent hydrogen peroxide (1.72 X 10 ohm cm. ) are very
similar, but the curve of specific conductivity versus weight per cent

hydrogen peroxide shows a gentle maximum at approximately 50 per cent

hydrogen peroxide, The shape of this curve can be explained by the

varying concentrations of the dissocation products of water and

hydrogen peroxide and the states in which they exist.









II. Historical Background


In an attempt to elucidate and systematize the differences

between hydrogen peroxide and water as solvents and to point up

the similarities, Floyd and Gross (6) studied the solubilities

of sodium chloride, potassium chloride, lithium nitrate, sodium

nitrate, potassium nitrate, and sodium fluoride over the full range

of concentrations of hydrogen peroxide water solutions at 25,

15 and 0 degrees centigrade. The results:of these studies indicated

that ion size considerations should be of first importance in

determining the amount and type of solvation and the solubility

as the composition of the solvent changes from pure water to pure

hydrogen peroxide. For example, they observed that the solubility

of potassium nitrate increases with increasing hydrogen peroxide

concentration. Further, Turner (7) has presented evidence for the

existence, at 25 degrees centigrade, of a salt having the composition

=no3 1)02.
Everhard and Gross (8) extended these studies to rubidium

nitrate and cesium nitrate and reported three solid phases in the

RbN O3 1 H202 120 system: RbNO3*3/7H202, RbNO3.'I202 and anhydrous

PRb 03. Their interpretation of these date was in terms of a

preferential solvation of the larger ions (K+ and larger) by hydrogen

peroxide rather than water. They also noted that there seemed to be

no simple relation between dielectric constant and solubility

differences. This is interesting since Akerlof and Truck (9)

interpreted their solubility studies of NaCl, KCI, HaN 3, KBr, KI,

I:204, NII4C and PbO(03)2 in CH30H-H,20 and H202-H20 ..Ytures as due










to the changes of the dielectric constant of the medium.

A general trend of increasing solubility in the hydrogen

peroxide water system with increasing hydrogen peroxide con-

centration in Group la was found by Everhard and Gross (8). They

expressed this trend by a linear relationship between difference

in solubility in pure water and in pure hydrogen peroxide and ionic

radius. The only ion which did not follow this relationship was

the cesium ion. This was explained as due to the lower charge

density of the cesium ion,

Further evidence of complex formation in hydrogen peroxide

solutions was obtained as early as 1902 by Jones et. al. (10).

Their results on the study of the freezing point lowering of water

and aqueous hydrogen peroxide by KC1, Na0IO3, and KN03 showed that

the salts produced a greater depression in water than in aqueous

hydrogen peroxide. In the case of KIO3 this could only be

interpreted as due to the combination between the salt and hydrogen

peroxide. The fact that NaNO3 shows different behavior in hydrogen

peroxide was noted by Naass and Hatcher in 1922 (11) when they

found that the amount of ionization of NaNO3, NaCi, and Na2SO4

wras approximately the same in hydrogen peroxide as in water, but

there was no evidence of complex formation except in the case of

Na2S04.2H202 (presumably an anion solvate).









III. Various Physical Chemical
Measurements Which May Indicate the Existence of Solvates


In general the method used to determine the existence of

hydrates or solvates involves examination of the variation of

some property of the solution as a function of the solution compo-

sition. Variations in the curves of vapor pressures, salt

solubility, and electrical conductivity, such as maxima, minima,

inflection points, etc.indicate the existence of some interaction

between the solute and solvent. If it is possible to crystallize

solid solvates from the solutions, further characterization can be

carried out by conventional analysis or by x-ray diffraction.

Solubility studies have the double value of indicating the existence

of solvated species by discrepancies in the curve of solubility

versus composition and also of immediately giving data to compare

one solvent to another and mixtures of the two.

A relatively new method used to detect solvation and/or complex

formation in solutions and solid phases is nuclear magnetic resonance

spectrometry. Jardetzky and Wertz (12) studied the absorption of the
23
Ha ion in aqueous solution as a function of concentration, viscosity,

and nature of the anionic species. The fact that the effect of a

charge decreases with the square of the distance (Coulomb's law),

leads one to believe that only molecules in immediate contact with

the ion can appreciably distort the symmetry of the electric field

at the nucleus. The result of this distortion is that one can

distinguish three types of interactions in solutions of ions, i.e.

interactions of the completing agent with the cation being studied

(which in Jardetzky's case was Na ) which are weaker than those between










solvent and cation, interactions which are of the same order of

magnitude, and interactions which are stronger than the interactions

with the solvent. The stronger interactions give rise to a broadening

of the cationic line due to the distortion of the electric field.

The interpretation, then, of broadening in dilute solutions is an

actual preference of the cation for the given complexing agent over

the solvent molecule. Broadening in highly concentrated solutions may

be interpreted as due to ion pair formation enhanced by relative

unavailability of solvent molecules. This latter type of broadening

may find some application in studies of activities.

In addition to investigating cations which have nuclear moments

one can also study 017 spectra as Jackson et. al. (13) have done.

Their study of the 017 spectra of aqueous solutions has shown that

it is possible to distinguish solvent water from water in the hydration

sphere of certain cations. They suggested a method for determining

the hydration number of paramagnetic cations by measuring the change

in the water available to interact with the ions due to water retained

in the hydration sphere of the cations.

The NMR absorption of protons in hydro-yl and hydronium ions is

displaced about 10 ppm. to lower field from the proton resonance

absorption in water (14). If the hydroniun ion is viewed as a

hydrated proton it is analagous to a hydrated cation and the above

observation leads to the conclusion that a study of the proton

magnetic resonance spectra of solutions of salts in protonated solvents

will give some information concerning the type and amount of inter-

action between ions and solvent molecules.









IV. Scope of the Present Study


The fact that the perchlorate ion does not interfere with

complex formation results in a simplification of processes taking

place in solutions of perchlorate salts; in general, complex or

solvate formation involves only the cation. With this in mind and

with hope of extending the relationships to other cations in addition

to those of Group la, the present study of the solubilities of the

perchlorates of Group la and IIa was conducted. The construction

of ternary phase diagrams for the various salts in mixed i202 120

solutions gives a good insight into interactions occurring in the

systems and by the use of the Schreinmacher method of wet residues

the composition of the salt in equilibrium with the solutions can

be determined (15). The basis of this method is that any wet residue

composed of a given solid phase and a saturated solution must lie on

a straight line joining the composition of the solid phase and that

of the saturated solution. Consequently extension of a tie line

constructed between any corresponding pair of residue (R) and solution

(S) points must pass through the composition of the solid phase (onm

extension past R).

Using this method plus other information derived from these

solutions, such as density measurements, it should be possible to

draw some conclusions concerning the relationships between water and

hydrogen peroxide and mixtures of the two as modified by salt solutes.















EXPERIMIhTAL PART


I. Materials



Hydrogen peroxide was:obtained from the Becco Chemical

Division of the Food Machinery and Chemical Corporation, Buffalo,

New York, as Bocco Hydrogen Peroxide SP "93" of 98 per cent

minimum concentration and used directly without further purification.

Dilution was carried out with conductivity grade water obtained by

passing distilled water through an ion exchanger resulting in

water with a conductivity equivalent to less than one part per

million NaC1.

All perchlorates e::cept the rubidium and cesium salts were

purchased from the G. Frederick Smith Company, Columbus, Ohio, and

used without further purification for most of the solubility

measurements. Any magnesium perchlorate which was usrd in the

more concentrated solutions of hydrogen paro2-ide was recrystallized

from water as the hex:ahydrate and then dehydrated at 250 degrees

centigrade, under vacuum for several hours. Rubidium and cesium

perchlorates were prepared by treating the respective carbonates

with an excess of 72 per cent perchloric acid. The salts precipitated

from the resulting solutions and were filtered, dried and then

washed with deionized water until the pHi of the washings was constant.

They were then redried.










Dowex 50W-xl2 cation exchange resin was purchased from the

J. T. Baker Chemical Company, Phillipsburg, l7ew Jersey, and was

sized as 200 to 400 mesh. The resin was prepared and regenerated

according to the directions of Day (16). The ion exchange column

was made from a pyrex glass tube of one centimeter inner diameter

and twenty centimeters in length. Near the bottom there was a

B porosity Corning glass frit; below the glass frit the tube was

joined to a teflon barrel stopcock used to control flow. This

column was mounted on a small bell jar (of a size suitable for

containing a one hundred millimeter beaker) which was connected to

a water aspirator. The vacuum created by the aspirator was sufficient

to pass one hundred milliliters through in about twenty minutes.

The hydrazine used was Hathieson technical grade, minimum

97 per cent 12124, and was dehydrated according to the directions of

Bock and Rudolph (17). Figure one represents the apparatus used

for this dehydration.










To Vacuum Pump

JL____T


BaO


Figure 1. Diagram of Apparatus Used For Dehydration of Hydrazine










II. Treatment of Glassware



In all cases any glassware which was to come in contact

with concentrated solutions of hydrogen peroxide was scrupulously

cloned in the following manner: each piece was allowed to stand

in a concentrated solution of Le:ocCel coacercial laboratory

glass cleaner for a number of days and then rinsed thoroughly

with tap water. Ne:xt the glassware received a thorough treatment

with concentrated nitric acid and another rinse with tap water.

The final rinse was with deionized water of the type used to make

up all solutions.









III. Apparatus


All samples used in the solubility determinations were Kept

immersed in a thermostatted bath held at 30.00-:0.01 degrees

centigrade by the use of a Sargent "Thermonitor" in conjunction

with a Sargent heating and circulating tower. The initial

temperature setting was obtained by the use of a Bureau of Standards

certified thermometer and thereafter was observed on a Beclkann

type thermometer. Fluctuations in temperature were controlled to

jO.01 degrees centigrade according to literature accompanying the

equipment. A check of this fluctuation indicated that this was the

case as long as the temperature of the room was held at approximately

25 degrees ccntigradc.

X-ray diffraction patterns were run on a Norelco x-ray

diffractometer with a Cu-K alpha target and a nickel filter.

Data were recorded on a continuous strip recorder by the use of

a counter rather than obtaining powder patterns.

Nuclear magnetic resonance data were obtained by using a

Varian high resolution spectrometer operating at 60 megacycles per

second and 14,000 gauss. Benzene, external to the sample, was

used as a reference material, and sidebands were applied in the

conventional manner, monitored continuously with a Hewlett-Packard

frequency counter,









IV. Procedure


A. Solubility and Density


The solubilities of LiCI04, NaC104, KC104, RbC104, and CsC104

at 30.00 degrees centigrade were determined over the full range

of 11202-1120 compositions and ternary phase diagrams were constructed

for these systems. Saturated solutions of the salts were prepared

by equilibrating an excess of solid anhydrous salt with the

H202-I120 mixtures in a long necked 50 milliliter flask, resulting

in a saturated solution in equilibrium with a wetted solid. These

samples were kept immersed in a thermostatted water bath held at

30.00+0.01 degrees centigrade and agitated continuously for several

days. Duplicate samples were then pipetted for 1202 analysis,

density determinations and perchlorate determinations. All samples

were taken through an Ace Glass fritted filter tube of D porosity.

The analyses were then repeated on new samples from the same flask

a few hours later as a test for equilibrium.

Hydrogen peroxide was determined by titration with approximately

0.1 Normal potassium permanganate. Perchlorate concentration was

determined by passing the sample through an ion exchange column

containing Dowex 50U-x12 cation exchange resin. The effluent

(BC104) was then titrated with approximately 0.1 ITormal sodium

hydroxide solution by the use of a Sargent-1almstadt automatic

titrator, Model SE, E . Sargent Company, Chicago, Illinois.

Day (16) reports that it is possible to determine perchlorate ion

to a deviation of 0p.025 per cent by this method. Analysis of

standard solutions of KC104 indicated that this was true. Caution










should be taken to insure that any salts analyzed by this method

are pure since the method is not specific for perchlorate but

rather measures total anion concentration. The density was found

by weighing a sample delivered from a volumetric pipette. All

sampling apparatus which came in contact with the solutions was

first preheated to the temperature of the sample before being used.

Wet residue scrpling was accomplished in one of two ways,

either by rapid filtration of the equilibrium mixture through a

fritted glass funnel, the filtration apparatus being immersed in

the constant temperature bath, or by merely removing a portion

of the wet solid from the equilibrium mixture with a long handled

nlass sampling spoon. Both methods gave comparable results, the

former resulting in a drier wet residue and hence a shorter

extrapolation to the composition of the solid phase on the ternary

phase diagram.

Analysis of cesium perchlorate on the ion exchange column

did not appear to be as facile or as accurate as merely drying

the samples at 110 degrees centigrade in a small oven to determine

total salt concentration. The limited solubility of cesium

perchlorate in both solvents was one of the determiining factors

in this choice.


B. X-ray Diffraction Data Studies


Wet residues for x-ray diffraction studies were filtered as

mentioned above, but also with dry nitrogen flushing the filter,

and then immediately placed in the sample holder and covered with

a plastic film (commercially called "Parafilm," and purchased from









Fisher Scientific Company, Silver Cpring, Maryland). These

sample holders were then placed in a vacuum dcssicator and analyzed

as soon as possible. "Parafilm" by itself gave a broad peak

between scattering angles of 5 and 10 degrees and two sharp peaks

at d-spacings of 4.14 and 3.73 but did not interfere with the

determination of the positions and intensities of the sample peaks.


C. Iluclenr I-irnetic Reconrnce Drt~


Solutions for INtR investigation were prepared specifically

for these determinations and not taken from samples used in

solubility determinations. Five peo!:s were recorded for each

sample run and an average taken of the five to obtain the tabulated

value. In order to minimize hydrogen peroxide decomposition all

sample tubes were treated by first allowing them to stand with the

respective solutions for a short while; after discarding this

solution a new sample was taken and run immediately.


D. Reaction of 112114 with Iagncsium Perchlorate


The reaction of hydrazine with magnesium perchlorate was

carried out in the same apparatus used to dehydrate the N2H4 and

described in Figure one. flgnesium perchlorate was substituted

for the BaO and the hydrazine was condensed cnto it. The quantity

of (C0104)2 used was sufficiently small that a solution in the

hydrazine resulted. The hydrazine was then distilled back into

its own trap (pr;calcghcd) until ((C104)2 coiaenced to precipitate

from the solution. At this point a pressure reading was made,

corresponding to the vapor pressure of the saturated solution.









Frequent checks of the vapor pressure were made during the

evaporation. Whcn it was observed that the pressure dropped, the

distillation was discontinued and some of the hydrazine was

recondensed onto the ~g(C104)2. This procedure allowed selection

of a pressure to complete the distillation which was somewhere

between the vapor pressure of the saturated solution and of any

solvate which may have formed. Distillation was finished at this

selected pressure. Both the weight of the It~(C104)2 flask and the

N2114 flask were checked occasionally until their weights were

constant, which signified the completion of the distillation.















PRESENTATION OF DATA


I. Solubility and Density Data



The experimental data for the solubilities of the Group la

perchlorate, at 30 degrees centigrade, over the complete range of

H202-H20 concentrations is presented in Tables one through five.

These tables show the weight per cent of each of the three components

in solution as well as the data for the hydrogen peroxide and

perchlorate concentrations of wet residues. Tables six through ten

show the data for the mole fraction of hydrogen peroxide in the

solvent, the morality of the salt, and the density of the solutions,

Table eleven shows the data for the 13(0104)2*-H202- 20 system over

the range of hydrogen peroxide concentrations studied. Table twelve

contains the data for the calcium perchlorate system. Table thirteen

contains the data for ammonium perchlorate, strontium perchlorate,

and barium perchlorate at 30 degrees centigrade. In the case of the

latter Group IIa perchlorates it was found that their saturated

solutions in hydrogen peroxide were very unstable with respect to

hydrogen peroxide decomposition. For this reason less data were

collected for these compounds than for other perchlorates. One

point was determined for ammonium perchlorate in concentrated hydrogen

peroxide solution to see if its solubility was increased in this

solvent analogously to that of potassium perchlorate. Since such was

not the case nothing further was done with this system.










Figures two through nine are ternary phase diagrams representing

the data tabulated in Tables one through thirteen. Figure nine

represents the data for Sr(C104)2, Ba(C104)2, and H40C104. Note

that no solubility curves are drawn since the data were not

sufficiently detailed; however, lines connecting wet residue points

and solubility points are drawn to indicate the composition of the

solid in equilibrium with the solution.

Figures ten through fourteen are plots of the molalities of

the Group la perchlorates as functions of the mole fraction of

11202 in the solvent. Figures fifteen through twenty-two show the

data for the densities of the solutions of the Group la perchlorates

(except flaC104) as functions of the mole fraction H202in the solvent

and molality of the salts. Since the densities of the solutions

of NaC104 varied little with concentration the plot of density

versus mole fraction H202 in the solvent (Figure twenty-three) is

made on a much expanded scale, which gives the apparent scattered

appearance to the data.









II. X-ray Data


Table fourteen shows the d-spacings, relative intensities,

and angle of reflection (29) for the two wet residues taken from

the Mg(ClO4)-H202-H20 systems, Table fifteen shows the d-spacingc

for salts prepared in this laboratory corresponding to g(C104)2

and its three hydrates, in order of decreasing intensity (see

discussion of results for comments concerning reliability of

these data). Figure twenty-four is a representation of the

diffraction lines in the patterns for the two wet residues and

the salt having a composition corresponding to gs(C104)2.2H20.









III. Nuclear Magnetic Resonance Data


Table sixteen represents data for the chemical shift, relative

to benzene, of the proton resonance line for mixtures of 11202 and

H20. It would normally be expected that, with sufficiently slow

exchange rates, the spectrum of solutions of hydrogen peroxide and

water would consist of two separate lines (18), however, only one

line is observed over the complete range of concentrations of H202

and 1120 at room temperature. In this table the parameter l

represents the distance, in cycles per second, which this single

proton resonance line is removed from the proton resonance line of

benzene, the external standard. In the cases where the line is at

a higher field than the benzene line & is positive, When the

line is shifted to lower fields, by the addition of H202, and falls

below the value for benzene, the sign of 4 is negative. Figure

twenty-five graphically presents this data as a linear relationship

between & and mole fraction H202. The line drawn through the

points is that calculated by the method of least squares (19).

Table seventeen presents data on line width at half height, in

cycles per second, as a function of mole fraction hydrogen peroxide.

These data are also presented in Figure twenty-six where it can be

seen that the width attains a maximum value at approximately 0.73

mole fraction H2202.

The data in Table eighteen are the result of proton resonance

measurements on solutions of water and hydrogen peroxide to which

various quantities of 6(C104)2 had been added. The results are







21


reported as normality of 13g(C104)2, mole fraction 11202 in the

solvent, distance from the benzene proton resonance line, and

width of the line at half height.










TABLE I:

SOLUBILITY Or, LiClO4 ii1; 11202-1120 M4XTURES AT 300C


Wr. % LiC1Q4 Wt. % H202 Wt. % 120 Wet Lrcsidue



38.67 0.00 61.33 -
38.47 0.34 61.19 58.2-0.1
38.35 1.49 60.16 56.7-0.4
38.05 1.54 60.41 53.4-0.7
34.05 11.79 54.16 63.8-0.9
31,02 13..0 50.18 53.2-6.2
31.05 19.21 49.74 55.2-6.3
30.15 22.34 47.51
30.91 23.10 45.99 51.9-9.2
30.37 23.87 45.76 57.3-6.3
26.46 39.26 34.28
25.83 43.97 30.20 46.7-20.9
25.32 45.41 29.27
47.15 45.80 7.05 71.9-23.2
41.40 46.32 12.28 65.9-:2,9
39.20 46.55 14.25 57.4-14.4
40.15 46.57 13.28 49.3-28.7
26.90 47.50 25.60 51.3-17.3
37.86 43.67 13.47
36.62 49.19 14.19
26.07 49.67 24.26
26.45 49.75 23.80
46.78 50.05 3.17
25.94 50.90 23.16
46.32 51.17 2.01 71.G-26.5
33.21 51.78 15.01 49.6-24.8
26.80 51.99 21.21 44.3-27.1
28.04 52.51 19.45 46.5-26.7









TABLE 2

SOLUBILITY OF NaC104 IN 11202-H20 IXTURES AT 30C


Wt. % NaC104 Wt. %11202 Wt. % 20 Wet Residue
NaC104-H202



68.92 0.00 31.02
68.16 0.91 30.9M3
68.62 1.21 30.17 83.3- 0.3
66.55 3.42 30.0:3
64.97 8.20 26.83
64.42 9.75 25.,3 83.0- 2.0
63.84 10.45 25.71 86.1- 0.8
63.73 11.06 25.21
62.91 11.75 25.34
62.33 13.57 24.10
62.89 14.35 22.76
62.17 14.59 23.24
60.93 16.77 22.30
59.97 18.22 21.81
58.26 20.45 21.29
57.02 22.72 20.26
55.27 25.09 19.64 74.1-14.3
54.40 26.52 19.08 85.2- 9.1
53.18 28.44 18.33 -
52.47 29.19 18.34
51.69 30.84 17.47
49.18 35.00 15.82 -
48.30 35.86 15.84
47.59 36.84 15.57
45.96 40.51 13.53
43.30 44.71 11.99
42.79 45.38 11.33 90.5- 7.3
39.29 51.47 9.24 93.1- 5.0
33.97 52.16 8.87
35.31 61.93 2.76 89.6- 8.5
35.16 61.97 2.87 89.0- 9.7

















TABLE 3
o
SOLUBILITY OF KC1O4 IN H202-H20 4'IXTURES AT 30 C


Wt. % KC104 Wt. % I202 Wt. % H20 Wet Residue
KC104-.202


2.42 0.00 97.5 -
3.11 36.66 60.23
3.29 46.37 50.34
3.42 50.31 46.27
3.75 60.80 35.45
3.74 61.95 34.31
3.9.2 67.20 2.88 -
3.97 67.96 28.07
4.20 72.16 23.64
4.25 73.94 21.81
4.47 78.43 17.10
4.70 83.35 12.45
4.76 84.30 10.94
5.06 88.74 6.20 96.7-3.2
5.38 90.92 3.70
5.48 91.08 3.44
5.32 91.79 2.89 96.7-3.3



















TABLE 4

SOLUBILITY OF RbC104 IN H202-H20 M rTURES AT 300C


Wt. % RbCl04 Wt. % O22 Wt. % -20 Wet Residue
RbC104-11202


1.60 0.00 98.40 -
3,09 29.46 67.45 96.3- 1.1
4.48 68.30 27.21 97.0-2.1:
5.03 74.85 20.12 94.1- 3.2
6.41 78.31 15.28 68.4-26.0
6.84 88.04 5.12 95.5- j.2
7.21 89.61 3.18 94.4- 2.1
6.93 92.19 0.98 96.8- 2.7



















TABLE 5

SOLUBILITY OF CsCl104 IN H202-H20 ~I-IXTIMES AT 300C


wt. % CsCIO4 W t. % 1202 Wt.% II20 Wet Residue
CcC104-1202



2.23 0.00 97.77
2.62 6.51 90.91 62.6-24.1
2.69 8.77 88.54 79.9- 1.8
6.92 50.25 42.83 64,5-19.4
8.22 56.30 35.48 73.9-15.3
10.08 64.91 25.01 52.2-34.5
12.05 69.05 18.90 63.9-27.1
18.00 77.00 5.00 72.6-25.9
17.24 77.13 5.63 65.1-32.4












TABLE 6

DENSITIES OF THE LIQUID PHASE IN THE SYSTEM
LiC104-1202-H20 AT 300C


Sble Fraction
H202 Ia Solvent olnlity LiClO4 Density of Solution
H22In Solvent


0.000
0.003
0.013
0.013
0,103
0.166
0.170
0.199
0.210
0.216
0.377
0.435
0.451
0.496
0.520
0.525
0.538
0.565
0.533
0.634
0.646
0.647
0.650
0.657
0.666
0.775
0.893
0.931


5.929
5.876
5.846
5.773
4.852
4.226
4.232
4.057
4.205
4.099
3.382
3.273
3.186
3.459
3.314
3.380
3.292
3.441
3.662
6.059
6.028
5.430
6.305
5.726
6.651
8.384
8.261
8.274


1.2796
1.2776
1.2833
1.2,5
1.2970
1.3121
1.3113
1.3211
1.3216
1.3198
1.3717
1. 336
1.4003
1.4242
1.4274
1.4428
1.4493
1.4589
1.4776
1.6106
1.5329
145811
1.6059
1.5807

1.7681
1.7644
1.7914


-- -----











TABLE 7


DENSITIES


OF THE LIQUID PHASE II THE SYSTE1
NaC1O4-H202-H20 AT 300C


1oble Fraction
H202 In Solvent oblality NaC104 Density of Solution


0.000
0.015
0.021
0.057
0.139
0.167
0.177
0.188
0.197
0.230
0.249
0.250
0.285
0.307
0.337
0.373
0.403
0.424
0.450
0.457
0.483
0.5:39
0.545
0.556
0.613
0M664
0.670
0.747
0.757
0.920
0.922


18.144
17.482
17.858
16.248
15.146
14.786
14.418
14.349
13.852
13.513
13.421
13.840
12.736
12.234
11.399
10.834
10.091
9.742
9.276
9.015
8.738
7.903
7.629
7.415
6.945
6.236
6.108
5.285
5.215
4.423
4.458


1.7003
1.7030
1.6900
1.6925
1.6932
1.6993
1.7077
1.6582
1.7031
1.7017
1.6882
1.7104
1.6934
1.6869
1.6844
1.6877
1.6876
1.6732
1.6799
1L6788
1.6571
1.6692
1.6685
1.6584
1.6631
1.6437
1.6669
1.6564
166542
1.6627
1.6556















TABLE 8

DENSITIES OF THE LIQUID PHASE IN THE SYSTEM
KC104-H202-H20 AT 300C


Hlole Fraction
1202 In Solvent 1alality KC104 Density of Solution



0.00 0.179 1.0105
0.24 0.232 1.1521
0.33 0.246 1.1970
0.37 0.256 1.2108
0.49 0.280 1.2770
0.48 0.281 1.2736
0.55 0.294 1.3114
0.56 0.298 1.3348
0.64 0.316 1.3212
0.64 0.320 1.3665
0.71 0.338 1.3767
0.78 0.354 1.4109
0.80 0.361 1.4013
0.88 0.385 1.4333
0.93 0.410 1,4554
0.93 0.419 1,4645
0.94 0.406 1,4547



















TABLE 9


DENSITIES


OF THE LIQUID PHASE IN THE SYSTEM
rbC104-Ho2-H12 AT 300C


Mole Fraction
H202 In Solvent iMlality RbC104 Densityof Solution



0.000 0.088 1.0040
0.188 0.172 1.1282
0.571 0.253 1.2956
0.663 0.286 1.3525
0.731 0.371 1.3356
0.901 0.397 1.4665
0.937 0.420 1.4772
0.980 0.402 1.4760

















TABLE 10

DENSITIES OF THE LIQUID PHASE IN THE SYSTEM
CcC104-H202-H20 AT 300C


Mol Fraction
H202 In Solvent Eblality CsC104 Dancity of Solution



0.000 0.098 1.0129
0.037 0.116 1.0352
0.050 0.119 1.0485
0.383 0.320 1.2655
0.457 0.385 1.3042
0.579 0.482 1.3944
0.659 0.590 1.4186
0.879 0.896 1.5637
0.891 0.945 1.5718













TABLE 11

SOLUBILITY OF 1-(C104)2 II I202 -I20 -MIXTURES AT 300C


Wt. % ,(0C104)2 Wt. % Wt. % Wet Residue Density
H202 1h20 1g (C104)2-11202


50.36 0.00 09.63 1.4716
59.39 28.35 12.26 72.2-15.0 1.8008
57.80 32.02 10.18 72.4-13.6 1.7555
58.44 29.02 12.54 74.3-12.9 1.8252
56.35 34.05 9.50 70.6-17.4 1.3797
56.57 34.36 9.07 --
56.78 35.83 7.39 71.2-18.2 1.,543
56.51 36.07 7.42 74.0-13.2 1.8547
56.80 36.41 6.79 73.0-15.4 1.8565
58.66 36.45 4.89 72.6-17.1 1.8951
56.90 36.82 6.28 68.9-21.4 1.8425
56.07 36.34 7.09 70.0-19.0 1.9006
56.53 36.97 6.50 68.8-21.9 1.8759'
56.99 37.14 5.87 73.7-15.5 1.8597
56.73 37.27 6.00 70.4-19.0 1.8572
57.79 37.88 4.33 75.6-17.3 1.8629



















TABLE 12

SOLUBILITY OF Ca(C104)2 IN H202-H20 MIXTURES AT 300C


Wt. % (CaC104)2 Wt. % H202 Wt. % H20 Uet Residue
Ca (C104) 2-H202


66.14 0.00 33.87
75.06 12.96 11.98 90.3- 5.1
73.43 15.29 11.28 90.2- 3.6
70.69 15.59 13.72 74.2- 5.8
61.22 17.71 21.07 72.4- 5.8
66.60 19.14 14.26 70.2-10.3
66.45 29.21 4.34 91.4- 4.0
















TABLE 13

SOLUBILITY OF N81C104, Sr(C104)2, AND Ba(C104)2 IN
H202-H20 FIXTURES AT 30C


Wt. % H14C104 Wt. % 1202 Wt. % 120 Wet Residue
-H14Cd104-H202


11.56 77.88 10.76 94.7- 4.5


Wt. % Sr(C104)2 Wt. 7 H202 Wt. % H120 et Residue
Sr. (C04)2-H202


74.63 0.00 25.37
64.60 25.83 9.96 83.0-11.8



Wt. % Ba(C104)2 Wt. 7% 1202 Wt. % 820 Wet Residue
Ba(C104) 2-1202


67.63 0.00 32.37
53.96 38.82 2.22 87.6- 9.0










TABLE 14

X-RAY PATTERNS FOR TWO WET RESIDUES
FROHM g (C104)2--H202-H20 SYSTEM


SALT A SALT B
28 d-spacing intensity 29 d-spacing intensity


13.1
18.2



26.2
26.6
27.8

28.8
29.5

33.2
34.4
34,7




39.7
41.5

44.3
42.9
47.8
50.6
53.4
56.8


6.75
4.87


3.39
3.35
3.21

3.10
3.03

2.70
2.60
2.58




2.27
2.17

2.04
2.01
1.90
1.80
1.71
1.62


V..S
M.
1H.

V.W.
V.W.

V.W.

V.W.





V.H.

V.1W.
V.W.
Vw.
VW
V.*W,
V.W.
V.W,1


12.8
18.1
18.7
22.3
25.7
26.1
26.6
27.6
28.0
29.0
30.0
32.2
33.1

34.7
36.0
34.6

33.4
37.0

44.2
44,3
40.0
47.6
49.4

56.6


6.9
4.89
4.74
3.98
3.46
3.41
3.35
3.23
3.18
3.08
2.98
2.78
2.70

2.58
2.48
2.43

2.34
2.25

2.05
2,04
2.00
1.91
1,34
1.84
1.63


f.
3.
V.S.
V.S.

M.
S.
M.
N.

W.
VW.
V.11H.
V.l.
V.W.
V.W,
V.W.
W,
H.

W.
W,

W.
VW.
W.


I
















TABLE 15

X-RAY PATTER S FOR M(C104)2 AND ITS HYDRATES


Mg (C104)2 Mg(C104)2.2H202 Mg (C104)2.4H20 Mg (C104)2. -H20
d int. d int. d int. d int.
-


3.41
4.85
3.35
6.9
3.95
3.21
2.59
2.05
2.95
2.83
2.70
2. 13
1.90
1..4
1.80
1.72
2,41
2.32


VS.s

S.
M.
M.
M..

W.
W.
W.

W,
W,
w.o
W.
W.
Vw.
V.W.


4.74
4.90
3.24
3.98
3.42
3.36
3.11
2. 95
2.80
2.78
2.70
2.58
2.43
2.34
1.91
1.84
6.90
5.60
5.30
2.05
2.00
1.63


V.S.
S.
S.
M.
M.
1M.
W.


W.
W,
W.
W.
W.

wN.
V.W.

V.W.
V.W.
V.W.
V.H.


4.68
3.21
3.39
3.32
4.82
4.79
3.95
3.08
2.94
2.35
2.75
2.62
2.50
6.70
5.30
5.24
1.99
1.91
1.84
1.62


V.S.
V.S.
S.

M.
1M.
M.

M.
W.
W.
w.

W.
V.W.
V.W,
V.w.

V.W.
V.14
V.W.
V.W.
VW.

V~lw.
vI'W.


4.18
4.00
2.66
2.88
2.59
2.27
1.96
5.0
2.32
1.97
1.89
1,85
1,70
1.70
1.57


V.S.
V.S.
S.
M.
M.


W.
W.
W.
W.
W.
W.
W.
V.W.
















TABLE 16

CHEMICAL SHIFT I TIH PROTON RESONANCE SPECTRUM
OP 11202- H20 MIXTURES


Mole Fraction 1202 6In C.P.S.*


0.000
0.035
0.074
0.086
0.157
0.197
0.245
0.328
0.419
0.435
0.516
0.665
0.672
0.772
0.910
0.951
0.955
0.960
0.960
0.980


+106.1
+ 96.4
+ 89.8
+- 82.94
+ 63.54
+ 40.60
+ 41.09
+ 1.74
- 28.79
- 36.52
- 60.07
-107.55
-102.82
-132.74
-182.30
-189.80
-195.82
-188.27
-192.30
-196.50


*":Shift relative to benzene, shifts
positive values.


to higher fields having


- "' ~~' ' '
















TABLE 17

WIDTH OF PROTOII r-lSO:l;ICE LITE IN I1202-1120 11ITTURES


Mole Fraction 1202 Line iidth
In C.P.S.*


0.000
0.086
0.157
0.197
0.245
0.328
0.430
0.435
0.516
0.665
0.772
0.910
0.955
0.960


00.86
18.92
43.89
28.03
59.58
50.00
65.33
66.10

74.02
82.22
48.51
10.40
17.01


*At half height


-- ------

















TABLE 18

PROTO MNAGIETIC RESOG.IPI]CE DATA FOR DILUTE SOLUTIONS OF
1,(C10O4)2 IN H202-H120 IXTURES


SMole Normality Line Width
Fraction 11202 Ii (C104)2 6 In C.P.S.* In C.P.S.*
In Solvent X 10


0.00 3.77 +105.41
0.584 3.77 78.48 73.54
0.715 0.17 -136.74 80.73
0,720 0.17 -135.24 85.35
0.722 3.77 -128.22 74.33
0.723 0.10 -139.1, 82.14
0.728 0.03 -139.78 81.85
0.735 0.07 -148.16 82.03
0.748 8.49 -144.58 74.55
0.839 0.10 -172.18 63.95
0.893 0.07 -10 .42 46.19
0.922 0.03 -193.18 27.46
0.931 3849 -191.82 21.83
0.945 52.21 -199.3
0.954 499.00 -192.9 -
0.957 235.80 -193. -
0.964 865.50 -196.6 -


*1clntive to Benzene
7,At half height







100% LicI04


LiC104.3H20--


1202


100%7. HK


Figure 2. LiC104-HO02-H20 System at 300C.









100% NaC104


NaC104.H21


o H202


100% H20


Figure 3. NaC104-H202-H20 System At 300C.






































100%


Figure 4. KC104-H202-H20 System At 300C.


H202










































)% H202


100% H20


Figure 5. RbC104-11202-H20 System At 300C.











































H202


100% H20


Figure 6. CsC104-H202-H20 System At 300C.









































Figure 7. Mg(Cl10)2-H202-H20 System At.300C.
A represents Ig (C104)2.2H20.H202
0 represents Mg(C104)2.H20.H202


70% H20


% H202








. 100% Ca(C10)2


Ca(CI04)2.4H20-


Figure 8. Ca,


i2-H202-H20 System At


100% 1120,


100% H202





































100%


H202


Figure 9. Sr(C104)2-H202-H20; Ba (C104)2-202H-20; NH4C104-H202-:
Systems At 300C.
O Sr(ClO4)2, 0 Ba(C10O2, fNH4C104





















10.00












-t


;5.00

04













0.0(


Mole Fraction H202 In Solvent


Figure 10. -


Plot of Molality of Salt Versus Mole Fraction H202
In Solvent for the System LiCI04-H202-H20 at 300C.


0.50


0.00


1.00





49














20.00



0








O0



10.00



0 ,Y











0.00,
0.00 0.50 r.OC


Mole Fraction H202 In Solvent


Figure 11. Plot of Molality of Salt Versus Mole Fraction H202
In Solvent for the System NaC104-H202-H20 At 300C.












































0


)


0.00 0.5 1.0


Mole Fraction H202 In Solvent


Figure 12. Plot of Molality of Salt Versus Mole Fraction H202
In Solvent for the System KC104-H202-H20 At 300C.


0.500


0.250


n nnn


a'
pr
,,


V,





















0.50


0.25


V.*u 0.50 1.00

Mole Fraction H202 In Solvent

Least Squares Line: Y = 0.377X + 0.065

Figure 13. Plot of Molality of Salt Versus Mole Fraction H202
In Solvent for the System RbC104-H202-H20 At 300C.





















1.00

o

/











o 0.50

-4












0.00
.00 0.50 1.00

Mole Fraction H202 In Solvent


Figure 14. Plot of Molality of Salt Versus Mole Fraction H202
In Solvent for the System CsCl04-H202-H20 At 300C.























2.00










0

o


1.50

00
o






1.00







0.00 0.50 1.00


Mole Fraction H202In Solvent



Figure 15. Plot of Density Versus Mole Fraction H202 In
Solvent for the System LiC104-H202-H20 At 30C.

























2.00













0
r-4

0

4 1.50
0



a)













1 nn


8.50 6.00

Molality of LiC104


Figure 16. Plot of Density Versus Molalty of Salt for the
System LiC104-H202-H20 At 30 C.


3.00


1 00
























1.50


0
00
O



oo


0
) z
o

tI

o 1.25

-'-I



0 0












1.00
0.00 0.50 1.1

Mole Fraction H202 In Solvent


Figure 17. Plot of Density Versus Mole Fraction H202 In
Solvent for the System KC104-H202-H20 At 300C.







56
















1.50






0





o
01.00
0.00 0.25
4 0

:-
0

o 1.25

















1.001
0.00 0.25 0.

Molality KC104


Figure 18. Plot of Density Versus Molality of Salt for the
System KC104-H202-H20 At 30 C.


































































0.50


Mole Fraction H202 In Solvent


Figure 19. Plot of Density

Solvent for the


Versus Mole Fraction H202 In

System RbCO04-H202-H20 At 300C.


1.50













0
-,4
u
,-4
o

S1.25
o

4J
o 1













1.00
1.00


0.00


1.00








58

















1.50

O O
O

0

0



0



0
-4
0 1.00













the System RbC4-H202-H20 At 3025
0





1.00- 0--____________________
0.00 0.50

Molality RbC104


Figure 20. Plot of Density Versus Mblality of the Salt for
the System RbClO4-H202-H20 At 300C.




















1.60


60










*o/





4.1

o 1.30
.rl

Q)
















1.00 -
0.00 0.50 1.00

Mole Fraction H202 In the Solvent


Figure 21. Plot of Density Versus Mole Fraction H202 In
Solvent for the System CsClO4-H202-H20 At 300C.







60
















2.00












O
*r4
0 0

S1.50
0

.4



o











1.00
0.00 0.50 1.

Molality CsC104


Figure 22. Plot of Density Versus Molality of Salt for the
System CsC104-H202-H20 At 300C.























1.80


1.75


0

00 000 00
0
0 o 0o o


1.601 I


0.00


0.50


1.00


Mole Fraction H202


Figure 23. Plot of Density Versus Mole Fraction H202 In
Solvent for the System NaC104-H202-H20 At 300C.




















































1.50


Figure 24. -


A Representation of the Diffraction Patterns
for Salt A, Salt B and Mg(C104)2.220














+100.0(















C)

C: 0.00
0*





Ov


0 0










o1a
o


-200.0(


0.00 0.50 1.00
Mole Fraction H202

Figure 25. Chemical Shift in the Proton Resonance Spectrum of
H202-H20 Mixtures
Least Squares Line: Y = 108.03 315.17X

























100.00





0

0




0
oo








.0
i-4













0.00) I
0.00 0.50

Mole Fraction H202


Figure 26. Width of Proton Resonance Line Versus
Mole Fraction H202














DISCUSSION OF RESULTS

1. Solubility Data


Vapor pressure measurements and solubility measurements by

Turner (7) and Everhard (20) have indicated that the size of the

cation is the major factor affecting the solubility relationships

btutcn water and hydrogen peroxide as solvents and the alkali metal

nitrates as solutes. The general trend was that the larger the cation

the greater was the solubility in pure hydrogen peroxide. Their data

for solubilities is tabulated in .Table nineteen along with the

standard radii (21) of the cations.


TABLE 19

IIOLALITY IN H202 FOR THE GROUP Ia NITRATES



Cation Standard Radius (1) Molality of Salt*


Li+ 0.607 2.48
Na 0.958 4.12
K 1.331 12.34
Rbn 1.484 14.51
Cs", 1.655 13.04

*These values were taken from a plot of morality versus ionic
,radii, presented in Everhard's dissertation.




The interpretation of this trend is greater solvation of the larger

cations by hydrogen peroxide. Further support of this conclusion was









the isolation of peroxyhydrates with INO03 and RbN03.

The trend with the perchlorates, however, does not indicate that

the morality of the salt increases in hydrogen peroxide with increase

in the size of the cation. Rather, it appears that the molality

actually decreases:


TABLE 20

MLKALITY IN 1202 FOR THE GROUP Ia PERCHLORATES




Cation Standard Radius (1) bolality of Salt


Li+ 0.607 8.28
r 0.958 4.21
I+- 1.331 0.43
+bt 1.484 0.43
Cs+ 1.656 1.13




However, the solubility data (Figures ten through fourteen) clearly

indicate that the solubility in pure hydrogen peroxide compared to

that in water increases with the size of the cation. This trend

matches that of the nitrates. The fact that no solvates of KCl04

and RbCI04 were obtained analogous to those of KN03 and MbllO can be

explained as due to the temperature at which the determinations were

carried out. All work on the perchlorates was done at thirty degrees

centigrade, while the highest temperature at which the nitrates were

studied was twenty-five degrees centigrade.

Everhard (20) obtained a linear relationship between cation

radius and differences in molality of salts in H202 and 120.









Calculation of the same quantity for the perchlorates gave the

following data:


TABLE 21

MIOLALITY DIFFERENCES IN 1202 ANID H20 OF GROUP la PERCHLORATES



Cation Radius 1olality bolality m'
In H20 In 120


Li+ 0.607 11.90 8.28 -3.62
Na1 0.958 22.28 4.21 -18.07
K' 1.331 0.18 0.43 -0.25
RPJb 1.484 0.09 0.43 40.34
Cs' 1.656 0.10 0.13 +1.03




where m' = m(1202) m(Ii20). It must be noted that the solubility

values for LiC104 and HaClO4 are not merely the result of solubility

measurements on solutions of NaC104 and LiC104. The solubilities

presented here represent those of anhydrous salts and in the case of

LiCIO4 and NaCIO4 the solid phases in equilibrium with the saturated

solutions in water have been shown to be, respectively, LiC104. 3120

and IcC104.120 (Figures two and three). In order to obtain the

actual molalities of the solutions in equilibrium with anhydrous salt

it was necessary to extrapolate the lines in the plots of morality of

salt versus mole fraction 11202 (Figures ten and eleven). In the case

of NaC104 this extrapolation is rather short and can be assumed to be

reasonably accurate. However, for LiC104 the extrapolation is rather

long and the accuracy of the value is questionable. To circumvent

this difficulty the following thermodynamic cycle was used to calculate








the solubility of anhydrous LiCIO4 in H120:

LiClO(c) +- 320 Ar Li+(aq., al) + C10l(aq., a,)



LiClOz.3H20n(c) Fq Li'(aq., 3) +- ClO~(aq., a3)

where al represents the activity in a solution which is in

with anhydrous LiC104 and a3 represents the activity in a s

which is in equilibrium with LiC104.3H20.


equilibrium

solution


Since 1 and 3 represent equilibrium conditions AF1 and &AF3

both must be equal to zero. This means that AF4 = A 4F. AF

can be calculated from the thermodynamic expression AFl = AI -

T AS, if AHI is kno n and aS~. is known or can be estimated.

Tarkowitz (22) has determined the value of iH2 and reports it as

-14.14 Kcal/mole. at twenty'five degrees centigrade. Latimer (23)

has tabulated values of entropy contributions to solids for fifty-

seven elements and thirty-two negative ions. He also states that the

entropy of hydrates may be estimated by assigning the value 9.4 cal./

deg. to the contribution of a mole of hydrated water. He cites

experimental data which gives good agreement with entropy values

calculated in this manner. The use of these values gives the entropy

of LiC104 at twenty-five degrace centigrade as 29.5 cnl./deg. mole

and for iC104.31120, 57.7 cal./deg. mole. Thus the entropy change

for reaction 2 in the cycle must be equal to -21.95 cal./deg. It

can be assumed that this value is essentially the came as that at

thirty degree centigrade. The value for AH2 can also be assumed to

be constant over the temperature range involved since invectiCgtors (24)

prior to aZrkowitz had detcrninn-d An at eighteen degrees centigrade









and found the value to be -14.20 Kcal./mole, in good agreecnt with

Markowitz' value of -14.16 Kcal./mole.. Substitution of the value for

the entropy change into the expression for &AF, along with the value

for AH, gives a value of LFW = -7.51 Kcal. at thirty degreee

centigrade. From this value it should be possible to calculate the

activity of LiClO1 in the hypothetical solution which would result if

it were possible to have anhydrous LiClO4 in equilibrium with Li'(aq.)

and C104 (aq.), the situation described by reaction 1 from the

relationship AI- AF = 2RTalna(1)/an.(3). Since the solutes on

both sides of reaction 4 are in aqueous solutions their standard

states are the same, so AFg = 0 and Ar4 = 2RTlna..(1)/aj.(3).

Solving this equation for al(1)/a+(3) results in a value of 5.09X102.

In order to calculate the desired activity of LiC1O4 from the

above ratio it is necessary to have the mean ionic activity coefficient

of LiC104 in the solution in equilibrium with LiClO4.3H20, i.e.

that represented by equilibrium 3. The morality of this solution is

given in Table six as 5.296. Jones (25) has reported mean ionic

activity coefficients for aqueous solutions of LiClIO up to 4.0 molal.

A plot of log 1V versus molality, although a non-linear plot overall,

is linear from 2.0 molal to 4.0 roll. Using the equation for this

straight line it is possible to estimate the desired mean ionic activity

coefficient. The result is a value of 4.003, which results in a value

of 1.08X104 for the mean ionic activity of LiC1O4 in a solution

in equilibrium with anhydrous LiC104. The plot of log a versus molality

is linear from 3 to 4 rolal so it is poccible to estimate the molality

of this solution from the mean ionic activity which has been calculated









covDe. The mraolity cctinmtod in thick -.:y turn out to be 7.54.

The rc c energy of transfer froam 110 to I202 '.a-:ld bo a better

meCaure oZ thev difcrcnczC botwoen 1yJro cn pero-xide mad wator ae

solvents then. tmh. .L rencc in ~. olubl.itin. Ths tr nsfer can be

reproecanitd by tthe following cycle:


2C3(c) --"
gr t

ro If ( 1C02**) Cr+ c (-2 2,T. )

then &PFt = r 7, hiclh also -cui *2.-Tln p;.:.;() .

:a ,cc only u2ai..t ica aco knom in inOM per cent ,I70 cad the menal

ioric activity co.-fl'-,:.:t.r cr ro.e : 1hnoia, the CLYC:.ic:L nrmus be

mcadc thi- the m mn ionic Cti:'.-ity coe2F.:.cl-.in: in both cozvcat-. aro

equal. Thirs rccults 11 the equation, at 2L.Tin z-(l)/i(2), If

thi. transfer p-:occ..: Uore purely dependent on cl..i;rottatl e intr-

.-.clos. 1.-:'-a saolute nad ~ olvent it iould be possible to dtcrI:. .,Le

the c:n ` rei cf Pft on the cationic r'.ii by a suitable form of

'ih- Dorn cr-.ti (26). or the reaction ion(g) ;= Zon(coln.) the

CIE,=c I C"_rn equation has the foera:
A?. -.2-/)
2r

th~rc Z ic the cihrge- on the oon r e '.aonic radius and D the

dielectric conostnt of the cmdium, Of course this does not include

cntropy chrlncec, i.thri- only electrical ork. The process of tranfer

fro one Rcol-cnt to aT other c-n aloo be r cpr Z....,jd by Ithe rfo!.l.'L-?nG

theiZ c- .,n-.'-ic cycle:










~f(g) + c10(g) AF2 '-(IO202) +- ClO (0202)




I' (H20) + C104 (1120)


For this cycle '6Ft = AF + AF62. Since the anion, C010, is the

same in all the salts being considered here the contribution to the

free energy changes from it are constant from salt to salt. As a

result of this, these contributions need not be carried along in this

treatment. AF1 describes the opposite of the process depicted in

Born equation so it is represented by:


AF1 -=- 2 (1/D)
2r

and since AF2 represents the process described in the Born equation;

Z2
F2 ___(-l/D2)
2r,


Z2 2
thus &Ft A + AF2 "- (1-1/D1) (1-1/D2)
2r. 2r+


"2
r-


S at p ( 22r. 1
2r..

This equation predicts, from purely electrostatic principles, that

the free energy of transfer from one solvent of dielectric constant

D2 should increase as the function l/r+ increases.










AFt;'s calculated for Group la perchlorates from the molalities

are tabulated in Table twenty-two (for the morality in 1120 for LiCIO,,

the calculated value, 7.54 molal, was used) along with the Pauling (21)

standard radii, in angstroms, and corresponding values of /r...


TABLE 22

4Ft FOPR TIE GROUP Ia PERCHLORATES



o
Cation r+(A) 4 (Kcal.) /r()



Li 0.607 -0.113 1.647
Na 0.958 +2.008 1.004
K 1.331 -1.049 0.751
Rb 1.484 -1.885 0.674
Cs" 1.656 -2.922 0.604




Figure twenty-seven is a representation of these data. A linear

relationship between AFt and I/r+ is evident, excluding the data

for Li. The fact that the value of AF; for Li' is considerably

below what it should be can be explained as due to contributions

caused by partial covalent bonding between the lithium ion and hydrogen

peroxide. In order for this to be true the polarizability of hydrogen

peroxide must be greater than that of water. Little study has been

made of intermolecular forces in hydrogen peroxide such as ion-dipole

or dipole-induced dipole attractions (27). However, Gorin (28) found

hydrogen peroxide to be salted in to a nurbcr of electrolytic solutions.

This phenomenon was explained as due to the replacement of H120 around



















+ 2.


0.00


- 2.00


l/r+ in 9-'


Figure 27. Plot of
Cationic


AFt Versus the Reciprocal of the
Radius









the ions by H202 as a result of the higher dipole moment of hydrogen

peroxide, compared to water, in the vicinity of the ion, The theory

proposed was that the.hydrogen peroxide molecule assumes a planar

shape as it approaches an ion.. This would result in a dipole moment

of 3.0lXl"018 as compared to 1.84Xl0"18 for water,

In addition to the above observation, the polarizabilities

per molecule of t102 and H20 have been determined (29) at twenty-

five degrees centigrade from refractive index data and it was found

that the polarizability per molecule for H202 is approximately 1.6

times as great as that for H20,.

Moelwyn-Hughes (30) has given the following equation for the

state of minimum energy (the most stable state) of an ion surrounded

by c eqidistant molecules of solvent at a distance ao from its center:


-uo c( 5B 7+
(9ao 9a )

where B is constant depending on the polarizability of the solvent

molecule, and C is a constant depending on the dipole moment of the

solvent molecule. In other words, all other things being equal, the

solvent molecule with the higher dipole moment and polarizability

will have the lower energy when taking part in ionic solutions. The

fact that so appears in the denominator and is raised to the fourth

power in the first term indicates that, due to the polarizing power

of lithium, this term will be much larger than for other ions with

less polarizing power. This is also the case in the second term.











The polarizing power of the Li ion is well known and is often

used to explain anomalous behavior of lithium, e.g. the high

relative instability of Li2CO3 compared to the other alkali metal

carbonates (31) is explained by the high polarizing power of the

Li+ ion as compared to the other alkali metal ions. The fact

that the free energy of transfer for Li+ ion from water to hydrogen

peroxide is considerably lower than predicted by a purely electro-

static approach is apparently another example of the anomaly of

lithium. The fact that this anomaly was not noted by Everhard (20)

in his study of the nitrates is perhaps explained by his failure

to reduce all his data to one reference state, i.e. the anhydrous

salt, before making any comparisons. A second cause could be the

use of solubility differences as a standard for comparison, rather

than energy differences. Of course, agreement with Everhard and

Turner (7) is apparent in the case of those ions which do fit a

purely electrostatic model. It is evident from Figure twenty-seven

that ions with smaller radii than potassium ion favor solubility

in H20 rather than H202 (but not lithium) and that ions the size

of potassium ion and larger favor solubility in 1202.

The solubility data for the alkaline earth salts are not

detailed enough to enable rigorous comparisons analogous to those

made for the alkali metal salts. Table twenty-three presents the

data for the solubilities of the alkaline earth salts in water

and in the H202-H20 mixture of the highest concentration of 1202
for which data was obtained. The mole fraction H202 in the

solvent in these latter solutions is designated by the value in

parentheses.










TABLE 23

1MOLALITY OF GROUP IIa PESRCILORATES IN 1120
AID R202-I20 HIXT-URES


Salt Molality in 20 MIolality in 11202
Rich Solution

1((C104)2 4.55 7.44 (0.82)

Ca(C104)2 8.18 3.83 (0.78)

Sr(C104)2 10.26 6.30 (0.58)

Ba(C104)2 6.21 4.27 (0.90)



Since the molalities in O20 refer to saturated solutions in

equilibrium with their respective hydrates no direct comparisons

can be made. The one conclusion which can be drawm from this table

is that MI(C104)2 seems to behave analogously to LiCli4, its

solubility increasing as the transfer is nade from water to hydrogen

peroxide. This would be expected due to the similarity in size

between lithium and magnesium.










II. Density Data


Strong attractions between ions and solvent molecules should

have an effect on the volume of the solution, and, as a result, on

the density. It is difficult to draw any conclusions from Figures

fifteen through twenty-two because the separate trends are not

segregated. It is known that the densities of mixtures of water and

hydrogen peroxide vary with composition. It is also logical that

the density should increase as the concentration of salt increases

since this process adds weight to the system.

Nevertheless, it should be possible to define a parameter which

will give some insight into the effect of a solute on the solvent.

An apparent molar quantity is defined as

Y n1y
y2
n2
where Y is the property of the mixture which is being studied (in

this case, the volume of a given amount of solution), nl and i2

represent moles of substance 1 and 2 respectively, and yl represents

the molar value of the property being studied, for the pure substanca.(32),.

From the tabulated densities and density plots the values of the

apparent molar volumes of the solutes in pure water and pure hydrogen

pcro::ide can be calculated by the above equation. The value for he

density of pure water, at thirty degrees centigrade, is 0.9957 gnc./ml.

(19), and that of pure hydrogen peroxide is 1.4363 Smr./ml. (33).

These values result in 18.08 ml./mole for the molar volume of

pure water and 23.67 ml./mole for that of hydrogen peroxide. The data

in Table twenty-four represent the result of calculations of apparent










molar volumes for the Group la perchlorates in pure water and pure

hydrogen peroxide. Also tabulated is the difference between the

two, a measure of the contraction in volume in passing from one solvent

to the other. It is calculated as


Vo202 vO

TABLE 24

APPARENT MOLAR VOLUME OF GROUP Ia PERCILORATES
IN SATURATED SOLUTION


Solute v1o (ml.) VH202 (ml.) (VH202 V20



LiC104 45.59 38.58 7.01

NaC104 49.01 49.40 + 0.39

KC104 55.58 47.67 7.21

RbCIO4 87.86 76.81 -11.05

CsC104 55.23 48.28 6.95




The above trend is practically the came as that of the free energy

of transfer calculated in the preceding section. Figure twenty-eight

is a plot of AFt versus (v1202 V120) and shows a linear relation-

ship, with the exception of Cso. Li falls between K" and Na'.

Such a trend might have been predicted from the relationship between

F and pV: P = U TS + pV. Since the cycle for the calculation of

AFt has the following form:
















+2.00 NaT














0.00 -

Li0o





I< / OK4t





/Rb4+
-2.00






O 0 Cs+
-10.00 0.00
VH11202-VH20

Least Squares Line Y = 0.343X + 1.861
(with exception of Cs')

Figure 28. Plot of aFt Versus Apparent Molar Volume
Difference (VH202-VH20)









C104o (c) A M 202) 1- c10 (H202)
2A/(H202)



P(H20) + C10 (120)



then AFh F AP2 A0 F

SF= U T A S +p V

AFt a U\ T AS2 + p V2 ( &.U T A S' + p V1)

or &Ft = T S' aU' p(&Vl V2)


which would give a linear trend, under the assumption that T S' -

A U' is a constant. However, if it is assumed that T AS' A U'

represents the Y intercept of the graph values calculated by this

equation would be much lower than those obtained by calculation from

experimental data. This infers that T aS' AU' also is a linear

function of apparent molal volume difference. The general equation

for &Ft; is bFt = a + b(vII202-V120). Comparison of this with

aFt = T &S' AU' P(V1 V2) results in T AS' AU'

a + b(vHO2-vI2). Substitution gives AFt = a + b (vH202-H20)

PCV1-V2) or AFt = a + (b-p)(vH202-vH20)

The fact that the cesium ion does not fall on the line is in

agreement with Turner (7) who found that the deviation of the molal

solubility of the alkali nitrates, between hydrogen peroxide and

water at twenty-five degrees centigrade, was a linear function of

the radius, except for cesium nitrate. This was explained by the

assumption that the smaller charge density of the large cesium ion









results in less of an interaction with hydrogen peroxide than does

the charge density of the smaller ions. However, the actual apparent

molar volumes must be considered here in addition to the difference

between the two. It is noted in both solvents that a general trend

seems to be increased molar volume with increasing size of cation.

However, the values for Cso- are very much lower than those for Rb+

and, in fact, approximate those of K'. This can be explained by

noting that the apparent molar volume would not only be affected

by the strength of the attraction between the ion and the solvent

but also by the number of such attractions. The Cs+ ion is so large

it can have more solvent molecules adjacent to it, than the preceding

alkali metal ion, thus a contraction in the solution is caused which

would be greater than if cesium had the same number of solvent

molecules adjacent to it as rubidium had.

Moelwyn-Hughes (30) has tabulated values for the coordination

numbers in water of the alkali metal ions, excluding lithium. The

values indicate that there definitely is a chnnge in coordination

number in going from rubidium to cesium. The value for rubidium

is 5.6 and that for cesium is 6.2. Assuming the same trend to hold

in hydrogen peroxide the above argument would appear to be valid.

Another interesting result of the calculation of apparent molar

volumes is that the lithium ion again shows itself to be anomalous,

contracting the hydrogen peroxide solution almost as much as the

potassium ion. Similar arguments to those applied in the discussion

of free energy of transfer apply in this case also. The fact that

lithium does contract the hydrogen peroxide solution, compared to








82


water, is further support for very strong interaction between the

lithium ion and hydrogen peroxide.










III. Isolation of Solvates With Hydrogen Peroxide


Examination of the partial ternary phase diagram for the

Mg(C104)2 H202 H20 system (Figure seven) indicates two inflection

points. The first point is at 55.2 per cent Ig(CI04)2, 36.2 per cent

11202 and 8.6 per cent H20 and the second is at 58.6 per cent I~(C104)2,

36.5 per cent 11202, 4.9 per cent H20. The data for the hydrogen peroxide

rich region of the latter inflection is limited to one point which

extrapolated to somewhere between the calculated percentage corres-

ponding to the monohydrate. Actually this point is most logically

interpreted as extrapolating to a salt having the composition

Mg(C104)2.120.H202, which is designated on the phase diagram by a

square.

Concerning the other inflection point it is noted that on the

left of this inflection point (hydrogen peroxide poor region) the

solubility of magnesium perchlorate is decreasing with increasing

hydrogen peroxide concentration; on the right (hydrogen peroxide

rich region) the solubility is increasing with increasing hydrogen

peroxide concentration. The first irprccsion gained from the wet

residue data is that the solid in equilibrium with solution on both

sides of this inflection point is 11(C104)2.2H120 since this is the

approximate composition at which the extrapolations of the solution -

wet residue lines intersect the axis. However, on the hydrogen

peroxide rich side of the diagram the extrapolations also pass

approximately through the point corresponding to the composition

Ig (C104)2.2H20.H202, so that this latter formula could represent the









composition of the solid phase as well, in the 1202 rich region.

In an attempt to reconcile this apparent dilemma, samples of wet

residue were taken directly from the system flasks and x-ray

diffraction patterns run on them. The d-spacings, relative

intensities, and angle of reflection (29) for a sample from the

hydrogen peroxide poor side of the inflection point (Salt A) and one

from the hydrogen peroxide rich side (Salt B) are presented in

Table fourteen. These patterns are also represented in Figure

twenty-five along with a representation for a salt prepared in this

laboratory and analyzing as 7Zg(C104)2.2120, The intensities cannot

be given much weight in this comparison since the method of

preparing the samples (see experimental part) possibly did not

completely randomize the directions of the crystallites scanned by

the beam of x-rays. However, a peak by peak comparison of these

three patterns indicates that they are different in some respects

and similar in others. It seems just as logical to identify Salt B

with the dihydrate as it does Salt A (which is the salt that should

be the dihydrate).

In order that a comparison among the three patterns be easily

made they are tabulated below along with two other patterns which

have significance, i.e. a pattern obtained by IHnawalt (34) in 1933

for what he presumed to be 1' (C104)2.3H20 which salt was later proven

not to exist (35); and a pattern reported by Smith and Rees (36, 37)

as the pattern for NI(C04)2.2H120. These latter two salts are

designated by the investigator's name in the table.

















TABLE 25


X-RAY PATTERNS OF TnET RESIDUES ANI'D g(C104)2.2H202


Salt A Salt B 1 (C104)2.2H20 Rees IHanaIlt
d int. d int. d int. d int. d int.


6.75 S. 6.9 M. 6.9 V.11.


4.87 M. 4.89
4.74
3.98
3.46
3.39 V.S. 3.41
3.35 M. 3.35
3.21 1. 3.23
3.18
3.10 V.W. 3.08
3.03 VW. 2.93

2.78
2.60 V.-1.
2.58 V.11. 2.58
2.48
2.43
2.34
2.27 V.H. 2.25
2.17 V.W.


2.04
2.01
1.90
1.80
1.71
1.62


V.W.
V.11W.
V.W.
H.
13w,
V.W.
V.11.
v~lw.
vSw.


S.
V.S.
V.S.
S.
M.
C,
II.
S.
M.


4.90
4.74
3.93

3.42
3.36
3.24


S.
V.S.
M.
14.
M.

S.


W. 3.11
1U. 2.96
2.86
H. 2.78
2.60


6.9
5.5


V.H.
W.


4.72 V.S. 4.80 V.S.
3.95 S. 3.99 S.


3.36 V.S. 3.35 S.
3.22 V.S. 3.22 V.S.


3.06
2.95
2.85
2.77 M. 2.75


2.43
2.34 V.W. 2.33 H. 2.34
2.25
2.18


2.05 W. 2.05 V.H.
2.04 V.H, 2.04
2.00 W. 2.00 V.W. 1.97
1.91 V.W. 1.91 W. 1.89
1.84 1. 1.84 11. 1.82

1.63 V.1U. 1.63 V.1. 1.61


V.H.
.W
V.TJ.
V.11.



V.W.
V.W.
V.JL
V.W.
V 11.


H.

S.
S.

V.1.









A line by line comparison of these patterns indicates that

there are lines missing in one which appear in others and vise-

versa. If these were only weak lines then perhaps a comparison

could be made but they are in many cases very strong lines. It

would be expected that Salt A, Salt B and g((C104)2.2H20 would have

very similar patterns since the data in Table fifteen indicate that

the three hydrates of IM(C104)2 and the anhydrous salt have similar

patterns differing by the deletion or addition of a few lines in

each case. This infers that the addition of water to the crystal

structure does not have an overwhelming effect on the crystal structure.

The major difference between Salt A and the dihydrate is the apparent

absence of two very strong lines, 4.47 and 3.98, which, however, are

present in Salt B. The main difference between Salt B and the

dihydrate is the presence of a strong line at 3.46 and a medium

line at 3.10, both of which are missing in the patterns for the

dihydrate and for Salt A. The fact that these lines are present is

a more forceful argument than the absence of the other two lines in

Salt A since it is possible that due to the way the samples were

prepared these lines could have been missed. The presence of these

two lines in Salt B infer that something has been added to the salt

which, in this case, would be the H202 molecule. However, in the

final analysis it seems that very little can be drawn from these

patterns alone. The main conclusion which can be drawn is that

Salt A is different from Salt B. This is what one would expect

because of the inflection point in the curve, on the phase diagram.

Further support for the fact that this actually is the mixed

solvate 1g(104)2.2H120.H1202 is the fact that the next transition is







87

unquestionably to IMg(CI04)2.H20O.I202. If the salt iriaediately

preceding this transition contained no hydrogen peroxide of solvation

it would rea more likely that the transition would be to anhydrous

I-,(C104)2, since no monohydrate of magnesium perchlorate is known.

The above two solvates tcre the only ones for which any evidence

was obtained in the systems studied. There were indications of

strong interactions in other solutions (e.g. apparent molar volume

calculations) but no solvates were isolated containing hydrogen

peroxide. This is most likely due to the fact that thirty degrees

ccntigrade is above the decomposition temperature of any solvates

which might be capable of existence. Lower temperatures would favor

the existence of phases containing appreciable amounts of solvent.











IV. Nuclear Magnetic Resonance Data


It would seem likely that extensive solvation by protonated

species of ions would alter the electronic atmosphere around the

protons in the solvated species and thus give rise to some alterations

in the proton magnetic resonance spectrum of the species being

solvated. In an effort to discover if this were the case a study

was made of the proton magnetic resonance spectrum of solutions

of magnesium perchlorate in 1120 and 11202-1120 mixtures. As a basis

for comparison the position of the combined proton magnetic resonance

line of mixtures of water and hydrogen peroxide was determined as a

function of hydrogen peroxide concentration (Figure twenty-five) as

well as the widths of the proton magnetic resonance lines at one-

half the height of the line.

The data in Table eighteen indicate that small concentrations of

magnesium perchlorate in 202-1120 mixtures have more of an effect on

the proton magnetic resonance spectrum of such mixtures than do

larger concentrations. This effect also seems to be most pronounced

at an approximate mole fraction for hydrogen peroxide of 0.73.

There is a similar effect at higher concentrations of 11202, but it

does not seem to be so pronounced; there seems to be no effect at

lower concentrations of H1202.

It is possible that the chemical shift which appears to be

caused by Ig(CIO4)2 could actually be due to oxygen in a supersaturated

solution, produced by decomposition of 11202. It is known that oxygen

dissolved in samples being investigated for nuclear magnetic










absorption can affect the position of the signal in the magnetic

field (38). Paramagnetic species such as oxygen can also alter the

width of the nuclear magnetic resonance signal. Both of these

effects were observed in the dilute solutions of magnesium perchlorate,

as the data in Table eighteen indicate. Further, the direction of

the effects were such that they could be attributed to oxygen.

The presence of significant quantities of oxygen in solution

could be explained if catalytic decomposition were caused by the

magnesium perchlorate. In order to explain the fact that more

concentrated solutions of Mg(C104)2 do not show a similar effect it

is necessary to postulate that the presence of large amounts of

magnesium perchlorate in solution can catalyze more effective nucleation

of the oxygen molecules produced by the decomposition, thus increasing

their rate of escape from the solution. This would result in a

lowered steady state concentration of oxygen, which could approach

the equilibrium concentration existing in the solutions which contain

no magnesium perchlorate.

The alternative to catalysis of the decomposition by the

magnesium perchlorate is that the magnesium perchlorate does not

enter into the process at all, but that there is a certain optimum

concentration of hydrogen peroxide where decomposition is most

rampant. This theory, however, is contradicted by the fact that

the plot of resonance line position versus mole fraction hydrogen

peroxide, for the solutions not containing magnesium perchlorate

(Figure twenty-five) is linear over the complete range of 202 -20

concentrations.









Although the effect of magnesium perchlorate on solutions

of hydrogen peroxide and water was not that expected it seems

likely that the mechanism of decomposition catalysis by magnesium

perchlorate would involve solvation of the magnesium ion by H202.

This would lead to weakening of the 0-H bond and the production of

the species H30 This "peroxonium" ion has been postulated as

responsible for catalysis of the decomposition of hydrogen peroxide

in acid solutions (39). The method of production of this ion would

be directly analogous to the production of H30'' in aqueous solutions

of -, i.e. (1h0)6 + I120 = t (120)5(OH)*- + 130+. Such an inter-

action seems probably for magnesium perchlorate since its solubility

increased in hydrogen peroxide, indicating strong attractions

between magnesium perchlorate and hydrogen peroxide, and also on

the basis of the evidence for the existence of g((C104)2.2120.1202

and I (C104) 2. I20.202.

The fact that the effect of magnesium ion seems to be the

greatest at approximately 0.73 mole fraction hydrogen peroxide may

be related to the fact that the plot of line width versus mole

fraction 11202 indicates that there is a minimum in the rate of

exchange of protons between water and hydrogen peroxide at

approximately 0.73 mole fraction 1202. This is based on the

assumption that line width is inversely proportional to the rate

of exchange in a simple manner. In order to say whether this is

so a study would have to be made of the kinetics of hydrogen exchange

between hydrogen peroxide and water. Anbar et. al. (18) have

investigated the kinetics of hydrogen exchange between hydrogen









peroxide and water by proton magnetic resonance methods in the

pH range of 2.5 to 6.5, and found the reaction to be both acid and

base catalyzed. Their investigations, however, covered only a

limited range of rather low hydrogen peroxide concentrations

(0.00 0.20 mole fraction 11202). The data obtained and tabulated

in this work (Tables sixteen, seventeen, and eighteen) were,

unfortunately, not of a sufficiently quantitative nature to permit

analyses analagous to those made by Anbar. The line width data

obtained were merely taken as an adjunct to the study of solutions

of magnesium perchlorate so no particular attention was paid to

the refinements necessary to obtain data suitable for kinetic

studies.




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