THE EFFECT OF TARTARIC ACID ON THE
KINEHCS CF THE GXEDATEON OF
URANIUM (W) IY THALLEUM (III)

EN AQUEOUS PEECHLORATE MEDIA

Thesis for flu Dogma 0’ pk. D.
MICHEGAN STATE UNIVERSITY
Calvin Miles Love

1964

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‘ ABSTRACT
"L": £14.; ,
‘THE EFFECT OF TARTARIC ACID ON THE KINETICS

OF THE OXIDATION OF URANIUM(IV) BY THALLIUM(III)
IN AQUEOUS PERCHLORATE MEDIA

by Calvin Miles Love

The kinetics of the oxidation of uranium(IV) by thallium(III) in
the presence of tartaric acid were investigated in aqueous perchloric
acid solutions. Increases in initial rate with increasing tartaric
acid concentrations were observed, along with departure from apparent
second~order kinetics. The orders with respect to uranium(IV), thal—
lium(III), tartaric acid, and hydrogen ion were found to be 0.93, 0.05,
0.37, and -0.91, respectively. Thallium(III) in the presence of
tartaric acid was shown polarographically to be stable, yet under kin—
etic run conditions where thallium(III) is in excess over uranium(IV),
the reactions stopped because the thallium(III) was consumed. The
oxidation of uranium(IV) by oxygen of the air is catalyzed by the pres—
ence of tartaric acid, but this reaction is negligibly slow compared to
the uranium(IV)~thallium(III)-tartaric acid reactions. The presence

of oxygen, uranium(VI), and thallium(I) were shown to have no effect on
the reaction kinetics. A mechanistic interpretation has been developed

to account for the observed results.

     

. I)
.N-

   
 

. or TARTARIC ACID ON THE KINETICS
.L .'.-. , ‘ I,
. tfi'm summon or URANIUmIv) BY THALLIUM(III)
IN AQUEOUS PmmLORATs MEDIA

By

Calvin Miles Love

A THESIS

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

DOCTOR OF PHILOSOPHY

Department of Chemistry

 

 

”W DEDICATION

In memory of my father,
who taught me the value of

education and knowledge.

 

ACKNOWLEDGMENTS

    

jg: I wish to express my sincere appreciation to Professor Carl H.
ffi‘flrubaker, Jr. for his valuable guidance and friendly counsel throughout
N the course of this investigation.

Financial assistance from the Atomic Energy Commission is grate-
fully acknowledged.

Appreciation is also extended to my parents who provided some
financial support which made family living on a graduate assistantship
reasonable.

Special thanks go to my wife, Sue, for her patience, understanding,

and steady encouragement which helped to make this entire project

possible.

 

  

D.
E.

 

i III.

‘ THEORETICAL . . . . . . . . . . . . . . .
IV. EXPERIMENTAL . . . . . . . . . . . . . .
A. Preparation and Standardization of Reagents .
l. Perchloric Acid . . . . .
2. Sodium Hydroxide . . . . . . . . . .
3. Uranium(VI) Perchlorate . . . . . .
h. Uranium(IV) Perchlorate . . . . .
5. Potassium Permanganate . . . . .
6. Nitrogen . .
7. Thallium(I) Perchlorate
8. Thallium(III) Perchlorate ......
9. Potassium Bromate
lO.

pow

JVIWaI-oognoono
‘HISTORICAL

Aqueous Uranium Chemistry . . . . . . . . . .

W
+2?
.

“new mm

Aqueous Thallium Chemistry . . .
Electron Transfer Reactions of uranium and

1.
2.
3.
h.

Previous Work on the Uranium(IV)

Thallium Exchange Studies . . . . . .
Thallium(III) Reductions .....

Reaction . . . . . . .
Oxidations of Tartaric Acid . . . . . .

11.

Sodium Perchlorate . . . . .
Tartaric Acid .

The Uranium(IV)- Thallium(III) Reaction

Studies of Uranium Solutions Without Thallium .

l.
2.

Studies of Thallium Solutions Without Uranium

Spectrophotometric Studies of Uranium
Polarographic Studies of Uranium . .

a o c a a n o I a

—Thallium(III)

  

lium . ... . . . . . . .
Uranium Exahange Studies . . . . . . . . .
Uranium(IV) Oxidations . . . . . . . . . .

Measurement of Uranium(IV) Concentrations . . .

l. Spectrophotometric Studies of Thallium . .

2. Polarographic Studies of Thallium' . . . .

3. 2 ,h- -Dinitrophenylhydrazine Experiments . .
Polarogra hic Studies of the Uranium(IV)-

 

(P

Thallium III) Reaction . . . . . . . . . .

 
  
    

   

1 f It. 115111.375

A. Kinetic Observations . . . . .

WV ONY‘LP‘M’ A) H

9.
10.

B. Observations of Uranium Solutions
C. Observations of Thallium Solutions

DISCUSSION

SUMMARY . . .
LITERATURE CITED . . . . .

Merénoe‘ Kinetic Runs . .

TABLE OF CONTENTS (Cont.)

Effect of Water Quality . .

Effect of Iron(III) . . .

Effect of Dissolved Oxygen .

Tartaric Acid Variation
Uranium(IV) Variation
Tiallium(III) Variation .

Effect of Uranium(VI),'Iha11ium(I),.and

Solution Age . . . .

Hydrogen Ion Variation .......

Polarography of Kinetic Runs . .

APPENDIX -- ORIGINAL KINETIC DATA . .

a o o o a o n I n c

      
 

 

  

VIII I
IX.

 

Unaccelerated reference runs .
Accelerated reference runs .
I Iron(III) accelerated runs .
Nitrogen— and oxygen-swept runs
Tartaric acid variation runs .
Uranium(IV) variation runs .

Thallium(III) variation runs ..

Hydrogen ion variation runs

grandpa ions in aqueous solution . ..
Concentrations of stock solutions

Uranium(IV) standard solutions . .

Uranium(VI), thallium(I), and aged runs

Half—wave potentials of the uranium(IV)— (III) wave as
a function of tartaric acid concentration .

Diffusion current as a function of time for a thallium-
(III)-tartaric acid solution . . . . . . .

Polarographic monitoring of thallium(III) concentration
in reference runs . . . . . . . . . .

   

92

95
115
116
117
119
120
122
123
12h
125

126

 

LIST OF FIGURES

Figure Page
1. Formal potentials of uranium in l M HC104 at 250 C . . 10
2. Formal potentials of thallium in l M HClO4 at 25° C. . lb
3. The reciprocal of uranium concentration as a function

of time . . . . . . . . . . . . . . . . . . . . . . 33
h. Uranium(IV) apparatus . . . . . . . . . . . . . . . . no
5. Nitrogen purification train . . . . . . . . . . . . . A7
6. Thallium(III) apparatus . . . . . . . . . . . . . . . 51
7. Absorbancy XE“ uranium(IV) concentration for

uranium standards . . . . . . . . . . . . . . . . . 60
8. Typical second-order graphs . . . . . . . . . . . . . 70

9. Apparent second—order rate constants XE' iron(III)
concentration . . . . . . . . . . . . . . . . . . . 73

10. Absorbancy vs. time for some tartaric acid accelerated
runs . . . . . . . . . . . . . . . . . . . . . . . . 75

ll. Logarithm of initial rate vs. logarithm of initial

tartaric acid concentration . . . . . . . . . . . . 76
12. Absorbancy vs. time for uranium(IV) variation runs . . 78

13. Logarithm of initial rate vs. logarithm of initial

uranium(IV) concentration . . . . . . . . . . . . . 79
1h. Absorbancy vs. time for thallium(III) variation runs . 80
15. Logarithm of initial rate vs. logarithm of initial

thallium(III) concentration . . . . . . . . . . . . 82
16. Absorbancy XE' time for hydrogen ion variation runs. . 85
17. Logarithm of initial rate vs. logarithm of hydrogen

ion concentration . . . . . . . . . . . . . . . . 86
18. "Polarograms" . . . . . . . . . . . . . . . . . . . . 87
19- Thallium spectra . . . . . . . . . . . . . . . . . . . 9h

 

 

I. INTRODUCTION

Since about 1950 many investigations in the field of inorganic
chemistry have been directed toward the understanding of the kinetics
and mechanisms of electron transfer reactions. Within this class of
reactions, which includes both oxidation—reduction and electron exchange
processes, the mechanisms of a variety of reactions have been elucidated.
It appears that only in gaseous systems is there clear evidence1 for
direct transfer of an electron from reducing agent to oxidizing agent.
In the liquid state the situation is more complicated and there are
sound arguments? for believing that such a process cannot occur with
ions in solution. Although the nature and sequence of the elementary
steps comprising many overall electron transfer reactions in solution
have been explained, the more basic problem of just how the electrons
are transferred still remains.

Examination of numerous electron transfer reactions, particularly
electron exchange reactions, has shown that no single type of mechanism
can account for all the observed kinetics. However, these studies have
resulted in the recognition of at least two broad classes of mechan—
isms. By considering the type of activated complex involved, Taube3
has designated these two distinct classes as "outer—sphere" reactions
and "bridged" (or "inner-sphere") reactions. In the reactions involv—
ing an outer—sphere activated complex, the number and identity of the
groups comprising the first coordination spheres remain unaltered.on
electron transfer. In reactions proceeding via a bridged activated

complex, a common group is shared by the metal ions, so that a change

 

     
   

t’ebagiiifiétion‘ sphere of at least one of the partner's era's '

   

l
"itgtivated complex are, in general, those reactions that are conSidered

tekproceed by an electron tunneling mechanism. The results for some
rgactions of this type have been summarized by Wah1.4 The conclusion
is that the direct transfer of electrons by a tunneling mechanism is
believed to occur between very similar ions that are fully coordinated
with non—labile ligands, such as between hexacyanoferrate(II) and hexa-
cyanoferrate(III) ions5 and between manganate and permanganate ions.6
Here the electronic environments are almost identical and the species
are nearly the same size; thus only a small amount of energy is required
to form the transition state in which both species have the same bond
lengths and in which the electron is then exchanged. The theory for
electron transfer by a tunneling mechanism was developed by Weiss7 and
by (R. J.) Marcus, Zwolinski, and Eyring.e More recently (R. A.) Marcus9
has contributed to this theory and succeeded in predicting free energies
and entropies of activation for various reactions which are in satis—
factory agreement with experimental data.

In bridged (or inner—sphere) reactions, electron transfer is pre-
ceded by substitution into the coordination sphere of one of the ions,
with the formation of a bridged species in which the reactants.are ‘
linked by a common ligand. The conditions for detection of this
bridged species, which may be either an activated complex or an inter—

mediate, have been clearly described by Taube.3 He pointed out thwhjy;_hz» Q

 

   
     
  
   

i an, a direct demonstratipn of abriflep WW
, pamlex 3(or intermediate) can be, made by simple. H

E gflireaction products. So by judicious choice of chromous
;1' §Q;(a reducing agent, Taube and co-workers were able to demonstrate

1 §ggfin§nsfer of a large number of univalent atoms and groups.4° Since

     

l'qppgmium(ll) complexes are labile to substitution and chromium(III)
species are quite inert to substitution, any atom or group coordinated
in a product chromium(III) species must have been transferred from the
oxidizing agent during the electron transfer process. In this manner
the transfer of chlorine atoms has been demonstrated when the oxidiz-
ing agents were CrClz+,11 FeC12+, AuC14-, and Co(NH3)5C12+.1° By using
a variety of chemical and tracer techniques, Taube and his co-workers
have further demonstrated complete atom or group transfer to chromous

ion according to the equation:
C0(NH3)5X(3+“)++ Cr2+ + 5H+ —> CrX(3+n)++ Co2* + 5NH4+

where X = 1", Br‘, I‘, 5042',10 NCS', 113', P043“, P2074“, acetate,
butyrate, crotonate, oxalate, succinate, maleate,12 and hydroxide,13

and where n = the charge on X. It is significant that, of the organic
anions in the above list of X groups,.the two (oxalate and maleate)
which provide a continuous pathway of conjugated n bonds between the
metal ions allow electron exchange to occur about 100 times faster than
do the others. Considerably more work than can be briefly detailed here

has been done on the kinetics of this system with various carhoxylato

    

ligands, 14 15 with the general result that electron transfer through

  

A"Li_. conjugated systems has been clearly demonstrated.

   

    

‘.-,.—18 tracer experiments have shown that oxygen atom transfer occurs
flit the oxidation of sulfite with hypochlorite, chlorite, chlorate, and
machete and in the oxidation of nitrite with hypochlorous acid.16 In
.further studies using oxygen-18 labeling, oxygen atom transfer is indi-
cated in reactions of aqueous uranium(IV) with various oxygen-contain-
ing oxidizing agents17 and in the reaction between uranium(VI) and
chromium(II).16 In a study of the isotopic exchange reactions of phos-
phorous pentachloride, the exchange between phosphorous pentachloride
and labeled chlorine is believed to proceed by an atom transfer mechan—
ism.19

Some evidence has been adduced to support a mechanism of electron
transfer between aquo and hydroxo metal ions which is based on the
transfer of a hydrogen atom between the hydration shells.2°,21 For
example, the rates of the Fe2+-Fe3+ and Fe2+—F'eOH2+ reactions are
slowed by a factor of two in heavy water compared to ordinary water;21
such a reduced rate is characteristic of a reaction in which a hydrogen
atom moves. Furthermore, a large number of metal aquo oxidation-reduc—
tion reactions have an activation energy close to 9 kcal./mole and an
activation entropy close to -23 e.u., thus suggesting they proceed by
a common mechanism which probably involves water.22 Still another point
is that certain oxidation-reduction reactions of metal complexes seem
to require that at least one of the inner shell ligands be a water

molecule,22 for example, a cyanide ligand must be replaced by a.aatg;, 5 ~

‘3
5‘
g
5'
H
H:
2
0
i
‘3
8
Ft
3
E".
.“L
E
8
:3
8'
ii.
a:
8
a.
S
*5?

  

ltfleall this evidence, there is as yet no direct preqfcv

    
 

<v;n atom transfer from the hydration sphere of the reductant

feels and that this merely constitutes an effective path for the trans-
fer of an electron by "conduction."

To summarize, it can be said that the actual mechanics of electron
transfer reactions are, in general, rather poorly understood and that
no single type of mechanism can describe all electron transfer reactions.
In some cases the bridged transition state complexes are compatible
with explanations based on atom transfer, but in many cases it seems
more likely that the bridging phenomenon serves to facilitate tunnel—
ing or may permit "electron conduction" through the bridge.24 So the
question is posed:25 Does the formation of a bridged species merely
serve to hold the reactants together in similar environments and in
proximity until some other process can occur26 or does the bridge
serve as a path for electron conduction?Z7

In hopes of answering this question, work in this laboratory has
dealt with electron transfer reactions in which anions can serve as
possible bridging groups between the oxidizing and reducing species.
In this manner, we expect to obtain information which will help distin-
guish among tunneling, atom (or group) transfer, and conduction as

paths for the transfer of electrons. We have studied the effect of

   
 

     

sulfate,ze chloride,29 and various organic oxy-acids30 on the Tl(I)- “ ‘9 ‘

T1(III) exchange. The Sb(III)-Sb(V) reaction has been examined in

  

fl-:_Sulfuric—hydrochloric acid mixtures31 and the Sn(II)-Sn(IV)

  

; f insurable acidsolutlpns‘.” The acceleration of":

     
    

Filfifili‘e (III) exchange"3 and the HOW—”(VD exchange“ by

”il'INEV)SU(VI) exchange, now with respect to photo-catalysis.35 We
-13'#§”also reported a study on the exchange between the lZ—tungsto—

' tgiohhaltateUI) and l2-tungstocobaltate(III) anions.35 With respect to
Poxidation-reduction reactions, we have studied the kinetics of the
lSn(II)-Ce(IV) reaction in sulfate media37 and also the effect of some
organic acids on the U(IV)-Tl(III) reaction.36

Our work has shown that anion bridges often produce pronounced
kinetic effects and we believe that, in most cases, conduction and/or
tunneling is involved in the electron transfer. Thus it seems par-
ticularly important to evaluate the effects of various kinds of anion
bridges and to study, in detail, the mechanics of some of these anion-
bridged electron transfer reactions.

The work by Quinn,36 which will be discussed in detail later, con-
cerning the effect of some organic acids on the U(IV)-Tl(III) reaction
showed that under appropriate conditions tartaric acid accelerates the
reaction. Considerable difficulty in interpreting the data was encounter-
ed. As a result, it was suggested that this system be examined in de-
tail. The problem then defined was to study the effect of tartaric
acid on the kinetics of the oxidation of uranium(IV) by thallium(III)

in aqueous perchlorate media, with the goal of elucidating the reaction

mechanism.

    

 

II. HISTORICAL

Some of the current thinking concerning the mechanisms of electron
transfer reactions has already been discussed. In addition, there are
a number of reviews on this subject.2,14:25,39‘47 Consequently, it is
the intention of the present chapter to report only previous work in
this area that is directly concerned with uranium(IV) and/or thallium(III).
To establish a background, the chemistry of aqueous uranium and thal-
lium solutions will be briefly examined. Then uranium and thallium
electron transfer processes will be reviewed, followed by an examina-
tion of previous investigations of the uranium(IV)-thallium(III) re-
action, including a discussion of the effects of various anions and

cations on this reaction. oxidatlxm of tartaric acid will also be dis—

cussed.
A. Aqueous Uranium Chemistry

The literature on the chemistry of uranium ions in solution is
voluminous since the subject deals with hydrolytic reactions, oxida-
tion-reduction equilibria, and complex ion formations of the ions of
this important element in all their various oxidation states. Discus—
sions on uranium chemistry, including its solution chemistry, are given
by Rodden and Warf,46 by Hoekstra and Katz,49 and by Katz and Seaborg.5°
The most up-to—date and comprehensive review on the analytical aspects

Of uranium chemistry is to be found in "Treatise on Analytical Chem-

istry."51

A

W-
8
Uranium in solution exhibits four well—defined oxidation states:
+3, +h, +5, and +6. Of these, the tetra— and hexapositive states are
most important. For the various oxidation states of uranium in aqueous
acid solution, Table I displays data on the principle ionic species

present, their color, and a brief statement on their stability.

Table I. Uranium ions in aqueous solution

 

 

Oxidation Ion in

 

State Solution Color Stability
+3 U3+ Deep red Slowly oxidiged by H20, rapidly
by air to U4
+h U4+ Deep green Stable; slowly oxidized by air to
U022
+5 U02+ Usually a Disproportionates to U4+ and U022+;

muddy brown stability greatest pH 2-b
(?)

+6 U022 Greenish— Very stable; difficult to reduce
yellow

 

Uranium ions have characteristic absorption spectra. They ex-
hibit sharp peaks reflecting Sf electron transitions.52 As a consequence,
environmental factors which affect the outermost orbitals are without
material effect upon the wavelengths of these bands; thus spectra in
the presence of complexing groups are strikingly similar to those of
the simple (aquated) cations. Measurement of the absorption due to the
various oxidation states of uranium without addition of special reagents
is useful for following changes in oxidation state equilibria. The
technique is also useful for establishing the stability of samples.
Despite the extensive studies of uranium in solution, no completely

definitive set of spectra appears in the literature. Stewart53 and

 

    
    

' .7 ..repart the absorption Spectrum of the» mfim)
iris acid. Kraus and Nelson-55 and Cohen and Cart-ell“

has, h95, 550 and 650 mp. Absorption spectra of uranium(V) (except in
:hvdrofluoric acid solution55) apparently have not been published,
probably due to the instability of this oxidation state in solution.
Among others, Arhland,57 Betts and Michels,55 and Silverman and Moudy59
have reported Spectrophotometric investigations of uranium(VI) in
perchlorate solutions.

Aqueous solutions of uranium salts have an acid reaction due to
hydrolysis which increases in the order U3+ < U023+ < U4+. The degree
of hydrolysis increases with increasing charge and decreasing ion size;
however, the uranyl ion which has a charge of +2, nevertheless behaves
in hydrolysis more like an ion of charge +h. Due to the ease of hydrol—
ysis, a high hydrogen ion concentration is necessary in order to reduce
the extent of hydrolysis and to maintain monomeric species in solution.
With decreasing hydrogen ion concentration, the extent of hydrolysis
increases, then polymerization of the hydrolyzed species occurs, and
finally precipitation of hydrous oxides results.

A great many complexes of uranium are known. Some of these will
be mentioned below under discussions of the chemistry of the individual
oxidation states. The extent of complex ion formation in a particular

case can often be studied by one or more of the following methods:‘ . >~‘3

spectrophotometry, potentiometry, polarography, distribution equgv:

 

 

5

10

The details of these methods are thoroughly described by Rossotti
and Rossotti.6° All of these methods have been applied to uranium
complexes.

The formal potentials of the various oxidation states of uranium
in acid solution are shown in Figure 1.51:52 The potentials are strongly
affected by complexing with anions, but nevertheless they are of great
value in predicting the oxidation—reduction behavior of uranium. The
U3+/U4+.and U02+/U022+ couples are reversible; the U4+/UOZ+ and U4+/U022+
couples are irreversible. This is in accord with the criterion for
the reversibility of a couple, which requires that no metal-oxygen

bonds be formed or broken during oxidation—reduction.“5

Figure 1. Formal potentials of uranium inl.fl HClO4 at 25° C.

U +1.80 113+ +0.631 U4+ -o.58 uoz+fluozz+

I —o.32

Aqueous solutions of uranium(IH) can be prepared by electrolytic
reduction of the +h or +6 species or by dissolving a uranium(III) halide.
Solutions containing the intense red +3 ion are unstable and oxidized
by water with the formation of uranium(IV) and hydrogen. Uranium(III)
is hydrolyzed, but the exact extent of hydrolysis is difficult to deter-
mine because of the oxidation reaction, however, qualitative evidence
based on pH measurements of uranium(III) solutions indicates hydrolysis
of U3+ is about that expected for a +3 ion. Information on complexes
of uranium(III) is scant, but it appears to behave like the lanthanides.

Uranium(IV) solutions are most conveniently produced by electrolytic

reduction of uranyl solutions or by dissolving a uranium(IV) halide.

t.

 
      
   
 

' ' “mouthwash erodes}: greenrare quite stable but
‘ .' nth; dissolved oxygen; if protected from air they can by -

  

, Towering periods of time. Although uranium(IV) undergoes exten—
.§h§vu hydrolysis, the existence of the unhydrolyzed species U4+ has

‘ Alldn unambiguously established.55 Hydrolysis of U4+ produces, at least
in the early stages, a monomeric species:

H“ + H20 <—'—'> UOH3+ + H+.

Kraus and Nelson55 found a value of 0.05 for the equilibrium constant,
RU, at 250(3 at ionic strength 0.19, and a value of 0.21 at zero ionic
strength, while Bett554 found a value of about 0.075 at 2h7O‘C at an
ionic strength of 0.19. For an ionic strength of 3.9 at 25°C, a suit-
able value for this hydrolysis constant is 0.021. ySome work has been
done on the study of uranium(IV) complexes. Katz and Seaborg65 review
work in this area. Measured stability constants have been tabulated“,67
Pertinent to the present work, it has been noted that hydroxy acids

such as tartaric, malic, quinic, citric, lactic, and glycolic form
complexes with uranium(IV) ions,58,69 but apparently the literature
contains no reports on the stability constants of uranium(IV) tartrate
complexes. Adams and Smith70 claim to calculate a stability constant
for the uranium(IV) chelate of citric acid. Since Th4+ is similar to
U4+, it is interesting to note studies on the complexes of thorium with
tartaric acid.71'75 These studies propose formulas for the complexes
formed in various acidity ranges and under various thorium and tartrate,

concentrations; and even some stability constants are reported. mot:

      

  

' " somewhat However, complexes with marinate“

 
     
   
 

'13“: or 1.1 and 1:2 are indicated and Katzin and Gulyas73

i‘filfi.thorium complexes bitartrate under fairly acid conditions.

‘ ‘39“ "uranium(V) is the least stable oxidation state of uranium in

'. Ifilution. The pentapositive ion, U02+, has a transitory existence

under most conditions, although evidence for its occurence has been

3 ‘ obtained polarographically.75‘78 Kraus and co-workers79,5° have
studied in detail the chemistry of aqueous uranium(V) solutions and have
shown that the U02+ ion is most stable in the pH range 2.0-h.0, where

the disproportionation reaction:

 

2U02+ + hH+ <———> 114* + U022+ +2H20

is very slow. Nelson and Kraus81 report the equilibrium constant for
this reaction to be 1.7 x 105. Millimolar solutions of U02+ can be
prepared by electrolytic reduction of U022+, by reduction of U022+

With hydrogen, zinc amalgam, uranium(III), chromium(II), or europium(II),52
or by dissolving uranium pentachloride.5° Heidt and Moon63 have shown
that uranium(V) ion is an intermediate in the photochemical oxidation
reactions of uranium(VI). The study of the hydrolytic behavior of U02+
is difficult due to the narrow pH range where U02+ is reasonably stable.
Apparently the hydrolysis constant for U02+ is considerably smaller than
10-3. This low extent of hydrolysis is consistent with the properties
expected for a large unipositive ion. Very little data is available

on complexes of uranium(V)

  
  

Hexapositive uranium is the familiar stable oxidation state ofdthisf,t'

     

7‘2. element in aqueous solutions.. Uranium(VI) in solution exists as ,3;

    

that is probably disturbed by strong local flatly“

   
  

. containing uranyl ion are bright greenish-yellow in color‘and

: may acidic as a result of hydrolysis. The exhaustive studies

oral agreement that hydrolysis leads to the production of polynuclear
complex species, 5.3., U2052+, U3052+, U306(0H)+, upswing, 0303(0103‘,
and U305(0H)42-. In solution, the uranyl ion forms complexes with a
very large number of anions. Work on uranium(III) complexes has been
reviewed by Katz and Seaborga6 and stability constants have been tabu-
lated.65,57 A number of studies have been performed on the uranyl-
tartrate system.‘57'92

One final item to be mentioned regarding uranium chemistry is its
polarography. This topic, which has been the subject of extensive in-
vestigation, has been thoroughly reviewed.93:94 Polarographic reduction
of uranium proceeds in as many as four steps: (VI) to (V), (V) to (IV),
(IV) to (III), and (III) to uranium amalgam. The (VD-(V) and (V)-(IV)
waves are obtained under conditions in which uranium(V) is stable.
The (VD-(V) wave in perchloric acid occurs around -0.17 v. E“ the
saturated calomel electrode, 8.0.5. The reduction of uranium(V) to
(IV) involves the breaking of oxygen bonds and is therefore acid-
dependent. Under conditions in which uranium(V) is not stable, a
pseudouranium(VI)—(IV) wave is obtained. This apparent uranium(VI)—
(IV) wave is due to the cycle of reduction of uranium(VI) to (V) with

subsequent disproportionation of the (V) to (V1) and (IV). The (330‘s:

         

(III) wave is clearly obtained in many solutions but is seldom u.
. for analysis because of its high negative potential (+0.9 to - ,
M "" .~.svisu1'..-tant..m11. plateau: . stanzas isms: ' -' .

E.
4!

 

. I

   

1 . , m
. I’iiflhfiiflfifimfiifliabfiui-thfl:Ollbififififiihhidej=-‘Ha:;.,.
, _ " ‘ .9 I}: to uranium anal-gain is fgemrany not. OhWMK
.fiwlfltfigpuns beyond the hydrogen wave at about —2 v. . y- e

      

      
 
 
 
 

 

h

     
      
 
 

B. Aqueous Thallium Chemistry
y gvuy‘

,;;i hycr The chemistry of aqueous thallium solutions is not nearly as com—

. ”pdex.as aqueous uranium chemistry since thallium exhibits just two com-

paratively stable oxidation states: +1 and +3. Standard treatises and

reference books offer brief discussions on thallium chemistry. The

. analytical chemistry of this element has been recently reviewed.9-",96

Thallium is in group III of the periodic table, so the trivalent

state is important. However, lower oxidation states become more prom-
inent in the heavier members of non-transition groups; thus the uni-
valent state of thallium is likewise important. In aqueous solutions
thallium(I) salts are more stable than thallium(III) salts. Indeed,
the Tl(I)-Tl(III) relationships are the dominant feature of the chem-
istry of thallium. The formal potentials of the thallium system in

acid solution are shown in Figure 2.97:98

Figure 2. Formal potentials of thallium in 1 fl HClO4 at 25° C.

T1 _0_3fi_ T1+ __L2h7_ no“

[7 -o.719

The electrode otential for the Tl+/Tl3+ couple clearly indicates
P

      

the stability of univalent thallium with respect to oxidation to {dingo ' 5;:

       
    
   

4' _sp§;i%Qly soluble, whereas the fluoride exhibits app:eciab‘le
, m sloflubiylity. Thallium(I) forms an insoluble chromate and sul-
.3}: the other hand, like the alkali metal ions, it shows compara-
inuhavqoly little tendency to form complex ions in solution. Thallium(I)
hydroxide, TIOH, is a strong, fairly soluble base. Thallium(I) can be
(”fluidized in basic solution to insoluble thallium(III) oxide with hydro-
‘ ; gen peroxide, potassium hexacyanoferrate(III), sodium hypochlorite, or
potassium permanganate. Potassium bronate or cerium(IV) sulfate can be
used in the presence of hydrochloric acid to oxidize thallium(I) to
thallium(III). Thallium(I) salts are oxidized by chlorine, bromine,
or aqua regia. An excess of chlorine or bromine can be removed by
boiling, but prolonged heating is to be avoided, for otherwise reduction
of thallium(III) ions takes place. Thallium(III) solutions free of
foreign ions can be conveniently prepared by anodic oxidation of thallium(I)
solutions.99:1°° It should also be noted that thallium(I) solutions
are exceedingly poisonous and trace amounts cause loss of hair.
Thallium(III) is colorless and exists as T13+ in strong non—com-
plexing acid solution. It resembles aluminum ion except that it is
larger. However, thallium(III) oxide, T1203, is not amphoteric, but
purely basic. It appears to form no hydrates and so cannot be formu—
lated as the hydroxide. Thallium(III) oxide can be made by heating
the metal in air below a red heat or by hydrolysis of a neutral or
basic thallium(III) solution. The sesquioxide is. extremely insoluble

:(with a Ks sp of 1.5 x 10 4‘ when written as the hydroxide). 97 It is

 

  

    

 
   
 

iiguighnlliun(111) ions are reduced conveniently by Sulfur
i) .;
Iodide ion or hydrogen peroxide in acid solution will also
h..;.

'm:however prolonged boiling is necessary to remove eXcess sulfur

‘ 1?? to the metal.

Soluble thallium(III) compounds are highly hydrolyzed. Hydrolysis

of T13+ can be written:

+

T13 + H20 <:> T10H2+ +H+.
In perchlorate media of ionic strength 2.9, a suitable value for the
equilibrium constant, KTl’ for this reaction at 25°C is 0.073, as given
by Biederman.1°° With decreasing acidity, further hydrolysis yields
T1(0H)2+, only mononuclear complexes being formed.”1

Thallium(III) has a strong tendency to complex with both inorganic
and organic liqands. In addition to hydroxide complexes mentioned above,
studies have been performed on thallium(III) complexes with nitrate,
sulfate, cyanide, halides, acetate, ethylenediaminetetraacetic acid,
various amines, and a variety of other organic reagents. The review by
Busev and Tiptsova95 briefly discusses these studies and cites the
numerous references. Stability constants have been tabulated“;67
Pertinent to the present work are the studies on tartrate complexes of
thallium(III). Pyatnitskii and Kostyshima102 have examined thallium(III)
tartrate complexes relative to those of aluminum, gallium, and indium.
Busev, Tiptsova, and Sorokinam3 report the overall successive stability
constants of the tartrate complexes of thallium(III) at unit ionic
strength and at 20° C to be $1 = 3.7 x 1011, $2 = 6.5 x 1012, and
.g gaff 2,2 x 1013, These values, however, were determined in peak ~.:

l ,.,, .

999““ P“ h-

 

  
   

 

 
  
  

“lflitibowe,'thallous and thallic ions are colorless ipdeiéfiyi‘

 

aa‘ho't absorb‘i'n the visible portion of the spectrum; They

:lfeerature contains several reports of spectrophotometric studies of
Cfiallium in perchloric acid solutions.1°1:1°4f1°6 Rogers and W'aind106
report changes in the ultraviolet absorption spectrum of thallium(III)
in perchlorate solutions due to acid, ionic strength, and temperature
variations, and also the effect of using heavy water. This study also
indicates the feasibility of determining thallium(I) and thallium(III)
concentrations Spectrophotometrically by utilizing differences in their
molar absorptivities at particular wavelengths. It is clear however,
that if other absorbing species (such as uranium or tartaric acid) are
present, the method would not be suitable.

Thallium polarography has been studied to some extent. Most of
the investigations deal with reduction of thallium(I) to thallium amal-
gam; and this work has been reviewed.1°7,1°a Reduction of thallium(I)
produces well—defined polarographic waves in different supporting
electrolytes, and the values of the half-wave potentials of these waves
are nearly constant, around —0.h7 v. gs. the S.C.E. The literature
contains only two reports on the polarography of thallium(III). Hughes
and Hush1°9 report that thallium(III) solutions in 0.1 to 1.0 fl HC104

and 0.01h g in NaN03 at 25° C (with 0.00h% sodium methyl red as a 1‘

  

maximum suppressor) display a reproducible and well-defined diffusion

current at zero volts 3;. the S.C.E., which corresponds to the two

 

electron reduction to thallium(I). The only wave found occurred atlas

 

 

1‘ also v. in 0.1 g limo4 and at -o.h71'v. in 1.0 game.“ this;

 
 

 

18
identified with the Tl(I)-T1(Hg) step. A more extensive study of thal-
lium(III) polarography was done by Smith and Nelson,110 but they chose
to work in chloride solutions. Their work is compatible with Hughes'
and Hush's, and furthermore they determined diffusion coefficients and
showed that polarograms from solutions containing both thallium(I) and
thallium(III) were suitable for analytical determinations of thallium(I)
and thallium(III).

Before leaving the subject of aqueous thallium chemistry, one more
item should be mentioned. That is the work by Catherino and Jordan111
on bivalent thallium. They studied the Tl(I)—Tl(III) electrode reaction
by hydrodynamic voltammetry and presented evidence that the electro—
oxidation of Tl+ and the electroreduction of Tl3+ proceed via a bivalent
thallium species. They showed that in l M HClO4 the rate of dis—
proportionation of thallium(II) is relatively fast compared to the elec-
tron transfer steps:

Tl(II) + e' —> Tl+

T13+ + e“ —> Tl(II).
This accounts satisfactorily for the fact that bivalent thallium has
never been prepared as a stable species in quantities amenable to chem-

ical characterization.
C. Electron Transfer Reactions of Uranium and Thallium

From previous discussions, it should be clear that from a stoichio-
metric point of view electron transfer reactions are of two types.
Those in which an electron is transferred from one metal ion to another
of the same element are called electron exchange reactions. When the

metal ions are of different elements, the process is called oxidation- ’

‘

19
reduction. To be di5cussed in this section are: (1) studies on uranium
exchange, (2) uranium(IV) oxidation reactions, (3) studies on thallium
exchange, and (h) thallium(III) reduction reactions.

It should be pointed out that kinetic and mechanistic studies are
often conducted in perchlorate media because of any anion, perchlorate
shows the least tendency to form complexes. Only at extremely high
perchlorate concentrations is there evidence for the formation of metal

ion perchlorate species.59

1. Uranium Exchange Studies

The electron exchange between uranium(IV) and uranium(VI) has
been studied by Betts112 in sulfate solutions; by Rona113 in chloride
solutions; and by King,114 by Masters and Schwartz,115 and by Benson34
in perchlorate solutions. Bettsllz studied the reaction in the presence
of constant external illumination and suggested that the active inter-
mediate in the exchange is uranium(V). Rona,113 who studied the exchange
in hydrochloric acid, proposed a mechanism after finding that the re—
action was first order with respect to uranium(VI), second order with
respect to uranium(IV), and minus third order with respect to hydrogen
ion concentration. She also observed no effect on the rate of exchange
with added inert salt, with chloride ion concentration, and with illum—
ination. Kingll4 reported that the exchange was slow and suggested
that this was due to the formation and breaking of metal—oxygen bonds.
Masters and Schwartzll5 studied the exchange in perchlorate media and
showed that in addition to Rona's path, there is another path which
predominates at low uranium(IV) concentrations and at temperatures of

25° C and higher. This second path was found to be first order in

i‘teaViolet irradiation. Benson34 studied the acceleration of the

 

we shawed that increasing tartaric acid concentration, increasing ionic
strength, increasing temperature, and ultraviolet irradiation all pro—
duced large increases in the reaction rate. He reports orders.of 1.3
in uranium(IV), O.h7 in uranium(VI), 0.90 in tartaric acid, and -2.9 in
hydrogen ion. A mechanism involving three paths for exchange is pro-
posed and calculated rates based on this mechanism are shown to agree
with experimentally determined rates.

The uranium(IV)-uranium(VI) exchange has also been studied in water,

 

ethanol, and water-ethanol mixtures by Mathews, Hefley, and Amis.116
This study has also been extended to water—ethylene glycol and water-
acetone solvent systems by Amis and co-workers.117 In all three systems
it is observed that the orders with respect to each reactant changes
radically as the composition of the solvent changes. Wear118 has

presented rate laws for the exchange in these three solvent systems.

2. Uranium(IV) Oxidations
The first really quantitative data to appear on a uranium(IV) oxida-
tion reaction was concerned with the uranium(IV)—iron(III) system. Bett3119 i}
i

examined this system in aqueous perchloric acid and showed that the

  
 
      
     

kinetics of the process are consistent with a mechanism which involves,
, in: f“
as a rate-controlling step, electron transfer between hydrolyzed i';}»".

  

4.» experimental evidence suggests, but does not prove, thatx§
H .hifi a;;‘ y . , . ‘.g j -7 2,

u-"y _ u

    

  
    
      
 

. 2:1;
war-asters“ + U(ai)23+—-_~a> premium-WM»
L‘?C _ygg§IU1ent path involving F‘eOH2+ and UOH3+. Baesizfi

  

  

' :ww:_uranimav) oxichtion by iron(III) in sulfuric and phos-

  

‘ ‘aeid solutions. He found uranium(IV) to be rapidly oxidized in
‘ éflgfihric acid and the reaction to be slower in phosphoric acid.
Kinetic studies on the oxidation of uranium(IV) by neptunium(VI),
t“ p1utonium(VI), and plutonium(IV) have been reported. The kinetics '.

of the reaction:

U(IV) + 2Np(VI) ——e> U(VI) + 2Np(V)
were studied in aqueous perchloric acid solutions by Sullivan, Zielen,
and Hindman121 and found to follow the rate law:
- d[NpOZ+]/dt = 2k[U4+][Np022+]/[H+].

Thermodynamic quantities of activation for the complex (UOH.NpOZ)5+ ~w

are reported, along with the effect of deuterium, ionic strength, and

the charge of the cation used to maintain the ionic strength. The

kinetics of the reaction between uranium(IV) and plutonium(VI) in

perchlorate media are reported by Newton.122 The hydrogen ion dependence

indicates that the reacting system passes consecutively through two

activated complexes: (HZO°U-OH-0Pu0)5+ and (HO-U-OH-OPuO)4+. Thermo-

dynamic activation quantities are reported and possible mechanisms are

discussed. Newton“:5 also reports a kinetic study on the reaction

between uranium(IV) and p1utonium(IV) in perchlorate media. The rate

law indicates that the activated complex is formed from water, U‘+, and

      
   

Pu4+ with prior loss of two hydrogen ions. Minor paths involving other

activated complexes appear unlikely. Thermodynamic quantities 01; AG

to acid greatly increase the rattler the ,:Jz'%ff

In"

22
Baker, Newton, and Kahn124 studied the kinetics of the reaction
between uranium(IV) and cerium(IV) in perchlorate media. The rate for
the principal reaction path was found to be proportional to both the
CeOH3+ and U4+ concentrations and inversely proportional to the hydrogen
ion concentration. Activation quantities are reported and a mechanism
is proposed.

The oxidation of uranium(IV) by molecular oxygen in perchloric
acid solution has been examined by Halpern and Smith.125 Over a con—
siderable range of conditions, the results are fitted approximately by
the rate law:

—d[U(lV)]/dt = k[U(IV)][Oz]/[H+].
The reaction is catalyzed by Cuz+ and inhibited by small amounts of Ag+
and C1-. The results are interpreted in terms of a chain reaction
mechanism involving U02+ and H02 as chain carriers. This appears to
be the first chain reaction in aqueous solution involving uranium ions
as chain carriers to be reported.

Some experiments were performed by Baker and Newton”6 to help
understand the reaction between uranium(IV) and hydrogen peroxide in
aqueous perchloric acid solution. The results show that the reaction

does not proceed directly by the simple net reaction:

+
J

U4+ + H202 —> U022+ + 2H
but that it is in part, at least, a chain reaction and that it is ac-
companied by a small amount of H202 decomposition. Possible mechanistic
interpretations are presented. The kinetics of the oxidation of uran—
ium(IV) by hydrogen peroxide in perchloric and sulfuric acid solutions

r

has been investigated polarographically by Sobkowski.137

  

 

 
       

I n n
‘ ”Whale“ 0f Willem) is sulfuric seminar W
, f‘i;;‘, hydrogen peroxide, and chlorite have been reported Hfi’
"1 95“]; and Fedorova.12° , _ .
‘,_:;ggxygen-18 tracer studies have been done by Gordon and Taubei7 on
'Jlthp.gxidation of uranium(IV) by various oxygen-containing oxidizing
' agents. The reactions studied were the oxidation of aqueous U4+ by
solutions of Mnof, (2042-, H202, and some halogenates; by solid samples
of PbOZ, MnOz, and NaBiOs; and by gaseous samples of Oz and 03. The
results are interpreted in terms of oxygen atom transfer accompanied
either by hydrolysis with solvent or by oxide ion transfer from the
oxidizing agent.

Studies on the oxidation of uranium(IV) by thallium(III) have been

done, but this work will be discussed separately later.

3. Thallium Exchange Studies
The electron exchange between thallium(I) and thallium(III) has
been the subject of much investigation. The first studies of this re-
action were doen by Zirkler129 and Majers,13° but the short half-life
(3.1 minutes) of the isotope employed (Tl—208) made their data diffi-
cult to interpret. Harbottle and Dodson“51 and Prestwood and Wah1132
then simultaneously reported, from studies using Tl—20h which has a
four year half-life, that this exchange reaction in perchloric acid
solution proceeded at a measurable rate. Subsequently, detailed kinetic
studies of this exchange were presented by Prestwood and wan1,133 Mc-
COnnell and Davidson,”1 Harbottle and Dodson,1°4 Dodson,134 and

Rossetti.135 In these studies, such aspects as the effect of 3'

 

  

 
   

   

”indicate that in perchloric acid the exchanging species are Tl+

     
   
 

 
 

both T13+ and T10H2+. A fairly recent re-examination by Roig and
a or'

{son1§‘ of the thallium(I)—thallium(III) exchange at various acidities

TA"

'1 IVII" '

f“; in perchlorate media shows that indeed the rate law can be written:
‘ Rate = k0[T1+][T13*] + k1[T1*][T10H2*].
Challenger and Mastersl37 have investigated the X-ray induced ex-
change between thallium(I) and (III) in 0.8 y H2504. Brubaker and
Nickel28 have studied the thallium electron exchange in sulfuric acid
solutions, as well as the effect of chloride on the exchange in this
medium.29 '
Brubaker and Andrade3° have reported the effects of some organic
oxy-acids on the electron exchange reaction between thallium(I) and
(III) in 3.50 E H0104. Reduction of thallium(III) is observed when
tartaric acid is added to the system, however, the extent of this
reduction is only significant after several days. Donasl:58 reports

another study on thallium exchange in the presence of organic acids.

h. Thallium(III) Reductions

The literature contains many reports on reactions of thallium(III),
but only a few detailed kinetic studies of metal ion reductions of
thallium(III) have been reported. Of these, the thallium(III)-iron(II)

reaction has drawn the most attention, namely by Johnson,139 Forchheimer

      
 
 

and Epple,14° and Ashurst and Higginson.141 The rate law which can be

written:

   
 

that the overall reaction must involve the.:

' J
-.7.. - . a.

25

k1

T1(III) + Fe(II) <:E:f> Tl(II) + Fe(III)
k2

T1(II) + Fe(II) "“> Tl(I) + Fe(III).

This system was further investigated by Duke and Bornong,14z who ex—
amined the reaction at high chloride concentrations.

The kinetics of the oxidation of mercury(I) by thallium(III) has
been studied by Halpern and co—workers.143 Here, the rate law:

-d[Hg(I)2]/ut = klkZIHg(I)2][Tl(III)J/k_1[Hg(II)]
for the reaction:

Hg(I)2 + Tl(III) ———+> 2Hg(II) + Tl(I)
points to the mechanism: k1 +
ngz+ <hz"? H92 + Hg°(aq)
-1

k
Hg°(aq> + T13*—2> Hg2+ + Tf’.

 

The kinetics of the reduction of thallium(III) by the trisdipyridyl—
osmium(II) ion has been reported by Irvine.144 The data can be inter-
preted by assuming that both T13+ and TioHZ+ react with Os(dipy)32+.

Higginson, Rosseinsky, Stead, and Sykesl45 have studied the reactions:

 

2V<III> + Tl(III) > 2V(IV) + Tl(I)
2V(Iv) + Tl(III) ———d> 2v(v) + Tl(I)

and discussed their mechanisms. The V(IV)-Tl(III) system was again
investigated by Sykes,146 now with the addition of excess Tl+ and found
to be independent of this ion. He was able to conclude that the
Tl(I)-T1(III) exchange, which is appreciable under conditions of his
experiments, is a single—stage two-electron transfer.

The reduction of thallium(III) by uranium(IV) will be discussed

in the next section.

26

D. Previous Work on the Uranium(IV)-Thallium(III) Reaction

The principal study on the uranium(IV)—thallium(III) reaction was
reported by Harkness and Halpern.147 They examined the kinetics of this
oxidation—reduction reaction, which involves a two—equivalent oxidant
and a two-equivalent reductant, with the objective of determining whether
reactions of this type occur in a single step or through successive one—
electron changes. Their results suggest, but do not demonstrate con—
clusively, that the reaction occurs through a single two-equivalent
step rather than through successive one—electron changes.

Harkness and Halpern147 made their kinetic measurements by follow—
ing uranium(IV) concentration with a thermostated Beckman DU spectro—
photometer using the 650 mu absorption peak. At this wavelength, absorp—
tion by all other ions in solution is negligible. They established that
the stoichiometry of the reaction in aqueous perchloric acid solutions
conforms to the equation:

U(IV) + Tl(III) ——> U(VI) + 11(1).
At constant hydrogen ion concentration and ionic strength, the kinetics
were found to be consistently of second order, first order in uranium(IV)
and first order in thallium(III). The apparent second—order rate constant,
k', as determined from the slopes of linear second-order kinetic plots,
were unaffected by variation of the initial uranium(IV) concentration
between 3.5 x 10—3 and 11 x 10—3 M and initial thallium(III) concentra-
tion between 5 x 10—:5 and 21 x 10—3 fl. Addition of uranium(VI) or
thallium(I) produced no effect. The ionic strength, u, was maintained
with sodium perchlorate. A typical kinetic run would have the following

conditions:

‘ i“
11'.

 

2?
IH+J = 1.76 m
u = 2.9
[U(Iv)]o = 3.5 x 10'3 g
[T1(III)]o = 9.0 x 10‘3 g
Temperature = 25° C.
Under these conditions, Harkness and Halpern147 find k' = 1.79 x 10—2
fl-lsec_l.

At constant ionic strength, k’ was found to depend inversely on
perchloric acid concentration in the range 0.5 to 2.8 fl. When fitted
to an empirical rate law of the form:

-d[U(IV)]/dt = k[U(Iv)][T1(III)][H+]‘“,
n was found to be 1.39 at 25° C. This non—integral inverse order in
hydrogen ion was interpreted in terms of simultaneous contributions
from two separate paths inversely first— and second—order, respectively
in hydrogen ion. The corresponding rate law is:

-d[U(IV)]/dt = [U4+][T13+](k1[H+]_l + k2[H+]‘2),

which can be rewritten:
k1[H+] + k2

-d[U(IV)]/dt = [U(IV)][T1(III)]([H+] + KU)([H+] + KTl)

 

Ru and KTl are the hydrolysis constants of U4+ and Tl3+, whose values
at 25° C and u ~ 3 are 0.021 and 0.073 E, respectively.55:1°° The last
factor in the above equation is just k', so a linear plot of k‘([H+] +
KU)([H+] + KTl) vs. [H+] yields k1 from the slope and k2 from the
intercept. Harkness and Halpern147 find k1 = 2.11 x 10‘2 sec_1 and

k2 = 2.13 x 10-2 M—lsec_1 at 25° C. The two rate constants are identi—

fied with reaction paths involving activated complexes of the compositions

‘

28
(U-OH-Tl)6+ and (U-O-Tl)5+, respectively. It is possible to interpret
the kinetics in terms of reactions involving the hydrolyzed ions U0H3+
and/or T10H2+.

The effect of ionic strength and the specific effects of various
anions and cations were examined. The anions were Cl- and 5042—; the
cations were Cu2+, Hg2+, and Ag+. In all cases the kinetics remained
first order in uranium(IV) and first order in thallium(III).

A brief mention of work on the uranium(IV)~thallium(III) reaction
was made by Mellon.148 He reports that research was conducted to
determine if the reaction proceeds by a chain mechanism; however, it
is noted that the work met with no success. Suggested experiments to
obtain information on this reaction include a study of the effects of
the copper(I) ion and a study of the system without argon treatment.

An investigation of the oxidation of uranium(IV) by thallium(III)
over an extended concentration range has been recently reported by Wear.149

A study of the electron transfer between uranium(IV) and thallium(III)
in water—methanol media is reported by Jones and Amis.15° The re—
action kinetics were studied at 1h, 20, and 25° C in 25, 50, and 75
Weight per cent methanol—water-perchloric acid solutions. The kinetics
of the reaction in these media did not follow second—order plots, as
in aqueous solutions, and the method of initial slopes was used to
determine the order with respect to each reactant. Orders with respect
to uranium(IV), thallium(III), and perchloric acid were found to de—
pend strongly on the solvent composition. Specific velocity constants
and thermodynamic quantities of activation are presented. A mechanistic

eXpression is developed to account for these orders in 25 per cent methanol

‘ i...

 

29
solutions. The effect of variation of ionic strength and the specific
effects of various anions and cations on the rate were examined in 75
per cent methanol solutions. The salt effects are strikingly different
from those found in aqueous solutions.

The uranium(IV)-thallium(III) reaction under Harkness' and Hal—
pern's conditions,147 and in the presence of various dibasic organic
acids was studied by Quinn.38 His work in the absence of organic acids
is in reasonable agreement with that of Harkness and Halpern.147 Under
the conditions described above, Quinn's value of k' is 2.21 x 10—2 M-1
sec—1. Oxalic and succinic acids were found to inhibit the reaction.
Malonic acid has no effect. Maleic, fumaric, malic, and tartaric acids
are all reported to increase the rate of the reaction under appropriate
conditions. With tartaric acid there appears to be strong catalysis,
but it was observed that the second-order kinetic plots Were very fre-
quently nonlinear. When the reported apparent second—order rate con-
stants are plotted as a function of initial uranium(IV) or tartaric acid
concentrations, much scatter is observed although possible trends are
indicated. In order to understand the nature of this system, it is

clear that more detailed research is necessary.
E. Oxidations of Tartaric Acid

The results of this present study indicate that oxidation reac-
tions of tartaric acid should be mentioned. d—Tartaric acid is sensi-
tive to many oxidizing agents. In general, when carefully oxidized,
g-tartaric acid yields dihydroxymaleic acid, tetrahydroxysuccinic acid,
and hydroxymalonic acid;151 stronger oxidizing agents decompose it into

various products, such as carbon dioxide and formic or oxalic acid.

 
  
  
  
   
 

I

   

t 1
Wea- at tartaric domains been semi-lbw .. :

iigdgyhof this subject is beyond the scope.ofrthis'thesdole~
" ”person: available, but "Beilstein"152 thoroughlyrecapitu-

I
3;
~11

o

lithe literature on this subject through 19h9. The following dis-
L iop will deal with some of the more important studies an oxidation

égogctions of tartaric acid, with emphasis on those studies in which the

   

rulf'
, i ' perennic intermediates have been characterized.

Much of the work on oxidations of tartaric acid is related to the
,fact that 1,2—glycols can be readily split by oxidation. Reagents such
as permanganate or acid dichromate react with 1,2-glycols to form two
moles of acid. By using a reagent, such as lead Udraacetate, periodic
acid, or sodium bimuthate, oxidative cleavage can be made to yield

aldehydes. However, tartaric acid does not behave strictly as a 1,2—

 

glycol.153

Strong oxidizing agents like permanganate153,154 and cerium(IV)155
react with tartaric acid to form formic acid and carbon dioxide. A
study of the oxidation of tartaric acid by selenious acid indirectly
showed that dihydroxymaleic acid is an intermediate oxidation product
in the tartaric acid oxidation by permanganate.156 Manganese dioxide
oxidation of tartaric acid is reported to yield acetaldehyde and carbon
dioxide.157 Oxidation of tartaric acid with dichromate in 3:5 mole

ratio in the presence of sulfuric acid occurs in two stages: a rapid

   
   
 

reaction to hydroxymalonic acid, and a slow reaction to dihydroxy-

nalonic acid.158 The kinetics of the reaction between tartaric acid

   

and peroxydisulfate has been studied and a mechanism proposed.159 .Ihg9*

i-R éthernal and photochemical reactions between tartrate and mangengfi'

  

31
Tartaric acid gives a silver mirror with an ammoniacal solution of a sil—
ver salt, the organic product being oxalic acid.161 Tartaric acid solu-

tions are photochemically oxidized in the presence of uranium salts to

primarily glyoxal, but also acetaldehyde, malic acid, succinic acid,

carbon monoxide, and carbon dioxide.162 A solution of tartaric acid is
oxidized by hydrogen peroxide in the presence of ferrous salts to di—
hydroxymaleic acid.163 In the electrolytic oxidation of tartaric acid,
as many as ten organic products are formed, which include dihydroxymaleic
acid, tetrahydroxysuccinic acid, and hydroxymalonic acid.164

in general, it appears that the first 2—electron oxidation of
tartaric acid yields dihydroxymaleic acid. This might decarboxylate
to yield hydroxypyruvic acid; but it appears that if more oxidizing
agent is available, the dihydroxymaleic acid is oxidized to tetrahydroxy—

succinic acid. Decarboxylation of the latter yields hydroxymalonic acid.

 

 

  

The great majority of work on reaction velocities is primarily
Rte-rned with trying to find out exactly in what manner the reactions
lé‘ proceeding. In other words, kinetic studies are generally aimed at
gzining insight into reaction mechanisms.

Standard kinetic techniques were employed in this.study, the

 

details of which are discussed in a great many textbooks}65 devoted
solely to the subject of kinetics. The only unusual techniques employed
will now be discussed.

It has been shown166 that under the conditions employed in the
present study (in the absence of organic acids), a graph of l/[U] vs.
time, t, should be linear at least over low values of t. This is
illustrated in Figure 3 for two experimental kinetic runs. Run A is
a rderence run performed under approximately Harkness' and Halpern's
conditions. A good linear second-order plot is obtained for this run,
and it is clear that the l/[U] 32. t graph is indeed linear. Thus it
is a simple matter to extrapolate the short distance to t = O and read
off l/[U]o, and hence obtain [U]o completely consistent with the other
values of [U]. Run B is a kinetic run performed under almost the same
conditions, except that tartaric acid is present. The second-order plot
at this tartaric acid concentration shows as much or more curvature than

any other kinetic run performed. The plot of l/[U] 3;. t for this run

as shown in Figure 3 is clearly nonlinear, however, it still is not dif-

 

 

 

33

 

 

 

 

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I .2 o-on I.§ mH u OHS/HHS .a use
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.oeHb Mo cpoucsm m mm coHumanoucou EdHcmns mo HmuoeaHuon och .m opSmHm
Am OH x .uomv oEHP
ca o.m as m H NH co 0 .
H _ _ _ Ln a

 

m5.

 

IV. EXPERIMENTAL
A. Preparation and Standardization of Reagents

various stock solutions were prepared in the course of these in—
vestigations. In allcases the starting materials were analytical re-

agents, except the thallium(I) nitrate which was C.P. grade. However,

'further purification was performed in the course of preparing the thal-

lium stock solutions, as was also the case with the uranium and sodium
perchlorate stock solutions.

The water used in all stock solutions was distilled demineralized,
1.3., distilled water which had been passed through a mixed bed of
cation and anion exchange resins, as supplied in a commercial unit
known as a Crystalab DEEMINIZER. This water contains less than 0.5
parts per million ionic impurities (measured as NaCl).

All glassware was kept scrupulously clean. Preliminary treatment
with aqua regia, alcoholic potassium hydroxide, or sodium dichromate—
sulfuric acid cleaning solution was employed when deemed necessary.
Ordinarily washing with a detergent was sufficient. The glassware was
thoroughly rinsed with hot tap water, then at least five times with
distilled water, followed by at least five rinses with demineralized
water, and finally set aside to dry.

Table II lists the important stock solutions with their determined
concentrations. A number of other stock solutions not noted in Table II
were prepared, however, these were either of only supplementary importance

or utilized only in preliminary experiments.

3h

 

35

Table II. Concentrations of stock solutions

 

 

 

Stock Solution Molarity

Hc104 b.872

NaOH (no. 1) 0.5385

NaOH (no. 2) 0.h922
uranium(VI) perchlorate:

[U(VI)] 0.1871

[U51v)] nil

[H J 1.022
uranium(IV) perchlorate:*

[U(IV)] 0.1263

[U(VI)] 0.0017

[H*] p 0.9275
potassium permanganate 0.1208
thallium(I) perchlorate:

[T1(I)] 0.09LS

[Tl(III)] nil

[H+] 1.002

thallium(III) perchlorate:

[Tl(III)] 0.0970

[T1(I)] 0.0001

[n+1 1.796
potassium bromate 0.1000
sodium perchlorate b.2b3
tartaric acid 0.1000

 

*The molarities cited here are for the freshly prepared uranium(IV)
stock solution. After almost ten months of storage, the [U(IV)] was
0.1079 M, the [U(VI)] was 0. 0201 M, and the [H+ ] was 0. 96b6M _

In many of the standardizations, micro pipets were used to sample
the fairly concentrated stock solutions. These pipets, which are cali-
brated to contain rather than to deliver, allow one to avoid errors

from too many volumetric operations and also to avoid viscosity effects

.d.‘ ‘

‘n'l'...

36
and drainage errors. The techniques involved in microvolumetric manipu-

lations have been described.167

1. Perchloric Acid

The perchloric acid stock solution was prepared by diluting Baker
Analyzed Reagent 70-72% perchloric acid with demineralized water. This
stock solution was used in the preparation of the kinetic run solutions
and either this solution or concentrated reagent perchloric acid was
used in the preparation of all the uranium and thallium stock solutions.
Both the perchloric acid stock solution and the concentrated reagent
gave negative chloride tests with silver nitrate.

The concentration of the perchloric acid stock solution was deter-
mined by titrating with standard sodium hydroxide solution to the methyl

red end point.

2. Sodium Hydroxide

Carbonate-free sodium hydroxide solutions approximately 0.5 M were
prepared from reagent grade sodium hydroxide pellets according to the
method described by Kolthoff and Sandell166 and stored in tightly stop-
pered polyethylene bottles. These solutions were used to determine the
hydrogen ion concentration in the perchloric acid stock solution and in
all the thallium and uranium stock solutions. Sodium hydroxide solution
was also used to check the molarity of the tartaric acid stock solution.

The sodium hydroxide solutions were standardized against primary
standard reagent potassium acid phthalate, which had been dried for two

hours at 125° C, using phenolphthalein as the indicator.

   

'gfidfiiiip::chlorate solution was prepared according to a pro;

 
 
  
 

*Vioiflapted from the work of Quinn,36 who used a method modified

  

' A that‘round in "Gmelin."159
:urnfilfiaker Analyzed Reagent uranium(VI) nitrate hexahydrate was dis-
bdlved in demineraliZed water and filtered through a fine fritted fun-
nel to remove the considerable cloudiness present. The filtrate was
diluted with demineralized water to make a 0.3 M solution. Reagent
grade 30% hydrogen peroxide solution was diluted with demineralized
water to give a 3% solution. The uranyl nitrate solution was heated to
about 80° C and the hydrogen peroxide solution, in a 2:1 mole ratio of
hydrogen peroxide to uranyl nitrate, was added slowly with stirring. A
pale yellow precipitate of uranium peroxide, U04'xH20 (x = 2 and/or h),17°
formed and it was digested at 80° C for several hours. The precipitate
was then filtered and washed ten times with demineralized water, the
precipitate being collected in a medium fritted funnel. The uranium
peroxide was dissolved at 80° C in enough 1.5 M perchloric acid to make
a solution 0.3 M in uranyl ion. Under these conditions dissolution took
about one day.
After the precipitate had completely dissolved, the solution was
again filtered through a fine fritted funnel and treated with 3% hydro-
gen peroxide at 80° C. This time twice as much peroxide was used and

it was observed that the rate of precipitation was much slower under

  
     

this more acid condition. Again the precipitate was digested, filtered,

      
      

washed, and dissolved in 1.5 g perchloric acid. Then the whole pry»
I.

‘ 3 and repeated for a third time. Finally the uranium(VI) poi-chm...“ " ' ‘

 
   
    
  
 

38

'an but most of it was used in the preparation of the uranium(IV)

,.gflalorate solution. This solution was quantitatively analyzed for

{f u;ppium(VI) and hydrogen ion; uranium(IV) was shown to be absent. Also
apnegative chloride test was obtained with silver nitrate solution.

The concentration of uranium(VI) was determined by reducing ali-

quots in a Jones reductor and titrating with potassium permanganate.
The Jones reductor was prepared according to directions given by Kolthoff
and Belcher.171 For use in the reductor, zinc amalgamated with 2% of
mercury was made from 20 mesh reagent grade zinc and reagent grade
mercuric chloride. The reduction was performed according to the pro-
cedure described by Kolthoff and Belcher.171 Since uranium is partially
reduced to uranium(III) under these conditions, air was bubbled through
the reduced solution for five minutes at room temperature to oxidize
uranium(III) to uranium(IV).172 Quadrivalent uranium is fairly stable
under these conditions and was titrated with standardized potassium
permanganate to the appearance of a pink tint. The titration reading
was corrected for the volume of permanganate required by the reagents
and reductor, as well as for the volume required to obtain an end point
color in the yellow solution. These corrections were made according to

the procedures given by Lundell and Knowles.”2

uranium(IV) was shown to be absent by performing a.permanganate

titration on a solution identical to the ones employed above, except

     
   

that the uranium aliquot was not passed through the Jones reductor.

The volume of permanganate was the same, within experimental error,}‘,,y"

  
  

; thregen infi‘concentration was determined'by replacing

N

“vii ions with hydrogen ions from an ion exchange resin and\ti-

Jay the resulting solution with standard sodium hydroxide to the

  
  
  

“Viiphthalein end point. The ion exchange was made in a 10 cm. col-
‘tiuiprepared from 100- 200 mesh DOWEX SON—X12, which is a hydrogen ion
' ‘ ‘ farm, strongly acidic, cation exchange resin. The free hydrogen ion

concentration is then equal to the titrated hydrogen ion concentration

minus twice the uranium(VI) concentration, all concentrations being ex-

pressed in units of molarity.

h. Uranium(IV) Perchlorate

Uranium(IV) perchlorate stock solution was prepared by electrolytic
reduction of the uranium(VI) perchlorate solution. The method employed
was similar to that used by Arhland and Larsson,"3 Quinn,38 and Gordon
and Taube.17

Figure h is a diagram of the apparatus. Besides other components,
the apparatus consisted of two flasks: an electrolysis flask and a
storage flask. Associated with the electrolysis flask was equipment
used in the reduction, namely, the solution to be reduced, the anode,
the cathode, the power supply, and a cooling bath. Along with the stor-
age flask, there were tubing and stopcocks used for transferring the
reduced solution to the storage flask, for removing uranium(IV) samples,
and for providing a nitrogen cover to prevent atmospheric oxidation of

uranium(IV). Also a means was provided for air—oxidizing any uranium(III)

   
 

igfermed during the electrolysis.
I The electrolysis flask was a modified 1000 m1. one- -neck round

‘%

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Uranium(IV) apparatus.

Figure A.

 

b1
the bottom of the flask permitted electrical connection to the mercury
cathode. Attached to the side of the flask was an anode compartment
that consisted of a fifteen millimeter "L" shaped tube containing a fine
fritted disc three centimeters from the flask proper. The details of
the fitting atop the flask are shown in Figure h, where it can be seen
that the center tubing extended to just above the surface of the mercury
pool.

After the mercury had been placed in the electrolysis flask, the
solution to be reduced was added. This solution was uranium(VI) per-
chlorate stock solution that had been diluted to 860 ml. with perchloric
acid and demineralized water, such that its uranium(VI) concentration
was 0.13 M or 0.26 E and its hydrogen ion concentration was 1.2 fl.
Electrolytic reduction of this solution, under the conditions described
later, proceeded without complication.

However, in an electrolysis attempt where the uranium(VI) concen-
tration was initially O.lh M and the hydrogen ion concentration was
only 0.77 M, a black precipitate was observed at the surface of the
mercury pool after about thirty minutes of electrolysis. Such behavior
has been noted by various workers.38:174,175 This material was almost
certainly a hydrous oxide of uranium. Casto,”6 who discusses the
electrolysis of uranium solutions, says the reduction of an aqueous
solution of uranyl salt deposits uranium on the cathode as a hydrous
oxide. He also reports that Heall77 has stated that the properties of
the deposit vary according to the conditions. If the solution is nearly
neutral (with a pH of about 6), a fluffy yellow oxide having the composi-
tion of a hydrated U03 forms rapidly at the cathode, and if the solution

is somewhat more acid, with a pH near 3, the deposits may be green or

b2
black. These deposits will contain more or less of the lower oxides,
approximating the composition U308 and U02. Casto176 further reports
that Pierle,”8 in his extensive investigation of the electrochemistry
of uranium, has concluded that uranium is deposited as an oxide of
variable composition. In this present work, no effort was made to
isolate or identify the black precipitate. It seems likely that despite
the over-all hydrogen ion concentration of about 0.77 M, the hydrogen
ion concentration in the vicinity of the cathode was decreased suffi-

ciently by the cathode reaction:

 

U022+ + 1m" + 2e- > 114+ + 2H20

so that the hydrous uranium oxide could precipitate. The dissolution

of the black precipitate was observed to be slow (of the order of a day)
even after uniform acid conditions had been obtained. It could, however,
be dissolved more rapidly by adding additional acid.

The anode compartment of the elctrolysis flask was filled to the
level of solution in the flask proper with perchloric acid of the same
concentration as in the flask, namely 1.20 M. Dipping into the perchloric
acid was a platinum electrode made by "hammer-sealing" a piece of 20
gauge platinuni wire to a one centimeter square of extra heavy platinum

foil. The major anode reaction occurring during electrolysis is:

 

H20 > 2H+ + 1/202 + 2e".

Reagent grade mercury was cleaned and used as the cathode. The
mercury was initially cleaned by passing it three times through a pinhole
in the vertex of dry filter paper cones placed in a dry funnel. Then
the mercury was washed five times with dilute nitric acid and five times

with demineralized water according to the procedure described in the

b3
"Handbook of Chemistry and Physics."179 The mercury, which was only
blotted to remove excess water, was poured into the electrolysis flask
until it formed a pool three inches in diameter. To provide an electrode
terminal, a piece of 2h gauge platinumx wire was placed into the mercury
side arm. In addition to the major cathode reaction, which was given
above, the reaction:

U4 + e- > U3

 

also occurs.

The direct current power supply was a "Heathkit" Battery Eliminator
Model BE-b from the Heath Company. Two volt—ohm-milliammeters were ar-
ranged to read voltage and amperage as shown in Figure A. The electrol-
ysis was conducted at ll v. and 0.7 amp.

The electrolysis flask was submersed in an ice—salt bath in order
to keep the solution at 00 C. The bath was circulated with a magnetic
stirrer. Gordon and Taube”:82 report that the reduction of uranium(VI)
to uranium(IV) goes to completion if the solution is kept cold; other-
wise a few per cent uranium(VI) is left.

Prior to electrolysis, the system, including the uranium(VI) solu-
tion, was swept with nitrogen for 2h hours. For this operation, the
stopcocks were set as shown in Figure h and the nitrogen flow rate from
the tube in the elctrolysis flask was two large bubbles per second. One
hour before starting the electrolysis, the ice-salt bath was applied.
This bath was replenished three times during the electrolysis. The
reduction was performed with mixing accomplished by nitrogen bubbling
at the rate mentioned above. The starting time, voltage, amperage, and
bath temperature were noted. The electrolysis took about eight and one-

half hours as predicted by Faraday's Law for the number of equivalent

u
weights of uranium present and the current employed. The end of the
reaction was indicated by the incipient formation of uranium(III), whose
deep red color when mixed with the dark green color of uranium(IV)
rendered the solution almost black. The red of uranium(III) is clearly
visible when the solution is viewed with a light behind it.

The uranium(III) was removed by air oxidation. This was accomplished
by stopping the nitrogen flow, closing stopcocks A, B and D (see Figure
h), attaching an "air cleaner" composed of a series of drying tubes
containing Ascarite and glass wool (in order to remove carbon dioxide
and dust) at the ball—and-socket joint C, and then opening stopcock B
and turning stopcock A to the vacuum. Incidently, flexibility in the
apparatus at ball-and—socket joint E, between stopcocks H and I, and
between stopcock J and the electrolysis flask, permitted easy opening
of ball-and—socket joint C for the air cleaner connection. Air was
allowed to bubble slowly through the solution for three minutes. Step-
cock A was returned to its original position, stopcock B was closed,
the air cleaner was disconnected and ball—and-socket joint C reconnected,
stopcock D was turned to connect only joints C and E, and the nitrogen
flow was started. Nitrogen was allowed to flush the system for two
hours at a fast rate in order to remove any oxygen.

Next the pure uranium(IV) solution was transferred from the elctrol-
ysis flask to the storage flask. First a rubber stopper was fitted
tightly into the anode side arm. The nitrogen was almost stepped, stop-
cock A was closed, stopcock G was opened, and stopcock J was turned to
direct incoming nitrogen to the electrolysis flask. The nitrogen pres-
sure was increased and this forced the uranium(IV) solution up the tube,

through B, C, D and E, down through the coarse fritted disc F and into

us
the storage flask.

The uranium(IV) stock solution was stored under a nitrogen atmos-
phere by closing stopcocks D, G and H. Incoming nitrogen was vented at
stopcock A.

To remove uranium(IV) solution, nitrogen is directed along the
route J-I-D-E-F, with stopcock G closed and stopcock H open to the
sample exit. The tube for sample removal is then emptied by opening
stopcock G and directing nitrogen along the route J-I-H-G. In normal
sampling, the tube for sample removal and the sample exit are rinsed
twice in this manner before collecting the sample of uranium(IV) stock
solution.

The uranium(IV) perchlorate stock solution was analyzed for uran-
ium(IV), uranium(VI), and hydrogen ion.

The concentration of uranium(IV) was determined exactly as it was
determined in the uranium(VI) perchlorate stock solution. Here, of
course, the uranium(IV) constituted almost all the uranium, whereas
before its concentration was nil.

The concentration of uranium(VI) was taken to be the difference
between the total uranium concentration and the uranium(IV) concentra—
tion. The total uranium concentration was determined exactly as the
uranium(VI) concentration (3.3., total uranium concentration) was deter—
mined in the uranium(VI) perchlorate stock solution.

The free hydrogen ion concentration was determined in the same man-
ner as it was determined in the uranium(VI) perchlorate stock solution.
However, since both uranium(IV) and uranium(VI) are present, the free
hydrogen ion concentration was taken as the titrated hydrogen ion molar-
ity minus four times the uranium(IV) molarity and minus twice the uran-

ium(VI) molarity.

b6
5. Potassium Permanganate
A 0.1 N_potassium.permanganate solution was prepared from reagent
grade potassium permanganate crystals according to the procedure given
by Kolthoff and Belcher.18° This solution was standardized against
primary standard reagent grade sodium oxalate, using the procedure of

Fowler and Bright181 as described by Kolthoff and Belcher.152 'This

potassium permanganate stock solution was used in the various uranium

determinations.

6. Nitrogen

Specially purified nitrogen was used to protect the uranium(IV)
perchlorate solution from atmosphereic oxidation, to effect the trans-
fer of this solution from the electrolysis flask to the storage flask,
and to remove uranium(IV) samples from the storage flask. The nitrogen
purification train, which was adapted from the works of Gordon32 and
Quinn,35 is pictured in Figure 5.

In order to have oxygen-free gas, prepurified nitrogen was passed
over COpper gauze heated to L650 C in a tube furnace and then down a
heated column containing active copper deposited on kieselguhr. The
temperature in the tube furnace was controlled with a variable auto-
transformer and measured with a chromel-alumel thermocouple and potenti-
ometer, using standard calibration tables.183 The kieselguhr column
was electrically heated and glass insulated as shown in Figure 5, the
degree of heat being controlled with another variable autotransformer.
The temperature in the center of the column was 2050 C, as read on the
installed thermometer. The copper on kieselguhr was prepared by the

method of Meyer and Ronge,184 as described by Dodd and Robinson.185

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After considerable usage the copper gauze and copper on kieselguhr be-
came oxidized and had to be regenerated. This was accomplished by
substituting a hydrogen cylinder for the nitrogen cylinder and allow—
ing hydrogen to pass over and reduce the heated copper oxide. The
unused hydrogen and the water produced were removed through the stopcocks
located at the end of the furnace tube and at the bottom of the kiesel-
guhr column.

Further down the train, the nitrogen.passed through three gas wash-
ing bottles equipped with fritted discs.4 The first bottle contained
chromium(II) sulfate solution over amalgamated zinc. This is a very .
efficient means of removing traces of oxygen186 and would be effective
should an electrical failure render the other oxygen purgers ineffec-
tive. The method of Stone and Beeson,157 which is described by Dodd
and Robinson,188 was used for the preparation of the chromous sulfate
solution and the amalgamated zinc. The second bottle contained deminer—
alized water which served to wash the nitrogen. In order to minimize
volume changes and consequent concentration changes in the uranium(IV)
stock solution during storage and sampling, it was necessary to have the
vapor pressure of water in the incoming nitrogen approximately equal to
the water vapor pressure of the uranium(IV) stock solution. To accom—
plish this, the third gas washing bottle contained a perchloric acid
solution whose concentration was approximately equal to the sum of the
total uranium and free perchloric acid molarities of the stock solution.
After the nitrogen left this bottle it was directed into the apparatus
for uranium(IV) preparation and storage.

Even when the nitrogen was not being used in the uranium(IV) appar-

atus it was allowed to flow at a very slow rate. This prevented any of

b9
the gas washing bottles from backing up and also slowed the rate at which
the fritted disc in the chromium(II) sulfate bottle would become clogged

with deposits.

7. Thallium(I) Perchlorate

Thallium(I) perchlorate stock solution was prepared by dissolving
thallium(I) nitrate in an excess of hot concentrated perchloric acid,
fuming to volatilize nitric acid, and then recrystallizing the precipi-
tated thallous perchlorate three times from perchloric acid. The start-
ing material was purified thallium(I) nitrate from the Fisher Scientific
Company. The expulsion of nitric acid was accomplished by fuming in
the temperature range 125 to 11100 C and nitrate was shown to be absent
from the precipitated thallous perchlorate by a negative brown ring test.189
The thallous perchlorate was slowly recrystallized three times from
dilute perchloric acid and then some of this salt was dissolved in de-
mineralized water. Perchloric acid was added to give a solution about
0.1 M in thallium(I) and 1.0 M in acid. This solution gave a negative
chloride test with silver nitrate solution.

The thallium(I) perchlorate stock solution was analyzed for thal-
lium(I) and hydrogen ion; thallium(III) was shown to be absent. The
concentration of thallium(I) was determined by potentiometric titration
with standard potassium bromate in the presence of 5-8% hydrochloric
acid at 10-700 C. This method was adapted from the works of Kolthoff190
and of Zintl and Rienacker.”l To determine the end point a Beckman
Laboratory Model G pH Meter was employed with Beckman electrodes of
saturated calomel and platinuni. The end point is very sharp. Hydro-

chloric acid is added because apparently chloride ion catalyzes the

50
otherwise slow oxidation-reduction reaction. The elevated temperature
enhances the speed of the reaction, keeps the saturated calomel electrode
from becoming clOgged with solid potassium perchlorate when the perchlor—
ate ion concentration is high, and prevents thallous chloride from pre-
cipitating. The thallium aliquot should be added to a preheated aqueous
hydrochloric acid solution to avoid the precipitation of thallous chlor-
ide, which dissolves only slowly on heating.

The hydrogen ion concentration in the thallium(I) perchlorate stock
solution was determined by titrating with standard sodium hydroxide to
the methyl red end point. 1

Determination of total thallium in the thallium(I) solution, by the
method employed with the thallium(III) solution, gave the same value

as was found for the thallium(I) concentration, thus proving the absence

of thallium(III).

8. Thallium(III) Perchlorate

Thallium(III) perchlorate stock solution was made from the pre—
viously prepared thallous perchlorate crystals by a method only slightly
modified from that described by Biedermann.1°° The procedure consists
of the anodic oxidation of a perchloric acid solution of thallous per-
chlorate. This method has the advantages that there is no need to intro-
duce foreign substances, the concentration of thallium(III) ions can
be varied within wide limits, and a very simple apparatus is sufficient.
The apparatus is shown in Figure 6. An 800 m1. beaker served as the
anode compartment and contained the thallous perchlorate solution with
2.0 M perchloric acid to avoid hydrolysis. The beaker was covered with

a stopper that could support a platinun1 anode and a cathode compartment,

S1

 

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Figure 6. Thallium(III) apparatus.

S2

besides having a hole for gas exit. The anode was 52 mesh platinum
gauze in the form of a cylinder ho mm. in diameter and 35 mm. in height.
The cathode compartment was made from a 20 mm. straight sealing tube
containing a medium fritted disc by cutting off the excess tubing just
below the disc. This compartment was filled with 2.0 M perchloric
acid and a piece of 2b gauge platinum. wire in the form of a coil was
supported in it by a stopper provided with a hole for gas escape. The
power supply and meters were the same as described in the preparation
of uranium(IV). The thallium solution was stirred efficiently with a
teflon-coated bar driven by a magnetic stirrer. The electrolysis was
initially conducted at ll\h and 2.6 amp. However, under these condi-
tions oxygen evolution from the anode commenced when there was still
about 13% thallium(I) left. In order to get more complete oxidation
the current was progressively reduced so as to avoid oxygen evolution
and continue thallium oxidation. The end of the electrolysis was con-
ducted at 2.h v. and 0.1 amp. Under these conditions very slow oxygen
evolution from the anode indicated the completion of the thallium
oxidation. The yield was found to exceed 99.9%. The total length of
time necessary for the electrolysis could be reasonably predicted from
Faraday's Law by knowing the number of equivalent weights of thallium
present and the current used for each portion of the electrolysis.
During the course of the electrolysis, considerable hydrogen gas was
liberated at the cathode and additional 2'0.M perchloric acid had to
be added several times.

The electrolyzed solution was passed through a fine fritted funnel
and gave a negative chloride test. This thallium(III) stock solution

was stored in an amber bottle and protected from sunlight. Biedermannm0

53
reports essentially no increase in thallium(I) under these conditions
even after four weeks, whereas thallium(III) solutions exposed to diffuse
sunlight are reduced with considerable speed. The stock solution was
quantitatively analyzed for thallium(I), thallium(III), and hydrogen
ion.

The concentration of thallium(I) in the thallium(III) perchlorate
stock solution was determined in the same manner as was thallium(I) in
the thallium(I) perchlorate solution.

The concentration of thallium(III) in the thallium(III) perchlorate
stock solution was taken as the difference between the total thallium
concentration and the thallium(I) concentration. Total thallium con-
centration was determined by reducing thallium stock solution aliquots
with sulfur dioxide, boiling to remove excess dissolved sulfur dioxide,
and then titrating the thallium(I) potentiometrically with bromate as
described above. The reduction was accomplished by bubbling C.P. sulfur
dioxide through the solution at a moderate rate for 20 minutes. The
sulfur dioxide expulsion was effected by boiling the solution in the
presence of three carborundum boiling chips for at least one hour, though
acidic vapors could not be detected with wet pHydrion paper after about
25 minutes.

The hydrogen ion concentration in the thallium(III) perchlorate

stock solution was determined as in the thallium(I) solution, except that

prior to titrating, the thallic ion was complexes with about a hundred—

fold excess of bromide.

9. Potassium Bromate

Reagent grade potassium bromate was dried for 12 hours at 165° C.

Sh
To prepare a 0.1000 N solution, 2.7835 g. was weighed, dissolved in
demineralized water, and diluted to one liter. This stock solution

was used as titrant in the potentiometric determinations of thallium.

10. Sodium Perchlorate

Sodium perchlorate stock solution was prepared by dissolving twice
recrystallized sodium perchlorate in demineralized water. To start
with, reagent grade hydrated sodium perchlorate from 0. Frederick Smith
Chemical Company was dissolved in demineralized water and the solution
was filtered through a fine fritted funnel. The filtrate was placed
on a hot plate and boiled gently (b.p. approximately lh2° C) until
solid began to appear on the surface. A small amount of demineralized
water was added to dissolve the solid and the solution was placed in an
oven preheated to 60° C where it was allowed to cool undisturbed for
approximately 2h hours. Mellor192 reports that A. Potilitzin showed
that under these conditions, 3.3., crystallization of an aqueous sodium
perchlorate solution above 50° C, nondeliquesent prismatic crystals of
anhydrous sodium perchlorate, NaC104, are formed; whereas below 500 C,
deliquesent monohydrated sodium perchlorate, NaC104-H20, is formed.
If a hot solution of sodium perchlorate was allowed to cool undisturbed
to room temperature, crystals of anhydrous sodium perchlorate slowly
separated; but the solution was clearly supersaturated with respect to
monohydrated sodium perchlorate because when an attempt was made to de-
cant the supernatant liquid, crystals of the hydrate suddenly appeared.
Such rapid recrystallization was deemed undesirable because of the like-
lihood of occluded impurities; hence the method of recrystallization

above 50° C was employed. The crystals were collected in a coarse fritted

SS
funnel but not washed because even a quick rinse with a minimum amount
of ice cold water rapidly dissolves the crystals. No solvent other
than water was used in order to avoid contamination. Cold aqueous per-
chloric acid could perhaps have been used to wash the crystals, but
then determination and corrections for the amount of acid left on the
crystals would have to be made and the simple method of standardization
described below could not be employed. In order to increase the yield
which is low because of the high solubility of sodium perchlorate in
water, the recrystallization procedure was applied to the supernatant
liquids three more times and the four crops of crystals were combined
to produce an 8h% yield of once recrystallized sodium perchlorate. The
last supernatant liquid was discarded. Since the recrystallization
method does not involve washing the crystals, the whole procedure was
repeated again. The then twice recrystallized sodium perchlorate was
dissolved in demineralized water and the resulting stock solution was
seen to give a negative silver nitrate test for chloride.

The concentration of the sodium perchlorate stock solution was
determined by evaporating aliquots, drying at 165° C and weighing the
anhydrous salt. One milliliter aliquots were measured with 1000 lambda
micro pipets and delivered with three rinses into weighed porcelain
crucibles with lids. The crucibles were placed in the oven with their
lids slightly askew and allowed to dry at least 2h hours, a length of
time sufficient to produce constant weight. The crucibles were then
covered, removed from the oven, placed in a desiccator, allowed to cool
30 minutes and finally weighed. The anhydrous sodium perchlorate was
observed to be slightly hydroscopic, but this introduced no difficulry
in weighing because the salt would only absorb water very slowly in the

covered crucibles.

56

ll. Tartaric Acid

To prepare a 0.1000 M tartaric acid stock solution, 3.7523 g. of
Mallinckrodt granular Analytical Reagent d-tartaric acid (m.p. 169-1700 C)
was weighed, dissolved in boiled demineralized water in a sterilized
flask, and diluted to 250 ml. Titration of this solution with standard
sodium hydroxide to the phenolphthalein end point gave excellent agree-
ment (within 0.1%) with the expected value for the molarity.

Quinn38 has observed the appearance of growths, that appeared to
be fungi, in tartaric acid solutions on standing. In this present work,
such growths were observed in unsterilized solutions in as few as nine
days. By taking the simple precaution of using boiled water and a

sterile flask, no growths were observed in as long as one year.
B. Measurement of Uranium(IV) Concentrations

Despite elaborate precautions taken to exclude oxygen and light
from carefully purified uranium(IV) perchlorate stock solutions, it has
been observed that the uranium(IV) concentration still gradually changed
over the course of days. Quinn's35 change amounted to about 0.013 M in
120 days; the change in this present study was 0.0184 M in 295 days.
This means that in order to know accurately the uranium(IV) concentra-
tion at ahy time, it would be necessary to analyze the stock solution
periodically. In addition to this drawback, the uranium(IV) is oxidized
in the interval between sample removal from the storage flask and initia-
tion of the reaction. This is probably due to exposure to air. Thus
the concentration of uranium(IV) at the beginning of the reaction could

not be satisfactorily determined from the concentration of the uranium(IV)

57
stock solution; and consequently another method of measuring the initial
uranium(IV) concentration had to be employed.

As was indicated in Chapter III, a plot of 1/[U] vs. time under
appropriate conditions yields a straight line, at least early in the
reaction, for a reaction with the second order rate law:

—d[U(IV)]/dt = k[U(IV)][Tl(III)].
The initial uranium(IV) concentration can be determined from this type
of graph (see Figure 3) by extrapolating to zero time and reading the
value of the reciprocal uranium concentration. Many of the reactions
in this work yielded good straight lines for this type of plot; others
had some curvature, but still a reliable value of initial uranium(IV)
concentration could be obtained by extrapolating the slightly curving
line the small distance necessary. Thus the initial uranium concentra-
tion could be determined quite accurately from the reaction data and the
value obtained was, of course, consistent with these data.

The course of the reaction was followed by measuring the changing
uranium(IV) concentration. This concentration was determined by measur—
ing the absorbancy as a function of time. Readings were made with a
Beckman Model DU Quartz Spectrophotometer (with cell compartment thermo-
stated to 25.0 i 0.10 C) using the 650 mu absorption maximum, which is
the most intense visible uranium(IV) peak.55:193 At this wave length

the interference from other ions was found to be nil, as has been pre-

+ +

viously observed.147 These other ions were U022+ Tl3 , T10H2+, Tl ,

)

H+, Na+, and 0104-. Tartaric acid is also colorless at this wavelength.

Silica absorption cells of 1.00 cm. thickness were used, with demineral-
ized water in the reference cell. The need for making cell corrections

due to absorbance by the cells themselves and for varying thicknesses

58
of the cells employed was examined. At 650 mu and in the range of
absorbancies encountered, cell corrections were never more (and gen-
erally much less) than 0.h% of the absorbancy reading. This is easily
within half the per cent relative error reported for the Beckman DU
spectrophotometer under these conditions.194 Hence cell corrections
were not made.

In order to obtain a relationship between the uranium(IV) con-
centration and the absorbancy, a series of solutions varying in uran-
ium(IV) concentration but identical in hydrogen ion concentration and
ionic strength were prepared. The uranium(IV) was obtained from very
recently standardized uranium(IV) stock solution and exposure to air
was kept at a minimum. These solutions were examined spectrophotometri-
cally under the same conditions as were the kinetic runs. Table III
' lists the concentrations of these standard solutions and the correspond—
ing absorbancy values. An excellent straight line plot of these data
(as shown in Figure 7) confirms Kraus' and Nelson's statement55 that
the absorption band obeys Beer's Law. The data were treated by the
method of least squares and the following equation obtained:

U = 17.39 A + 0.00283,
where U is the uranium(IV) molarity times 103 and A is the absorbancy
at 650 mu. This equation was used to calculate concentrations corres-
ponding to measured absorbancy readings.

This method of following the uranium(IV) concentration is still
completely valid in the kinetic runs performed at different hydrogen ion
concentrations. Kraus and Nelson55 report that the variation in molar
absorptivities of uranium(IV) solutions is negligible in the high

acidity range employed in this study.

59

Table III. uranium(IV) Standard Solutions. [n+1 = 1.76 .u; u = 2.9;
Temp. = 25.00 C; k = 650 mu.

 

 

 

[U(IV)] (M x 103) Absorbancy
0.00 0.000
0.13 .008
1.26 .072
2.53 .le
3-79 .217
5.05 .291
6.31 .361
7.58 .h37

 

C. The Uranium(IV)-Tha11ium(III) Reaction

All kinetic runs were conducted in essentially the same manner.
This consisted of mixing equal aliquots of two solutions, one contain—
ing thallium and the other containing uranium, but otherwise identical
in all respects. The reacting solution was transferred to an absorption
cell, placed in a Beckman DU spectrophotometer, and the concentration
of uranium(IV) was followed spectrophotometrically as a function of time,
The two solutions for each run were as identical as possible so as
to minimize dilution effects, heat of mixing, Slow hydrolyses, slow
complex formation, etc. In order to make comparisons possible, the
conditions chosen were similar, for the most part, to those of Harkness
and Halpern147 and Quinnse. In all runs except those of the tartaric
acid variation, the initial tartaric acid concentration was 6.00 x 10-3
M. Thus both solutions for each run were prepared so as to have that

tartaric acid concentration. In those runs excluding the uranium(IV)

l

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.NN xocmonOmnd .m ouumwm

 

.mcnmccmpm asacmns now coapwnpcoocoo A>Hvasmcmns

$2 x ed 353
p m a m
_ A .14 _

 

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61

variation, the initial uranium(IV) concentration was approximately

3.50 x 10-3.E° In those runs excluding the thallium(III) variation,
the initial thallium(III) concentration was 9.00 x 10’3'm. Since both
uranium(IV) and thallium(III) were diluted two-fold upon mixing equal
aliquots, their concentrations in their respective solutions were twice
the values cited above. In all runs except those of the hydrogen ion
variation, the hydrogen ion concentration was 1.76 M. Thus both solu-
tions for each run were made 1.76 M by adding perchloric acid stock
solution and recognizing that the uranium(IV) and thallium(III) stock
solutions contributed some to the hydrogen ion concentration. All
solutions were adjusted to an ionic strength of 2.9 M by addition of
sodium perchlorate stock solution. In 1.76 M perchloric acid, uranium(IV)
is hydrolyzed approximately 1.2% and thallium(III) is hydrolyzed approxi-
mately b.0%. These small extents of hydrolysis applied to the low
uranium and thallium concentrations means that hydrolysis has a neglig-
ible effect on the ionic strength. Since tartaric acid is only 0.063%
ionized in 1.76 M acid, the low concentration of tartaric acid added
certainly produces only a negligible effect on the ionic strength. Thus
the ionic strength, u, of the thallium and uranium solutions were cal-

culated using the following relationships:

u %H[TH+] + [Na+] + 9[Tl(III)] + [Tl(I)] + [C104_]}

and

H

A ”n+1 + [Nan + 161mm] + Mun/1)] + [C104']} .

The perchlorate ion concentration was taken as:

[c104‘] = [H*] + [Na+] + 3[Tl(III)] + [11(1)]

62
and _
[clog] = [n+1 + [Nah + Mum/>1 + 21mm].

The two solutions for each run were prepared in 10 ml. volumetric
flasks. The reagents were added in the following order: perchloric
acid, sodium perchlorate, thallium(III) or uranium(IV), tartaric acid,
and lastly demineralized water to the mark. The additions were made
with initially dry graduated pipets, care being taken to wipe the out-
side of the pipets with a piece of tissue, to add the reagents without
wetting the ground glass neck of the flask, and to add the reagents
slowly so as to minimize drainage errors. During all solution prepara-
tion, room temperature proved to be within 2 C0 of 25° C, so no temper-
ature equilibrations were performed prior to diluting to the mark with
demineralized water. To perform the dilutions, water was added with
a squeeze bottle to a point below the mark, the flasks were swirled to
mix the reagents but not wet the stopper, and then more water was added
with a medicine dropper, Splitting drops if necessary, to bring the
solution volume to the mark. The flasks were then inverted at least
ten times to thoroughly mix the reagents and subsequently allowed to
equilibrate for approximately 75 minutes in the same constant tempera-
ture bath that equilibrated the cell compartment of the spectrophotometer.
During these 75 minutes, the cell holder containing a reference cell
filled with demineralized water and a covered empty cell for each run
were placed in the cell compartment and allowed to equilibrate.

All the kinetic runs were started about two hours after the thal-
lium and uranium were added to their respective solutions. Prior to
starting the reaction, a h ml. aliquot of the thallium solution was

delivered with a volumetric pipet into a 25 m1. Erlenmeyer flask,

63

allowing extra seconds to insure the pipet had thoroughly drained. The
empty Erlenmeyer flask had been allowed to equilibrate:fix*75 minutes in
the constant temperature bath before the thallium solution was added
and then it was returned to the bath for the few minutes until the uran-
ium solution was added. To start the reaction a h ml. aliquot of the
uranium solution was delivered with a class A volumetric pipet into the
thallium-containing Erlenmeyer flask, allowing five seconds in excess
of the specified 19 seconds delivery time. Class A pipets were used
because in addition to their acknowledged accuracy, they provided a
reproducible drainage rate. Timing of the reaction commenced when the
pipet containing the uranium solution was half drained. Time was moni-
tered with electric timers which read direct to tenths of seconds and
totalize to 10,000 seconds. The flask was swirled throughout the uran-
ium addition and swirling was continued for the first 30 seconds of
timing, except for the seconds when the pipet was being removed. Then
the reacting solution was carefully poured into an absorption cell and
placed in the SpectrOphotometer. The concentration of uranium(IV) was
then followed by measuring the absorbancy as a function of time.

Kinetic runs were generally performed in sets of three. A time
schedule for starting the reactions and taking the absorbancy measure-
ments was devised. In planning the schedule a number of factors were
considered: (1) runs were to be started as close together in time as
possible, (2) the runs were to be followed closely for one hour, (3)
thirteen absorbancy measurements were to be taken for each run so as
to avoid misinterpreting results should one or two points be in error,
(b) absorbancy readings for a particular run should be taken at inter-

vals of 100 seconds early in the reaction and at larger intervals later

6b

in the reaction, (5) for esthetic reasons and to aid comparisons, ab-
sorbancy readings in all runs should be taken after the same lengths
of time, (6) a minimum of 100 seconds would be necessary to prepare for
and start each run, and (7) a minimum of 50 seconds would be necessary
to standardize the spectrophotometer, take each reading, and record the
result. It was found that if the runs were started at intervals of
850 seconds and readings were taken at 100, 200, 300, 500, 700, 1000,
1300, 1600, 1950, 2300, 2650, 3100, and 3600 seconds, all the above
factors were satisfied. The first hour of the reactions was monitored
in this manner, but results from early runs indicated that it would
be profitable to follow the reactions for a longer length of time.
Consequently further readings were taken at 5h00, 7200, 16200, 18000,
and 21600 seconds. Thus most of the reactions were observed for Six
hours.

After concluding each set of runs, the cells were rinsed with
water and stored in 3 M hydrochloric acid, which removed any thallic
oxide deposited on the cells during rinsing. All thallium and/or

uranium solutions were collected for later reclaimation.
D. Studies of Uranium Solutions Without Thallium

The early kinetic results indicated that it would be advantageous
to examine uranium(IV) solutions containing tartaric acid, but no

thallium. This was done spectrophotometrically and polarographically.

l. Spectrophotometric Studies of Uranium
Uranium(IV) solutions containing varying amounts of tartaric acid

were examined Spectrophotometrically under the standard kinetic run

65
conditions previously described. Scans of the visible portion of the
spectrum were made on both Beckman DK—2 and DB recording spectrophotom—
eters. Precise examinations over the 650 mu peak (from 6&0 to 660 mu)
were performed on such solutions with the Beckman DU spectrophotometer.
These solutions were examined freshly prepared and after standing one

day.

2. Polarographic Studies of Uranium

Uranium(IV)—tartaric acid polarographic investigations were per—
formed in order to estimate the extent of complexation under the kinetic
run conditions. Polarograms were recorded at warm room temperatures
(27 t 2° C) with a Sargent Recording Polarograph, Model XXI. An ”H“
cell equipped with fritted disc below the sample solution arm (to
enable rapid oxygen outgassing) was employed. Two minutes of efficient
Sweeping with water-pumped nitrogen removed dissolved oxygen. The
sample solution arm contained uranium(IV) solutions about 3.5 x 10-3
M, with and without tartaric acid, 1.76 M in hydrogen ion, and having
an ionic strength adjusted to 2.9 with sodium perchlorate. No maxi-
mum suppressor was needed in the uranium polarography. The other arm
of the "H" cell contained perchloric acid—sodium perchlorate solution
of the same acidity and ionic strength. A saturated potassium chloride-
agar salt bridge connected this arm to a saturated calomel reference
electrode. The potential of the S.C.E. was checked against a Beckman
fiber junction S.C.E., both dipping into a saturated potassium chloride
solution, and found to agree within 0.2 mv. The resistance of the cell
system was found to be 255 i h ohms as measured with a Serfass Conduc-

tivity Bridge Model RC M15 on the 1000-cycle a.c. setting and using a

66

small platinum electrode in place of the dropping mercury electrode, D.M.E.
Within the precision desired, this small cell resistance made it un-
necessary to correct for iR drop in the cell. Capillary characteristics
of the D.M.E. were not determined in the presence of uranium solutions,
but these constants for the capillary were measured in the thallium
polarography work. Various initial voltages, span voltages, and sensi-
tivity settings were used; but the principal work was done with an
initial voltage of —0.30, a Span voltage of 1.0, and a sensitivity of
0.150 ua./mm. No damping was used. Measurement of the fifitial and

final potentials on some of the polarograms with a potentiometer proved
that the voltage read from the polarograms was precise enough, con-
sidering the precision desired. All current measurements were of appar-
ent maximum current, 143., the tOp of the pen excursions was used. Half-
wave potentials were simply measured from the polarograms (with resid—
ual current corrections), rather than from plots of log i/(id — i).X§'
applied potential. This less precise method was used for a number of
reasons: (1) the shifts in half-wave potentials, 51/2, were large,

(2) the uranium(IV)-(III) reduction wave was not "clean" because the
solutions contained some uranium(VI) and consequently the wave showed
some distortion at the low end because of a small apparent uranium
(VI)-(IV) reduction wave; and (3) the diffusion current, id, could not

be determined accurately because of the small plateau which slopes

into the hydrogen wave.

67

E. Studies of Thallium Solutions Without Uranium

In addition to spectrophotometric and polarographic examinations
of thallium(III) solutions containing tartaric acid—-but no uranium,
some qualitative experiments with 2,h-dinitrophenylhydrazine reagent

were performed on these solutions.

1. Spectrophotometric Studies of Thallium

Various spectrophotometric examinations of thallium solutions were
performed in the course of these investigations, most of which were
done with solutions the same as those used in the nflerence kinetic runs.
Careful examinations in the ultraviolet region were done with the
Beckman DU spectrophotometer using a matched set of Beckman stoppered
10 mm. silica absorption cells designed for use in the far ultraviolet
region. These solutions were examined freshly prepared and at later

times.

2. Polarographic Studies of Thallium

Polarographic investigations of thallium(III)-tartaric acid solu-
tions were conducted in essentially the same manner as were the polaro—
graphic studies of uranium. The solutions contained 9.00 x 10-3 M thal-
lium(III), so a sensitivity setting of 0.600 ua./mm. was convenient.
An initial voltage of +0.50 and a span voltage of 2.0 were commonly
employed. Unusual maxima were observed with thallium(III) solutions
in the absence of a maximum suppressor, but with solutions containing
0.01% gelatin no maxima occurred. Capillary characteristics were deter-
mined at an applied potential of -0.26 v. XE' the S.C.E. A drop time,

t, of 3.660 sec. and a rate of mercury flow, m, of 1.993 mg./sec. were

68
determined; thus m2/3tl/6 = 1.966 mg.2/3/sec.1/2. Half-wave potentials
were unimportant in the thallium polarography; rather, the technique
was used to follow thallium(III) concentrations as a function of time.
This was done by either repeatedly scanning and measuring the diffusion
current at a particular voltage, or more simply by just working at a
constant applied potential of 0.00 or -0.26 v. 23. the S.C.E., where

the diffusion current is a measure of thallium(III) concentration.

3. 2,b-Dinitrophenylhydrazine Experiments

A solution of 2,h-dinitrophenylhydrazine, DNP, was prepared ac-
cording to the procedure given by Shriner, Fuson,and Curtin.195 This
reagent was used in attempts to test qualitatively for products of
tartaric acid oxidation by thallium(III). These products could con-
ceivably be aldehydes or ketones and should consequently give a posi-

tive test with DNP.
F. Polarographic Studies of the Uranium(IV)—Thallium(III) Reaction

Standard kinetic runs were examined polarographically. The tech—
niques employed were similar to those previously described, except
that now outgassed aliquots of uranium(IV) and thallium(III) (both
containing 0.01% gelatin) were mixed, transferred to the "H" cell, and
diffusion currents were observed as a function of time. The disappear-
ance of thallium(III) was quantitatively followed. Appearance of

uranium(VI) was also observed.

V. RESULTS
A. Kinetic Observations

Except for some polarographic studies of the reacting system,
where thallium(III) concentration was followed, all the other kinetic
observations are the result of spectrophotometric monitoring of uranium(IV)

concentration.

1. Immerence Kinetic Runs

Two types of reference kinetic runs were performed. The first
type (called "unaccelerated") were done under Harkness' and Halpern's
conditions147 (described in Chapter II); the second type (called "ac-
celerated”) were done under these conditions plus 6.00 x 10-3 M in
tartaric acid. The kinetic data are presented in the Appendix, the
unaccelerated runs in Table VI, the accelerated ones in Table VII.
Graphs of l/[U] vs. t for these two types have been shown in Figure 3.
If the data are plotted as a second-order reaction with a rate law:

-d[U(IV)]/dt = k’[U(IV)][Tl(III)],

the experiments without tartaric acid yield linear second—order plots
(as is shown for a typical run in Figure 8 by line A). However, if
the same procedure is applied to the data from the experiments with
tartaric acid, nonlinear second—order graphs result (as shown by
line B in Figure 8). Examination of the absorbancy data as a function
of time reveals that the absorbancy approaches zero for the unaccelerated
runs, but levels off at an appreciable value for the accelerated runs.
Calculation of the apparent second-order rate constant, k', for the three

unaccelerated studies gave the values: 0.79 x 10-2, 0.7b x 10-2, and

69

7O

 

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([Hl/[IL]

70

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.I o .Hlomm IE NIOH X NN..O u .x m2 MIOH M Q4.m n o_HA>HvDH_ "d. Cdm
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([nJ/[Tl]

71
0.72 x IO-Z‘M-lsec-l. These are appreciably less than those found by
Harkness and Halpern147 and Quinn.38 Apparent second-order rate con—
stants could not be determined, of course, when the second—order graphs
were nonlinear. Identical runs were in very good agreement. Any dis-

agreement between duplicates was primarily due to inability to produce

identical initial uranium(IV) concentrations.

2. Effect of Water Quality
A set of unaccelerated runs using Quinn's stock solutions36 and
three different sources of water were performed. The waters were: dis—

tilled, demineralized, and conductance. The apparent second-order rate
2 —1

)

2.17 x 10‘2, and 2.12 x 10‘2 M

constants obtained were 2.18 x 10-
sec-l, respectively. The agreement of these values shows the source

of water had little or no effect on the kinetics of the reaction. The
much higher values of k' obtained in this set of experiments are in
agreement with the values obtained by Quinn using these same stock solu—
tions. The lower values of k' obtained using the present experimenter's

stock solutions, indicates that the value of the rate constant is de-

pendent upon the stock solutions employed.

3. Effect of Iron(III)

In an effort to explain the considerably lower values of k' ob-
tained in this work, a series of runs was performed in the presence of
low concentrations (1.00 x 10-5 to 25.0 x 10‘5 M) of iron(III). Iron(III)
is a possible impurity in the thallium(III) stock solution when it is
made from T1203 which has been prepared by oxidizing thallium(I) with
hexacyanoferrate(III) in base. The original kinetic data is shown in

Table VIII of the Appendix. Linear second—order graphs were obtained

 

72
in all cases. A graph of k' XE' iron(III) concentration is shown in
Figure 9. It is clear that small amounts of iron(III) produce a con-
siderable increase in k'; however, it seems likely that other impurities
may also contribute to the higher rates. Impurities such as lead and
zinc could be present in the thallium(III) stock solution. An electro—
lytically prepared thallium(III) solution probably has fewer impurities
than thallium(III) solutions prepared by chemical oxidations. The
present studies show that the use of electrolytically prepared solu-
tions yields lower rate constants, and are probably more pure. Impur-

ities in any of the other stock solutions could also be responsible for

the higher rates.

A. Effect of Dissolved Oxygen

The curvature present in the second-order graphs of tartaric acid
accelerated runs might have been due to the presence of dissolved oxygen.
Hence solutions for two rderence accelerated runs were prepared. One
set of solutions was swept with nitrogen for ten minutes prior to mixing;
the other set was swept with oxygen for ten minutes prior to mixing.
The kinetic data are shown in Table IX of the Appendix. The nitrogen-
swept run gave results identical, within experimental error, to the
reference accelerated runs. The second—order plots still showed the
same amount of curvature. The oxygen-swept run had an initial uranium(IV)
concentration almost 20% less than the other run, and similar curvature
was observed in the second-order graph.

Thus oxygen appears to have no effect on the reaction kinetics
other than to lower the initial uranium(IV) concentration by oxidation.

The presence of trace amounts of oxygen when the solutions are not swept

73

.m n-oH x oo.m

02:52; mm .72 x m.m a 3953 Mo oo.mm n dame mm.w u 1 m: on; u :5

.cofimubcoocoo AHHHVcotfl .WM wfiwpmcoo 3mm povuoupcooom oedema“? .m 93m:

AnoH x my fiAHHHVmeH

 

 

0m ma ma an NH 0H m o . a m o .

a a _ _ A _ _ _ . _ _ m
OJOA

O
o lmé
0
low
0
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O.m

7b
(as was the usual procedure) has a negligible effect on the reaction
kinetics, except for slight lowering of the uranium(IV) concentration
in the interval between solution preparation and the start of the re-
action. Consequently, with the method employed for determining initial
uranium(IV) concentrations, no precautions for exclusion of oxygen

were necessary.

5. Tartaric Acid Variation
In addition to four accelerated reference kinetic runs (with
[HZTarJO = 6.00 x 10“3 M), twelve runs with varying initial tartaric

3 to 20.0 x 10-3 M) were performed.

acid concentrations (0.50 x 10—
The kinetic data for the latter runs is shown in Appendix Table X. A
graph of absorbancy vs. time for some of these experiments (and an
unaccelerated run) is displayed in Figure 10. The data for the experi-
ments not graphed fall approximately where expected. Examination of
the initial lepes shows that the reaction is initially faster, the.
higher the tartaric acid concentration. The order of the reaction with
respect to tartaric acid can be determined from the initial SIOpes by
plotting the logarithm of the initial rates (or slopes) as a function
of the logarithm of the initial tartaric acid concentration. The slope
obtained after a least squares treatment indicates that the order with
respect to tartaric acid is 0.37 i 0.02 (see Figure 11).

The absorbancy in experiments with tartaric acid concentration
less than or equal to 1.00 x lO_3 M appears to eventually go to zero
(see Figure 10). The absorbancies in experiments of higher initial
tartaric acid concentrations do not go to zero, but instead level off
at appreciable values. The higher the tartaric acid concentration in

the range 2,00 “0‘3 to 20.0 x 10‘3 111, the lower the final absorbancy.

75

. 2 x a E OTENE 3835 385% mm m 2 x cod u OZHHUH:
II ”I. o o h o o .N O “I
m: Tea x :48 8 TB x S m u 23:: .o 00 mm u game .m m u 1 allow; u w E
.mddp oopmpoaooom Ufiom ompmpump meow pom oewu .m> zocmnpomnd .OH oudwflm

AniQH x .oomv mswh

 

 

 

 

 

 

mm ON 2 2 fl .2 2 m e a N o
_ _ -_ a _ I J _ _ _ 8.
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Xoueqaosqv

76

.cosomcpcmecou aHum oHemoumo HmeHcH mo eonemmoH .mm meme HmHoHcH eo ecpHcmmoH .HH mesmHe

AnoH x a.cH oHeme~:_vmoH

 

 

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e H- w H o N N N H N o N m N . -
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O 0 — F _ r h P NHmn

 

 

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(L01 x I_oas M u.

01p/°[(n1)nlp)50I

77
6. Uranium(IV) Variation

Six kinetic runs at varying initial uranium(IV) concentration
(1.h6 x 10”3 to 11.17 x 10_3 M) and the four accelerated reference
constituted the uranium(IV) variation. The kinetic data are shown
Table XI of the Appendix. Figure 12 is a graph of absorbancy vs.
for the uranium(IV) variation experiments. As can be seen from th
initial slopes, the reaction is initially faster the higher the in
uranium(IV) concentration. A graph of the logarithm of the initia
rate as a function of the logarithm of the initial uranium(IV) con
tration is displayed in Figure 13. The slope obtained after a lea
squares treatment indicates that the order with respect to uranium
is 0.93 i 0.03.

In all runs of this series, the absorbancies as a function of
approach nonzero values, as can be seen in Figure 12. The final a2
ancies appear to be directly related to the initial uranium(IV) co
centrations. Each reaction appears to proceed for a length of tim'
and then essentially stop. The approximate time at which each rea
stops is indicated by the short vertical lines in Figure 12. The i

the initial uranium(IV) concentration, the sooner the reaction stq

7. Thallium(III) Variation

The thallium(III) variation was comprised of six runs at vary
initial thallium(III) concentrations (b.50 x 10.3 to 22.5 x 10”3 M
the four accelerated reference runs. Table XII of the Appendix sh:
the kinetic data and Figure 1b shows a graph of absorbancy vs, tim<
these runs. The initial slopes are all essentially identical. Th1

graph of the logarithm of the initial rate as a function of the 101

78

n

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72 x 8.0 n 02.3me H: H...

OH xood n

.noH x m 5 0353 womeHccH Saga
CZECH: mo oo.mN n deg KN u a “FEH u H E
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79

 

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32 x a. E 0223: m3
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1 a _ 4 _ 4 _ _ 9 a
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mm

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:0

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Roueqdosqv

81
of the initial thallium(III) concentration (as shown in Figure 15) has
a slope of almost zero. Actually the slope obtained after a least
squares treatment was 0.05 i 0.0L, this being the calculated order with
respect to thallium(III).

It is seen in Figure lb that the absorbancies as a function of
time approach nonzero values for all runs of this series. The final
absorbancies decrease on going from [Tl(III)]O = b.50 x 10.3 to 9.00
x 10-3 M; then they increase on going from [Tl(III)]O = 9.00 x 10_3 to
22.5 x 10-3 M. Each reaction appears essentially to stop after a
certain length of time, this time being indicated by the short vertical
lines in Figure lb. The lengths of time are directly related to the

initial thallium(III) concentrations; the higher the initial thallium(III)

concentration, the later the reaction stops.

8. Effect of Uranium(VI), Thallium(I), and Solution Age

All kinetic runs containing tartaric acid exhibited departure
from second—order kinetics. An apparent decrease in reaction rate as
a function of time was consistently observed. This feature has been
shown for a typical run by the curvature in line B of Figure 8. The
decrease in rate might be due to a reaction by the products: uranium(VI)
and/or thallium(I). Another possibility considered was that some kind
of slow complexation was occurring in the solutions prior to starting
the runs and that this was somehow responsible for the curvature ob—
served in the second-order graphs. In order to test these hypotheses,
three runs were performed. The first was done under the conditions of
an accelerated reference run, except that the reacting solution con-

tained added uranium(VI) at an initial concentration of 3.38 x 10"3 M.

 

82

 

 

 

 

.coHemeocooeoo HHHHVesHHHmee HoHoHcH 0o aeeHommoH .mw some HmHoHcH eo eeeHeomoH .mH ooomHh
AooH x a eH oHAHHHvHeIV ooH

0.H- N.H- 0.H- o.H- 0.N- H.N- N.N- m.N- H.N-Nm.m-

.11 _ _ _ _ _ _
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Ts

(LOT X

83
The kinetic data for this experiment (H-l) is shown in Appendix Table
XIII. The second run was likewise performed under accelerated refer-
ence run conditions, except this time an initial thallium(I) concentra-
tion of 9.00 x lo”3 M was introduced. The data for this experiment
(H—2) are also in Table XIII. The third run, performed to learn of the
presence (or absence) of any slow complexation equilibria, was simply
a run performed with the same solutions used in an accelerated reference
run (B-3), but which were allowed to age for 2h hours. The data for
this experiment (H-3) are also in Table XIII.

The absorbancy vs. time data for the runs containing extra uranium(VI)
or thallium(I) agrees closely with the data from accelerated reference
runs. The second—order graphs likewise show similar curvature. It is
clear that neither uranium(VI) or thallium(I) has an effect on the
kinetics of accelerated runs.

The run performed with aged solutions had an initial uranium(IV)
concentration about 28% lower than the run performed 2h hours earlier
with the same solutions. This is presumably due to air oxidation of
uranium(IV) that is accelerated in the presence of tartaric acid. The
second-order graph for this run exhibits curvature, but the line has a
consistently higher slope than did the plot for the unaged run. In
other words, the rate of dissappearance of uranium(IV) appears to be
faster in the aged run. The presence of some slow complexation reac—

tion is not proved by this evidence, but it is possible.

9. Hydrogen Ion Variation
The hydrogen ion variation was comprised of five runs at varying

hydrogen ion concentrations (0.53 to 2.78 M) and the four accelerated

 

8b
reference runs. The kinetic data are given in Table XIV of the
Appendix. Figure 16 shows a graph of absorbancy vs. time for these
experiments. Examination of the initial SIOpes proves that the re-
action is initially faster, the lower the hydrogen ion concentration.
A graph of the logarithm of the initial rate as a function of the log—
arithm of the hydrogen ion concentration is presented in Figure 17.
The slope obtained after a least squares treatment indicates that the
order with respect to hydrOgen ion is -0.91 i 0.06.

Again, as in all previous runs in the presence of tartaric acid,
the absorbancies as a function of time approach nonzero values. This
feature is shown in Figure 16, where it can also be seen that the fina
absorbancies are higher, the lower the hydrogen ion concentration.
Again, each run appears to proceed for a length of time and then es-
sentially stop. Examination of the curves in Figure 16 shows that the

reaction stops sooner, the lower the hydrogen ion concentration.

10. Polarography of Kinetic Runs

Before examining polarographically solutions under unaccelerated
reference run conditions, three polarograms were recorded. The same
polarograph settings (previously described) were used for all these
polarograms. First a residual current polarogram was obtained for a
solution 1.76 M in hydrogen ion, 2.9 in ionic strength, and 0.01% in
gelatin. This is shown by line A in Figure 18. Then a polarogram for
a solution of the same acidity, ionic strength, and gelatin concentra-
tion, but containing approximately 3.50 x 10‘3 M uranium(IV) and some
and some uranium(VI) was recorded. Line B in Figure 18 shows this

polarogram. A polarogram under the same conditions, but containing

 

85

 

 

 

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Hz n-0H x 00.0

 

 

 

 

 

 

 

 

 

 

a 5 rs is: is; a is x as
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86

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Current (scale division)

87

 

250 l T

200-—-

 

 

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011111111
.5 .3 .l -.l -.3 -.5 -.7 -.9 -l.l -l.3
Applied potential (v. 32' S.C.E.)

Figure 18. "Polarograms."

Maximum pen excursions are plotted. Sensitivity = 0.600 par/Scale
division.

No damping. Temp:F = 27 i 2° C.

All solutions: [H J = 1.76 M; u = 2.9; 0.01 % gelatin.

Scan A: Residual current; No HZTar and [HzTar] = 6.00 x 1073 M.

Scan B: [U(IV)] £53.50 X lO-éM; Some U(VI); No HzTar and [HZTar]=
6.00 x lO-3M.

Scan c: [T1(III)1 =‘9.oo x 10'3M; No HZTar and [HzTar] =
6.00 x 1073M.

Scan D: Expected "palarogram" for solution [U(IV)] at3.50 x 1073M,
Some U(VI); [T1(III)] = 9.00 x io‘éy, and with and withofit
HzTar.

88
9.00 x 1073 M thallium(III) and no uranium, was also recorded and is
shown by line C in Figure 18.

Before polarographically examining solutions under accelerated
run conditions, three more polarograms were recorded. These were per-
formed with the same conditions as just described, but each solution
contained in addition 6.00 x 10.3 M tartaric acid. The polarograms
were identical, within experimental error, to those obtained for the
solutions without tartaric acid.

A solution containing both uranium(IV) and thallium(III) (before
any oxidation-reduction had occurred) would be expected to yield a
polarogram like line D in Figure 18. Such a solution would be identi—
cal to a reference kinetic run solution at mix time (except for the
presence of gelatin). Solutions like this, with an without tartaric
acid, were prepared and scanned polarographically as soon as possible.
The polarograms corresponded to line D of Figure 18 except for devia-
tions due to the extents of oxidation-reduction that had occurred in
the time interval. Additional scans of these solutions clearly showed
thallium(III) disappearing and uranium(VI) appearing. The more negative
applied potential side of the polarograms could be qualitatively inter—
preted as showing the appearance of thallium(I) and the disappearance
of uranium(IV).

Another set of these reference kinetic runs was prepared and exam-
ined as a function of time at a constant applied potential of 0.00 v.
(25. S.C.E.), where the diffusion current is a measure of thallium(III)

concentration. Data for these experiments is presented in Table XV of

the Appendix.

89

In the unaccelerated experiment (J-l), the thallium(III) concen-
tration was observed to decrease appreciably faster than the uranium(IV)
concentration had been found to decrease in spectrophotometric observa—
tions. The faster rate can be reasonably attributed to the fact that
the room temperature was a warm 290 C, instead of the customary 25° C
used in spectrophotometric runs. The final diffusion current was in
good agreement with that expected for the calculated concentration of
excess thallium(III) expected to remain. This reconfirms Harkness' and
Halpern's statement:147 "It was established that the stoichiometry of

the reaction conforms, within experimental error, to the equation:

 

U(IV) + Tl(III) > U(VI) + T1(I)."
An important observation made in the experiment containing tartaric
acid is that the thallium(III) was completely consumed, though in ex-
cess over uranium(IV). Despite lack of temperature control (room temper-
ature = 27° C) and the presence of gelatin, the approximate length of
time required to use up all the thallium(III) was similar to the length
of time the spectrophotometrically monitored runs proceeded until the
decrease in absorbance essentially stopped. The disappearance of thal-

lium(III) appeared to be fairly rapid during the first several hundred

seconds of reaction.

B. Observations of Uranium Solutions

In order to understand more fully the nature of the complicated re~
action system, solutions containing uranium(IV) and tartaric acid, but
no thallium, were investigated. These studies were performed to gain

information about the stability of uranium(IV), about the Species in

90

solution, and about any slower reactions that might be competing with
reactions involving thallium.

Careful spectrophotometric examination of the 650 mu absorption
peak of uranium(IV) showed no significant shift in the position of the
maximum with varying concentration of tartaric acid. In fact, the whole
visible spectrum of uranium(IV) is essentially the same in the presence
and absence of tartaric acid, all other variables being held constant.
Incidently, spectrOphotometric examination of a spent solution from an
accelerated reference run showed the maximum still at 650 mu.

Solutions 3.6 x 1o'3 14 in uranium(IV) and 1.8 x 10‘3 to 7.2 x 1o‘3
M in tartaric acid showed a decrease in absorbance (or uranium(IV)
concentration) of from 1% to 10% in two hours, the decrease being greater,
the higher the tartaric acid concentration. No particular precautions
to exclude air were employed. Acid and ionic strength were maintained
at 1.76 M and 2.9, respectively. These same solutions (stored in
stoppered flasks) when examined 2b hours later, showed an overall de-
crease in absorbance of from 32% to b5%, the extent of decrease still
in the same order as before. These observations indicate that solutions
of uranium(IV) containing tartaric acid are less stable to oxidation
than uranium(IV) solutions having no tartaric acid; the more organic
acid present, the more rapid the disappearance of uranium(IV). Since
it is unlikely that tartaric acid is oxidizing uranium(IV), these re-
sults are interpreted as oxidation of uranium(IV) by oxygen from the
air, the reaction being catalyzed by tartaric acid. This.conclusion has
been suggested by Quinn.38 More evidence for this interpretation follows.

On several occasions uranium(IV) and uranium(IV)-tartaric acid solu-

tions were prepared for kinetic runs and the unused portions were allowed

 

91
to stand in stoppered volumetric flasks for a day or more. The color
of uranium solutions not containing tartaric acid was observed to be—
come progressively only slightly lighter green in the course of a few
days. On longer standing, the now pale green color persisted and, to
the eye, appeared to no longer become paler. This certainly could be
interpreted as due to slow air oxidation of uranium(IV), the oxidation
ceasing when all the oxygen in the stoppered flask had been consumed.
On another occasion, two identical uranium(IV) solutions containing
tartaric acid were prepared and aliquots removed for use in kinetic
runs. The initial absorbance in these runs indicated that in the two
hours between sample preparation and the start of the run, the uran-
ium(IV) concentration had decreased about h% from that eXpected and
observed with uranium solutions free of tartaric acid. The two solu—
tions remained in stoppered volumetric flasks which were each only
opened for a few seconds in order to remove aliquots for the kinetic
runs. The next day the solutions were both much lighter green indicat-
ing appreciable decrease in uranium(IV) concentration. Another kinetic
run was performed using one of these now one day old uranium(IV)-tar—
taric acid solutions. This time the initial absorbance was down about
28% from that observed with fresh uranium solutions free of tartaric
acid. The volumetric flask was, of course, opened for several seconds
to remove the second aliquot. The following day, 1.3., the third day,
the two solutions were observed. The one that had not been opened on
the second day was definitely darker green than the other one. It can
be concluded that the repeated exposure to air causes more oxidation of

uranium(IV). Had the disappearance of uranium(IV) been caused by anything

92
other than air oxidation, the two solutions would have been the same
pale green color on the third day.

This conclusion is confirmed by Hoeschele,35 who has prepared
uranium(IV)—tartaric acid solutions that were thoroughly de-oxygenated
by nitrogen bubbling. No significant changes in the absorbancies at
650 mu of these solutions were observed in three weeks.

Polarographic investigations of uranium(IV)—tartaric acid solutions
were performed in order to estimate the extent of complexation under
the kinetic run conditions. Half-wave potentials, 51/2: of the uran-
ium(IV)-(III) reduction wave as a function of tartaric acid concentra-
tion are shown in Table IV. Large shifts in El/2 were observed which
indicates complexation. The value obtained for 51/? with 1.76 M tar—
taric acid is in doubt because the wave has shifted enough to be par-
tially merged with the hydrogen ion reduction wave.

Table IV. Half-wave potentials of the uranium(IV)-(III) wave as a
function of tartaric acid concentration

 

 

 

[HZTar] QM) El/Z (v. XE' S.C.E.)
nil -.816
0.00600 -.817
0.0500 -.853
0.500 -.898
1.76 “.922 (7)

Conditions are described in experimental chapter.

 

When the data are treated by conventional methods196 for polaro-
graphically determining the coordination numbers and stability constants

for stable metal ion complexes, the following results are obtained.

93
The value for p - q is 0.7h, where p is the coordination number of
uranium(IV) and q is the coordination number of uranium(III). Appar-
ently uranium(IV) can coordinate one more ligand than can uranium(III).

The value of Biv/BEII is about hO, where Biv is the over—all stability

II is the over-all

constant for uranium(IV) with p ligands and B:
stability constant for uranium(III) with q ligands. Clearly uranium(IV)
complexes appreciably better than uranium(III), as would be expected.
Since uranium(III) is probably similar to thallium(III) in complexing

ability, it seems reasonable to assume that complexes of uranium(IV)

are more important in this study than are complexes of thallium(III).
C. Observations of Thallium Solutions

Some properties of solutions containing thallium(III) and tartaric
acid, but no uranium, were investigated.

The ultraviolet spectra of thallium(III) solutions under kinetic
run conditions, with and without tartaric acid, were recorded and are
shown in Figure 19. Line A shows the shoulder of the thallium(III)
peak in the absence of tartaric acid. Six days later the same solution
was re-examined and no significant change could be detected. Line B
shows the thallium(III) spectrum in the presence of tartaric acid. A
large shift toward the visible is obvious, indicating some sort of
thallium(III)-tartaric acid interaction. This solution was re—examined
lb hours and six days later; no significant change could be detected.
Thus thallium(III) in the presence of tartaric acid appears stable. For
reference, a thallium(I) spectrum is displayed in Figure 19 by line C,

a tartaric acid spectrum is shown by line D, and a perchloric acid-sodium

9b

.aoHonz can a0H0: ch0

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.E n 0H x 00 0» HccH~0_ H0

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N z mnoH x 00. m u hAHHHvHHH "d

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Aoueqdosqv

 

 

 

 

9S

perchlorate spectrum by line E. The spectrophotometric results shown
by lines A and C are in agreement with the observations of Rogers and
Waind.106

Some polarographic studies of thallium(III)-tartart acid solutions
were performed. Polarograms for thallium(III) solutions under kinetic
run conditions (plus 0.01% gelatin) and in the absence or presence of
6.00 x 10_3 M tartaric acid are essentially identical. Such a polaro—
gram has been shown in Figure 18 by line C. In order to see if any
reduction of thallium(III) occurs in the presence of tartaric acid, the
thallium(III)-tartaric acid solution was repeatedly scanned and the
diffusion current, which is proportional to thallium(III) concentra-
tion at —0.26 v. vs. S.C.E., was measured. The results are shown in

Table V. There is clearly no significant change in thallium(III) con-

centration in over eight hours.

Table V. Diffusion current as a function of time for a thallium(III)-
tartaric acid solution

 

 

 

Time (min.) Diffusion Current(scale divisions)
17 110.8
3b 110.2
63 110.1
106 111.8
128 110.7
187 110.9
22h 110.5
3h? 110.6

51h 111.0

 

96

Since the products of oxidation of tartaric acid could conceivably
be aldehydes or ketones, qualitative tests with DNP were attempted. All
solutions tested were made 1.76 M in perchloric acid. A tartaric acid
solution, uranium(IV) solution, and a thallium(I) solution all gave
negative tests, as did a tartaric acid plus uranium(IV) solution. Solu-
tions containing thallium(III) and thallium(III) plus tartaric acid
gave positive tests. Identical results were observed when the solutions
were allowed to stand for one day before applying the DNP test. It is
clear that the presence of thallium(III) makes the results inconclusive
as a test for aldehydes or ketones. A solution 0.01 M in thallium(III)
and 0.05 M in tartaric acid was prepared and allowed to stand for as
long as three days before applying the DNP test in order to allow time
for the thallium(III) to be reduced. A positive test was still obtained.
Lastly a DMP test was applied to a spent solution from an accelerated
kinetic run in which all the thallium(III) was presumably reduced; a
positive test was observed. This result appears to be due to the pres-

ence of an aldehyde or ketone.

VI. DISCUSSION

The results have indicated that the reaction system is not simple.
The unaccelerated reaction exhibits second-order kinetics, but this is
no longer the case when tartaric acid is present. Experiments conducted
to determine the orders with respect to each of the individual react-

ants yield the equation:
-d[U(IV)]/dt = k[U(IV)]°°93[T1(III)]°'°5[H2Tar]°-37[H+]-°'91.

The fractional orders indicate a mechanism involving more than one
path. All attempts to fit the data to a simple rate law were unsuc—
cessful. Nevertheless, a number of features of the system are evident
and a mechanism qualitatively in accord with all the experimental
observations can be proposed.

In addition to the orders reported above, several particularly
important results are to be noted. The initial rate of disappearance
of uranium(IV) is progressively increased with increasing tartaric acid
concentration. All reactions in the presence of more than just a small
amount of tartaric acid proceed for a length of time and then essentially
stop, even though uranium(IV) is still present and was initially ex-
ceeded generally by almost a three-fold concentration of thallium(III).
The reactions stOp because the thallium(III) is exhausted. The only
reagent present capable of reducing the excess thallium(III) is tartaric
acid, yet thallium(III)—tartaric acid solutions are stable for rather
long times. Oxidation of uranium(IV) by oxygen of the air, even though
accelerated by the presence of tartaric acid, is still negligibly slow.
Oxygen exclusion, the presence of uranium(VI), and the presence of

97

98

thallium(I) all have no effect on the reaction kinetics.

Polarographic

studies of uranium(IV)-tartaric acid solutions indicate appreciable

complexation.

The following mechanism seems reasonable:

 

 

Y
U(IV) + T1(111)i-> U(VI) + T1(III) } PathA
ka
HZTar <— HTar- + H+
k
114+ + HT ’ ‘3") (UHT 3+)
ar <17:- ar I
k Path B
3+ 3+ 2 + 4+
(UHTar )I + T1 f§§t7> Tl + U + X
. __>ks . l
(UHTar3 ) < (UHTar:5 )
I k-3 II
>Path C
(UHTar3+) + "113+ + 2H 0 —ki—> 1‘1” + 110 2* + Mar" + 234
II 2 fast 2 J

Path A is the Harkness and Halpern mechanism.147

 

The first equa-

tion of path B is a slowly approached equilibrium between U‘+ and bi-

tartrate to give a uranium(IV) complex.

This complex (or perhaps the

tartaric acid part of this complex) reacts with T13+ in the second

equation of path B to form T1+, an oxidized tartaric acid species, X,

and leave uranium(IV) unoxidized.

of excess thallium(III).

This path accounts for the reduction

The first equation of path C is a slow equi—

librium rearrangement of the first uranium(IV) complex into a second

uranium(IV) complex.

This rearrangement could conceivably be a change

from a mono- to a bidentate complex or from a bi- to a tridentate

complex.

The second equation of path C is a reaction of the second

complex with T13+ to form Tl+ and uranyl ion. This path accounts for

the accelerating effect of tartaric acid on the rate of disappearance

of uranium(IV).

99

The choice of bitartrate as the complexing species is not proved,
but is based on the following logic. Under the strong acid conditions
employed, the concentration of tartrate is extremely small; so complexa-
tion by tartrate seems unlikely. The unionized tartaric acid molecule
is not charged, so presumably it would not complex very well. The order
of almost minus one in hydrogen ion indicates an activated complex that
has lost one hydrogen ion, as would be the case in bitartrate complex
formation. Another possibility, undistinguishable kinetically, would
be formation of the complex from UOH3+ and unionized tartaric acid.

Results of the tartaric acid variation are in accord with the pro-
posed mechanism. With no tartaric acid, the reaction proceeds by path
A. As the tartaric acid concentration is increased, this path becomes
less important and deviations from second-order kinetics are eXplained
by contributions from path B, in which uranium(IV) is a product. At
very low tartaric acid concentrations 85 1.00 x 10-3 M) where the absorb—
ancies are observed to approach zero, there is not enough tartaric acid
present to reduce all the thallium(III) by path B. Also the amount of
(UHTar3+)II formed is limited, so only small increases in the rate of
disappearance of uranium(IV) by path C are observed. At higher tartaric
acid concentrations, the kinetic data are consistent with paths B and
C. The fact that the final absorbancies decrease with increasing
initial tartaric acid concentration (in the range 2.00 x 10-3 to 20.0
x 10-3 M) is due to an increase of reaction proceeding via path C.

In the uranium(IV) variation, increasing the initial uranium(IV)
concentration would increase the amounts of both complexes formed. This

being the situation, path C explains the increase in initial rates and

 

100
path B explains the observation that the reaction stops sooner, the
higher the initial uranium(IV) concentration. The order with respect
to uranium(IV) of almost one agrees with the proposed mechanism.

The order with respect to thallium(III) was found to be almost
zero; this is what the proposed mechanism demands. In the thallium(III)
variation, it was seen that the length of time the reaction proceeded
was directly related to the initial thallium(III) concentration. This
substantiates the conclusion that the reaction stops because thallium-
(III) is used up. The final absorbancy achieved in the thallium(III)
variation was a minimum when the initial thallium(III) concentration
was 9.00 x 10-3 M. Higher and lower initial thallium(III) concentra-
tions both produced higher final absorbancies. The higher final ab-
sorbancies observed when the initial thallium(III) concentration is
high are explained if path B (which produces uranium(IV)) is favored
under these conditions. At low initial thallium(III) concentrations,
the interconversion of the two complexes may be essentially at equi-

librium. Depending on the magnitudes of k2, k3, k and k4, the

-3)
minimum in absorbance (or remaining uranium(IV)) might result. In

other words, at very low initial thallium(III) concentrations, the
equilibirum between the complexes can be maintained as complexII is
reacting with thallium(III) to convert a certain fraction of uranium(IV)
to uranyl ions. At somewhat higher initial thallium(III) concentrations,
the equilibrium between the complexes is still essentially maintained
and so reaction with now more thallium(III) converts larger fractions

of uranium(IV) to uranyl ions, and the final absorbancies are lower.

At very high initial thallium(III) concentrations, the concentration of

 

101
complexII is lowered below the equilibrium value and cannot be maintained
because k3 is too small; thus now smaller fractions of uranium(IV) are
converted to uranyl ions and the final absorbancies are higher.

With respect to the hydrogen ion variation, one would expect that
the lower the hydrogen ion concentration, the more of each bitartrate
complex there would be formed. According to path B, the reaction should
stop sooner, the lower the hydrogen ion concentration. This feature
was experimentally observed. According to path C, the initial rate of
disappearance of uranium(IV) should be faster, the lower the hydrogen

ion concentration. This also was eXperimentally observed.

 

VII. SUMMARY

The kinetics of the oxidation of uranium(IV) by thallium(III) in
the presence of tartaric acid were investigated in aqueous perchloric
acid solutions. Increases in initial rate with increasing tartaric
acid concentrations were observed, along with departure from apparent
second-order kinetics. The orders with respect to uranium(IV), thal-
lium(III), tartaric acid, and hydrogen ion were found to be 0.93, 0.05,
0.37, and -0.91, respectively. Thallium(III) in the presence of
tartaric acid was shown polarographically to be stable, yet under kin-
etic run conditions where thallium(III) is in excess over uranium(IV),
the reactions stopped because the thallium(III) was consumed. The
oxidation of uranium(IV) by oxygen of the air is catalyzed by the pres-
ence of tartaric acid, but this reaction is negligibly slow compared to
the uranium(IV)-thallium(III)-tartaric acid reactions. The presence
of oxygen, uranium(VI), and thallium(I) were shown to have no effect on
the reaction kinetics. A mechanistic interpretation has been develOped

to account for the observed results.

102

 

O\UIJ:"\.U

\1

10.

ll.

12.
13.
1h.

15.

16.

17.
18.

19.

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Rossotti, F. J. C., J. Inorg. Nucl. Chem., 3, 159 (1955).

Roig, E. and R. W. Dodson, J. Phys. Chem., 65, 2175 (1961).

Challenger, G. B. and B. J. Masters, J. Am. Chem. Soc., Z8, 3012
(1956).

138.

139.
1&0.
1&1.
1&2.

1&3.

1&&.
1&5.

1&6.
1&7.
1&8.

1&9.
150.

151.

152.

153.
15&.

110

Donas, T. P., Anales Fac. Quim. Farm., Univ. Chile, 12, 1&9
(1961); C. A., 5g, 287&f.

Johnson, C. B., Jr., J. Am. Chem. Soc., 7_&_, 959 (1952).
Forchheimer, 0. L. and R. P. Bpple, ibid.,‘ZQ, 5772 (1952).
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Duke, F. R. and B. Bornong, J. Phys. Chem., 69, 1015 (1956).
Armstrong, A. M. J. Halpern, and W. C. E. Higginson, ibid. ,60,

1661 (1956); A. M. Armstrong and J. Halpern, Can. J. Chem.,
_§, 1020 (1957)

 

Irvine, D. H., J. Chem. Soc., 1957, 18&1.

Higginson, W. C. B., D. R. Rosseinsky, J. B. Stead, and A. G.
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Sykes, A. G., J. Chem. Soc., 1961, 55&9.
Harkness, A. C. and J. Halpern, J. Am. Chem. Soc., él, 3526 (1959)-

Mellon, E. K., Jr., "The Reaction Between U(IV) and T1(III),"
Summer Report, IID-ll&66, Los Alamos (1960); N.S.A., 15,

12907.
Wear, J. 0., ”U.S. Atomic Energy Commission," SC-&972(RR) (196&).

Jones, F. A. and E. S. Amis, J. Inorg. Nucl. Chem., 36, 10&5 (196&).

Haynes, L. J., in B. H. Rodd(ed.), "Chemistry of Carbon Compounds,"
Vol. I, Part B, Elsevier, New York, 1952, p. 1171.

Prager, B., P. Jacobson, and later F. Richter(eds. ), "Beilstein’ s
Handbuch der organischen Chemie, " &th Ed. ,Springer, Berlin;
Main series, Vol. III, 1921, pp. b86- &89; lst Supp. , Vol.
III, 1929, pp. 171- 172, 2nd Supp. , Vol. III, 1982 pp 313-
31&; 3rd Supp., Vol. III, Part 2, 1962, pp. 999-1001.

Levesley, P. and W. A. Waters, J. Chem. Soc., 1955, 217.

Cabello, A. S. and H. S. Garcia, Anales Real Soc. Espan. Fis.
Quim (Madrid), &88, 281 (1952) (see C. A. &6, 8&82b);S
S. Perez and H. S. Garcia, ibid. ,&9B, 187 (T953)(see C. A,
&7, 119l2f); H. S. Garcia and M. G. L. Lucini, ibid. ,&9B,
355 (1953) (see c. A. ,&8, 6212b); v. M. Bhale, P. G. Sant,
and S. L. Bafna, J. Sci. —Ind. Res. (India), 15B, &5 (1956)
(see C. A. ,50, 99209); G. V. Bakore and R. Shanker, Indian
J. Chem., _, 108 (1963)(see C. A. ,59, 15135f).

155.

156.
157.

158.

159.

160.

161.

162.

163.

16h.
165.

166.
167.

168.

169.

111

Benrath, A. and K. Ruland, Z. Anorg. Allgem. Chem., 11b, 267 (1920);
H. H. Willard and P. Young, J. Am. Chem. Soc., 52, 132 (1930);
N. N. Sharma and R. C. Mehrotra, Anal. Chim. Acta, 11, bl?

(195b).
Srivastava, T. N. and S. P. Agarwal, J. Prakt. Chem., g, 319 (1957).

Barakat, M. 2., M. F. Abdel-Wahab, and M. M. El—Sadr, J. Chem.
Soc., 1956, b685.

Datar, D. S. and S. N. Kelkar, Current Sci. (India), 15, 3L8
(l9h6)(see c. A., £3, 2308b); T. L. Rama Char,'§uii. India
Sect. Electrochem. Soc., Z, 89 (1958)(see C..A., 53, lb807e);
L. S. Vaidya and D. S. Datar, J. Sci. Ind. Res. (I—ndia),

292, 35 (1961)(see C. A., 55, 197829).

Saxema, L. K. and C. P. Singhal, J. Indian Chem. Soc., 38, 3L6
(1961). "'

Srivastava, T. N., Z. Physik. Chem. (Leipzig), 399, 22 (1958).

Bothe, P., J. Prakt. Chem., 22, 191 (186A); A. Claus, Chem. Ber.,

Seekamp, W., Ann. Chem., 218, 373 (189D); C. Neuberg, Biochem. Z.,
_1_3_, 317 (1908). .

Penton, H. J. H., J. Chem. Soc., 65, 899 (189b); 81, 80A (1905);
W. H. Hatcher and M. G. Sturrock, Can. J. Research, 3, 21A

(1930)-
Sihvonen, V., Suomen Kemistilehti, 2E, 32 (1936); C. A., 31, 252b7.

See for example: S. N. Benson,"The Foundations of Chemical Kinetics,"
McGraw-Hill, New York, 1960; or A. A. Frost and R. G.
Pearson, "Kinetics and Mechanism," 2nd Ed., Wiley, New York,

1961.
See reference 38, pp. 19-20.

Dodson, R. W., A. C. Graves, L. Helmholz, D. L. Hufford, R. M.
Potter, and J. G. Povelites, in A. C. Graves and D. K.
Froman (eds.), "Miscellaneous Physical and Chemical Tech-
niques of the Los Alamos Project," McGraw-Hill, New York,

1952, pp. 1-6.

Kolthoff, I. M. and E. B. Sandell, "Textbook of Quantitative
Inorganic Analysis," 3rd Ed., Macmillan, New York, 1952,
pp. 526-529.

Gmelin, L., "Gmelin's Handbuch der anorganischen Chemie," 8th Ed.,
System-number 55, Verlag Chemie, G.m.b.H., Berlin, 1936,

p. 136.

170.

171.

172.

173.
178.

175.

176.

177.

178.

179.

180.

181.

182.
183.
18k.
185.

186.

187.
188.

112
Sato, T., Naturwissenschaften, L8, 668 (l961)(see C. A. ,56,
9683g); I. Kobayashi, Rika Gaku Kenkyusho Hokoku, 37,
3h9 (l961)(see C. A. ,56, 11165d).

Kolthoff, I. M., and R. Belcher, "VClumetric Analysis," Vol. III
Interscience, New York, 1957, pp. ll-lb.

)

Lundell, G. E. F. and H. B. Knowles, J. Am. Chem. Soc., D7, 2637
(1925). '—

Arhland, S. and R. Larsson, Acta Chem. Scand., 8, 137 (l95b).
Young, R. C., J. Inorg. Nucl. Chem., 7, D18 (1958).
Benson, E. P., Jr., private communication.

Casto, C. C” in C. J. Rodden(ed.), "Analytical Chemistry of the
Manhattan Project," McGraw-Hill, New York, 1950, p. 52b.

Heal, H. G., Report MC-95, Oct. 19th.
Pierle, C. A., J. Phys. Chem., 22, 517 (1919).

Hodgman, C. D.(ed.), "Handbook of Chemistry and Physics," 38th
Ed., Chemical Rubber, Cleveland, 1956, p. 30b0,

See reference 171, p. 37.

Fowler, R. M. and H. A. Bright, J. Res. Nat. Bur. Std., 15, L93
(1935). "'

See reference 171, p. 52.
See reference 179, pp. 2385-2386.
Meyer, F. R. and G. Ronge, Angew. Chem., 52, 637 (1937).

Dodd, R. E. and P. L. Robinson, "Experimental Inorganic Chemistry,"
Elsevier, Amsterdam, 1957, pp. 166- 167.

Roscoe, H. E. and C. Schorlemmer, "A Treatise on Chemistry,"
V01. II, Macmillan, New York;, 1889, p. 159; S. R. Carter
and N. H. Hartshorne, J. Chem. Soc., 1926, 363; N. H.
Hartshorne and J. F. Spencer, J. Soc. Chem. Ind. (London),
h5, b7hT (1926); E. E. Aynsley, T. G. Pearson, and P. L.
Robinson, J. Chem. Soc. , 1935,58.

Stone, H. W. and C. Beeson, Ind. Eng. Chem. Anal. Ed., 8, 188 (1936).

See reference 185, pp. 167-168.

189.

190.

191.

192.

193.

198.
195.

196.

113

Moeller, TL, "Qualitative Analysis," McGraw-Hill, New York, 1958,
p. 06.

Kolthoff, I. M., Rec. Trav. Chim., Q}, 172 (1922).
Zintl, B. and o. Rienacker, z. Anorg. Allgem. Chem., igg, 276 (1926).

Mellor, J. W., "A Comprehensive Treatise on Inorganic and Theor-
etical Chemistry," Vol. II, Longmans, Green and Co., London,

1922, p. 395.

Fred, M. and C. J. Rodden, in C. J. Rodden(ed.), "Analytical
Chemistry of the Manhattan Project," McGraw-Hill, New York,

1950, p. 5&3.
Ayres, G. H., Anal. Chem., 21, 652 (19h9).

Shriner, R. L., R. C. Fuson, and D. Y. Curtin, "The Systematic
Identification of Organic Compounds," hth Ed., Wiley, New
York, 1956, p. 111.

Kolthoff, I. M. and J. J. Lingane, "Polarography," Vol. I, 2nd
Ed., Interscience, New York, 1952, pp. 217-22h.

 

IX. APPENDIX

ORIGINAL KINETIC DATA

1111

115

Table VI. Unaccelerated reference runs

+
All runs= [H J = 1.76 M; u = 2.9; Temp. = 25.00 C; [T1(III)]O =
9.00 x 10‘3 M.

0.79 X 10_2 M-lsec-l.

Run A-l: [U(IV)]O = 3.65 x 10‘3 M; k'

0.71 x 10‘2 M-lsec-l.

Run A-2: [U(IV)]O 3.51 x 10’3 83 k' _

0.72 x 10"2 M_lsec-l.

3.11 x 10“3 M; k'

Run A-3: [U(IV)]O

 

 

 

 

 

 

t(sec) Run A—l . Run A-2 ,_;, Run A-3 ,4
‘71 log Tl/U A log Tl/U A log Tl/U
100 .208 .3911 .262 .1071 .196 .1200
200 .205 .3978 .199 .1110 .191 .1226
300 .203 .1001 .197 .1136 .192 .1255
500 .200 .1013 .191 .1178 .189 .1298
700 .197 .1082 .191 .1219 .187 .1327
1000 .193 .1138 .187 .1277 .181 .1371
1300 .189 .1193 .181 .1322 .180 .1133
1600 .186 .1236 .181 .1366 .176 .1195
1950 .181 .1310 .177 .1128 .173 .1512
2300 .176 .1387 .173 .1192 .170 .1592
2650 .172 .1151 .169 .1558 .166 .1660
3100 .168 .1516 .165 .1626 .162 .1731
3600 .161 .1585 .161 .1695 .158 .1802
5100 .113 .5096
7200 .130 .5386
16200 .085 .6771
18000 .078 .7068
21600 .067 .7601
75600 .011 1.172

86100 .005 1.791

 

 

116

 

 

 

Table VII. Accelerated reference runs
All runs: [8+] = 1.76 g; 1 = 2.9; Temp. 25.00 c; [T1(III)]O =
9.00 x 10’3 g; [HzTarJO = 6.00 x 10'3 M.

Run B—l: [u(1v)]o = 3.17 x 10'3 8.

Run B-2: [U(IV)]o = 3.17 x 10‘3 8.

Run B-3: [U(IV)]o = 3.07 x 10‘3 8.

Run B-1: [U(IV)]o = 3.01 x 10’3 m.

t(sec) RunAB-l RunAB-2 RunAB-3 RunAB-h

100 .175 .171 .169 .167

200 .168 .166 .162 .160
300 .163 .160 .156 .151
500 .153 .150 .118 .115
700 .115 .113 .110 .110
1000 .135 .132 .130 .128
1300 .126 .121 .121 .123
1600 .119 .117 .111 .115
1950 .111 .109 .107 .107
2300 .106 .101 .102 .102
2650 .101 .100 .096 .096
3100 .096 .091 .090 .090
3600 .091 .090 .086 .085
5100 .086 .079 .078
7200 .081 .078 .076
16200 .077 .071 .069
18000 .076 .070 .067
21600 .071 .067 .065

 

 

117

Table VIII. Iron(III) accelerated runs

+1

 

 

 

 

 

 

All FUHS= [H = 1.72 M; u = 2.9; Temp. = 25.00 C; [T1(III)]o =
9.00 x 10 8-
Run c-1: [U(IV)]O = 3.18 x 10'3 g; [Fe(III)]O = 1.00 x 10‘5 g;
1' = 1.01 x 10'2 M-lsec-l.
Run c-2: [U(IV)]o = 3.17 x 10‘3 MS [Fe(III)]o = 3.00 x 10‘5 M;
k' = 1.27 x 10.2 M_lsec-l. _
Run c-3: [U(IV)]O = 3.61 x 10‘3 Mi [Fe(III)10 = 5.00 x 10‘5 13
k' = 1.51 x 10’2 M'lsec-l.
Run c-1: [U(IV)]o = 3.51 x 10’3 M; [Fe(111)1O = 10.0 x 10'5 g;
k' = 1.83 x 10.2 M-lsec-l.
Run 0-5: [U(IV)]O = 3.62 x 10‘3 y; [Fe(111)1o = 17.0 x 10'5 g;
k' = 2.28 x 10-2 M_lsec-l.
Run c-6: [U(IV)]O = 3.71 x 10'3 g; [Fe(III)]o = 25.0 x 10‘5 13
k' = 2.59 x 10"2 M_lsec-l.
t(sec) Run C—l Run C-2 Run C-3
A log Tl/U A log Tl/U A log Tl/U
100 .198 .1153 .196 .1186 .203 .1021
200 .195 .1193 .193 .1226 .199 .1076
300 .193 .1221 .190 .1268 .195 .1130
500 .189 .1278 .185 .1311 .189 .1213
700 .186 .1322 .181 .1101 .181 .1286
1000 .181 .1397 .175 .1195 .178 .1376
1300 .176 .1175 .170 .1577 .171 .1189
1600 .171 .1555 .165 .1661 .161 .1606
1950 .166 .1639 .159 .1770 .158 .1711
2300 .163 .1691 .153 .1880 .150 .1863
2650 .158 .1781 .119 .1958 .115 .1962
3100 .152 .1891 .111 .5123 .138 .5110
3600 .116 .5013 .135 .5253 .130 .5291

 

 

118

Table VIII. continued

 

 

 

 

 

 

t(sec) Ru“ C-b Run C-S Run c-6
A log Tl/U A 10g TI7U A 10g Tl/UV
100 .197 .1151 .201 .1015 .207 .3921
200 .191 .1231 .195 .1125 .199 .1026
300 .188 .1277 .190 .1195 .192 .1121
500 .181 .1381 .180 .1311 .183 .1219
700 .175 .1175 .171 .1131 .176 .1357
1000 .167 .1608 .165 .1583 .161 .1553
1300 .160 .1729 .158 .1707 .151 .1732
1600 .153 .1859 .118 .1897 .115 .1907
1950 .116 .1997 .139 .5083 .136 .5097
2300 .139 .5111 .131 .5262 .127 .5305
2650 .133 .5276 .121 .5131 .119 .5506
3100 .121 .5191 .111 .5693 .109 .5782

3600 .115 .5730 .105 .5956 .100 .6061

 

119

Table IX. Nitrogen—anxioxygen-swept runs

Both runs: [H+] = 1.76 M; p = 2.9; Temp. = 25.00 C; [T1(III)]o =
9.00 x 10'3 MS [HZTar]o = 6.00 x 10‘3 1.
Run D—l: [U(IV)]O = 3.13 x 10"3 M; nitrogen bubbled through thallium
and uranium solutions for ten minutes prior to run.
Run D-2: [U(IV)]0 = 2.52 x 10'3 1; oxygen bubbled through thallium

and uranium solutions for ten minutes prior to run.

 

 

 

t(sec) RunAD—l RunAD—2
100 .175 .138
200 .168 .130
300 .162 .125
500 .152 .116
700 .115 .109
1000 .135 .103
1300 .126 .097
1600 .119 .091
1950 .112 .087
2300 .108 .082
2650 .103 .078
3100 .097 .071

3600 .091 .068

 

 

120

 

 

 

 

Table X. Tartaric acid variation runs
All runs: [n+1 = 1.76 y; p = 2.9; Temp. = 25.00 c; [T1(111)]0 =
9.00 x 10'3 1.
Run B-l: [U(IV)]O = 3.36 x 10‘3 M; [HZTarJO = 0.50 x 10"3 1.
Run B—2: [U(IV)]O = 3.29 x 10‘3 13 [HZTarJO = 1.00 x 10‘3 1-
Run 5-3: [U(IV)]O = 3.22 x 10‘3 y; [HZTarJO = 2.00 x 10'3 1.
Run 5-1: [U(IV)]O = 3.19 x 10‘3 M5 [HZTar10 = 3.00 x 10'3 m.
Run E—5: [U(IV)]O = 3.22 x 10'3 y; [HZTarJO = 1.00 x 10'3 1.
Run 8-6: [U(IV)]o = 3.20 x 10’3 y; [HzTarJO = 5.00 x 10’3 M-
t(sec) BEBKELl BEEK§:2 33332;; '322KELE RunAE-S RunAE-6
100 .190 .185 .181 .178 .177 .177
200 .187 .180 .176 .171 .170 .170
300 .185 .178 .172 .168 .165 .161
500 .181 .175 .169 .161 .157 .155
700 .179 .172 .166 .155 .151 .118
1000 .171 .168 .161 .118 .111 .139
1300 .171 .165 .160 .111 .139 .130
1600 .167 .163 .157 .111 .135 .121
1950 .161 .160 .151 .138 .130 .117
2300 .160 .155 .151 .135 .126 .111
2650 .156 .152 .118 .132 .122 .107
3100 .153 .119 .115 .130 .119 .103
3600 .118 .115 .113 .128 .116 .098
5100 .135 .131 .135 .123 .105 .093
7200 .123 .121 .129 .121 .102 .091
16200 .079 .086 .123 .118 .100 .081
18000 .072 .080 .122 .117 .100 .083
21600 .061 .070 .120 .116 .099 .081

 

 

121

Table X. Continued.

Run E-7: [U(IV)]o = 3.18 x 10'3 Mi [HzTar10 = 7.00 x 10‘3 1.
Run 5-8: [U(IV)]O = 3.18 x 10‘3 83 [HZTar]O = 8.00 x 10‘3 8-
Run 5-9: [U(IV)]0 = 3.21 x 10‘3 83 [HZTar]o = 10.0 x 10’3 1-
Run 5-10: [U(IV)]o = 3.18 x 10'3 Mi [HzTar]O = 12.0 x 10‘3 1.
Run 5-11: [U(IV)]o = 3.20 x 10‘3 13 [HZTar]o = 15.0 x 10‘3 1.
Run E-12: [U(IV)]O = 3.19 x 10’3 Mi [HZTar]o = 20.0 x 10‘3 M-

For runs with zero [HZTar]o see Table VI.
For runs with 6.00 x 10-3 M [HzTarJO see Table VII.

 

 

Run B—7 Run E-8 Run E-9 Run E-lO Run E-ll Run E-12

 

 

 

 

t<sec) A A A A A A
100 .171 .173 .176 .171 .171 .175
200 .166 .166 .169 .161 .166 .166
300 .160 .160 .163 .157 .160 .159
500 .150 .119 .152 .116 .118 .116
700 .113 .111 .113 .110 .139 .137

1000 .131 .130 .130 .128 .126 .121

1300 .126 .120 .121 .120 .116 .112

1600 .120 .111 .115 .112 .107 .103

1950 .115 .105 .105 .105 .098 .092

2300 .108 .098 .099 .099 .091 .086

2650 .105 .092 .091 .095 .085 .079

3100 .100 .087 .085 .088 .077 .070

3600 .095 .082 .079 .081 .071 .061

5100 .081 .075 .071 .066 .061 .052

7200 .080 .072 .069 .061 .058 .050

16200 .075 .065 .063 .058 .050 .012
18000 .071 .061 .062 .057 .019 .011

21600 .072 .061 .060 .055 .017 .039

 

 

Table XI.

A11 runs:

Run F—l:
Run F-2:
Run F-3:
Run F—1:
Run F-5:
Run F-6:

122

Uranium(IV) variation runs

[11*] =

9.00 x 10‘3
[U(IV)]O = 1.16 x
[U(IV)]O = 1.73 x
[U(IV)]O = 6.27 x
[U(IV)]O = 7.86 x
[U(IV)]o = 9.67 x

[U(IV)]O =11 17 x _
For runs with [U(IV)]o = 3.01 x 10‘3 to 3.17 x 10-3 1 see Table VII.

10
10
IO
10
10
10

-3

‘3

1.76 M; 11 = 2.9; Temp.

‘3

l3 I3 I:

‘3
—3

n: he

-3

3

= 250 c; [T1(111)]O
13 [HzTar]O = 6.00 x 10‘

3

M.

n—
-.

 

 

 

 

t(sec) RunAF—l RunAF-2 RunAF-3 RunAF-1 RunAF-S RunAF-6
100 .080 .261 .316 .136 .536 .620
200 .076 .251 .335 .125 .520 .603
300 .073 .211 .327 .115 .510 .592
500 .068 .233 .311 .100 .192 .576
700 .061 .221 .302 .391 .182 .563
1000 .059 .211 .288 .382 .169 .519
1300 .051 .206 .277 .373 .160 .513
1600 .051 .199 .270 .369 .153 .510
1950 .017 .191 .265 .363 .150 .539
2300 .011 .185 .263 .357 .150 .539
2650 .010 .179 .260 .355 .150 .539
3100 .037 .173 .259 .351 .119 .539
3600 .031 .170 .259 .351 .118 .538
5100 .025 .167 .256 .352 .115 .535
7200 .019 .166 .251 .351 .113 .533
16200 .006 .159 .216 .316 .137 .521
18000 .006 .158 .215 .315 .136 .523
21600 .005 .155 .212 .313 .131 .521

 

123

Table XII. Thallium(III) variation runs

+
All runs: [H 1 = 1.76 85 1 = 2.9; Temp. = 25.00 c; [HZTar]o

6.00 x 10‘3 1.
Run 0-1: [U(IV)]O = 3.07 x 10'3 MS [Ti(111)]o = 1.50 x 10‘3 H.
Run 0-2: [U(IV)]o = 3.13 x 10“3 Mi [T1(111)]o = 6.00 x 10‘:5 1-
Run 0-3: [U(IV)]O = 3.13 x 10‘3 M3 [T1(111)]o = 7.50 x 10‘3 1.
Run 0-1: [U(IV)]O = 3.11 x 10'3 MS [T1(111)]O = 13.5 x 10‘3 1.
Run 0—5: [U(IV)]O = 3.07 x 10‘3 MS [T1(111)]o = 18.0 x 10‘3 1-
Run 0-6: [U(IV)]O = 3.15 x 10‘3 13 [T1(111)]0 = 22.5 x 10‘3 8-

For runs with [T1(111)]O = 9.00 x 10‘3 1 see Table v11.

 

 

Run G-l Run G-2 Run G-3 Run G—1 Run G-5 Run G-6

 

 

 

t<sec> A A A A A A
100 .169 .172 .172 .171 .170 .175
200 .163 .166 .166 .166 .163 .168
300 .156 .160 .160 .160 .158 .163
500 .118 .151 .151 .150 .119 .157
700 .112 .111 .113 .111 .111 .151

1000 .131 .135 .133 .136 .139 .118

1300 .128 .128 .121 .130 .135 .111

1600 .123 .122 .118 .125 .131 .111

1950 .118 .115 .110 .119 .128 .138

2300 .111 .110 .101 .116 .125 .135

2650 .110 .106 .100 .110 .122 .132

3100 .106 .102 .091 .107 .119 .129

3600 .103 .098 .090 .102 .116 .126

5100 .100 .091 .086 .089 .107 .118

7200 .099 .093 .081 .086 .103 .111

16200 .093 .087 .076 .080 .098 .110
18000 .092 .085 .075 .079 .097 .109

21600 .090 .083 .073 .077 .096 .108

 

121

Table XIII. Uranium(VI), thallium(I) and aged runs

3

A11 runs: [H+] = 1.76 H; 1 = 2.9; Temp. = 25.00 c; [T1(111)1O =
9.00 x 10‘3 H; [HzTar]o = 6.00 x 10'3 M-

Ruh H—1: [U(IV)]O 3.16 x 10‘3 H; [U(VI)]O 3.38 x 10‘3 1-

Run H-2: [U(IV)]O 3.13 x 10-3 1; [T1(1)]o 9.00 x 10‘3 M

Run H-3: Same solutions as used in Run B-3, except now 21 hours older;

[U(IV)]O = 2.21 x 10'3 M

[I
ll

 

 

 

t(sec) M M W
A A A
100 .171 .169 .119
200 .161 .161 .113
300 .157 .155 .108
500 .118 .116 .100
700 .111 .110 .091
1000 .131 .131 .086
1300 .121 .121 .079
1600 .119 .119 .073
1950 .112 .113 .068
2300 .107 .109 .063
2650 .103 .103 .059
3100 .097 .101 .055
3600 .091 .096 .051
5100 .082 .083 .010

7200 .080 .081 .037

 

 

125

Table XIV. Hydrogen ion variation runs

A11 runs: 1 = 2.9; Temp. = 250 c; [T1(111)]o = 9.00 x 10‘3 m;

[HzTarJO = 6.00 x 10'3 M.
+

 

 

 

Run 1—1: [U(IV)]o = 3.02 x 10 M; [H J = 0.53 M.
Run I-2: [U(IV)]o = 3.03 x 10‘3 M3 [H+] = 0.90 M.
Run I-3: [U(IV)]O = 3.05 x 10‘3 y; [H+] = 1.31 1-
Run 1-1: [U(IV)]O = 3.01 x 10‘3 M; [H+] = 2.25 M.
Run 1-5: [U(IV)]O = 3.01 x 10'3 M; [H+] = 2.78 M.
For runs with [H+] = 1.76 M see Table VII.
t(sec) RunAI-l RunAI-2 RunAI-3 RunAI—1 RunAI-S
100 .152 .160 .161 .168 .169
200 .111 .152 .155 .160 .162
300 .137 .116 .150 .151 .155
500 .129 .139 .112 .115 .116
700 .125 .132 .135 .139 .111
1000 .120 .125 .127 .130 .131
1300 .118 .120 .123 .123 .125
1600 .117 .115 .118 .117 .119
1950 .116 .110 .111 .110 .112
2300 .115 .106 .109 .105 .108
2650 .115 .105 .105 .101 .101
3100 .115 .103 .099 .096 .099
3600 .111 .102 .095 .091 .091
5100 .112 .100 .091 .079 .085
7200 .110 .099 .089 .070 .078
16200 .103 .091 .082 .065 .065
18000 .101 .089 .080 .061 .061

21600 .098 .086 .077 .062 .061

 

 

126

Table XV. Polarographic monitoring of thallium(III) concentration in
reference runs

Applied potential = 0.00 v; (E? S.C.E.). Scale divisions are for
maximum pen excursions. Corrections for redisual current
have been made. Sensitivity = 0.600 ua./sca1e division.

N0 damping. 9.00 x 10.3 M T1(III) = 109.8 scale divisions;
1 scale division = 0.08197 x 10‘3 M T1(III).

Both runs: [H+] = 1.76 M; 1 = 2.9; 0.01% gelatin; [U(IV)]O :5 3.50 x
10'3 M; [T1(111)]O = 9.00 x 10'3 m.

Run J-1: N0 HZTar; Room temp. = 290 C.

Run J—2: [HzTar]O = 6.00 x 10_3 M; Room temp. = 270 C.

 

 

 

 

 

 

Run J-l RU“ J’2
t(S€C) scale divisions scale divisions
100 107.3 92-6
200 105.1 82.6
300 103,7 71.7
500 101.1 62-0
700 100.0 53-1
1000 97-5 82'2
1300 95-3 33‘8
1600 93-1 2M
1950 90.7 21.1
2300 88.1 16-7
2650 87-3 13'3
3100 81-0 9‘7
3am 8L5 67
1500 79-5 8‘1
5100 78.7 3'2
7200 76.7 1.2
9000 75.2 0.7
16200 72-0
23100 71.5
27000 71.0
28800 70-8

 

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