Citation
Stability constants and structures of metal complexes of the antithyroid thiouracils

Material Information

Title:
Stability constants and structures of metal complexes of the antithyroid thiouracils
Creator:
Weber, Dennis Joseph, 1934-
Publication Date:
Language:
English
Physical Description:
xi, 241 leaves : illustrations; 29 cm.

Subjects

Subjects / Keywords:
Pharmaceutical Preparations -- analysis ( mesh )
Antithyroid Agents ( mesh )
Genre:
bibliography ( marcgt )
non-fiction ( marcgt )
Academic theses ( lcgft )
Academic theses ( fast )

Notes

Thesis:
Thesis (Ph.D.)--University of Florida, 1967.
Bibliography:
Includes bibliographical references (leaves 235-240).
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Dennis Joseph Weber.

Record Information

Source Institution:
University of Florida
Holding Location:
University of Florida
Rights Management:
Copyright Dennis Joseph Weber. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
Resource Identifier:
029765419 ( ALEPH )
26246820 ( OCLC )

Downloads

This item has the following downloads:


Full Text
STABILITY CONSTANTS AND STRUCTURES
OF METAL COMPLEXES OF THE
ANTITHYROID THIOURACILS
By
DENNIS JOSEPH WEBER
A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
December, 1967


DEDICATION
To my wife,
Shirley,
whose encouragement, help and patience
made this meaningful and possible.


ACKNOWLEDGMENTS
The author gratefully acknowledges Dr. Edward R.
Garrett, chairman of the supervisory committee, for his
research guidance and his assistance in the preparation
of this manuscript.
The author is most grateful for the laboratory assis
tance of his wife, Shirley A. Weber, and for her patience
and understanding.
Sincere appreciation is expressed to Dr. Melvin J.
Fregly for his willing help in the animal studies.
The author extends his sincere thanks to the many
people of the College of Pharmacy for their time and
energy which were very helpful.
The author is indebted to Dr. Alan Agren of the
University of Uppsala, Sweden, for the helpful information
and discussions he so willingly gave.
This investigation was supported by a Public Health
Service Fellowship from the Division of General Medical
Sciences, Public Health Service. The author is deeply
grateful for this support.
The author wishes to thank his typist, Mrs. Arthur
Grant for a superb job.
iii


Appreciation is extended to Dr. Melvin J. Fregly,
Dr. Werner M. Lauter and Dr. R. Carl Stoufer for serving
on the supervisory committee.
iv


TABLE OF CONTENTS
Page
ACKNOWLEDGMENTS iii
LIST OF TABLES vii
LIST OF FIGURES ix
INTRODUCTION 1
EXPERIMENTAL 9
I. Materials 9
II. Potentiometric Titrations.... 15
III. Ultraviolet Spectral Studies 19
IV. Polarographic Studies 29
V. Synthesis of Complexes 35
VI. Animal Experiments 41
VII. Magnetic Susceptibility Measurements 43
EQUATIONS 44
I. Potentiometric Titrations 44
II. Polarography 58
III. Solubility Analysis 69
RESULTS ?2
I. Potentiometric Titrations 72
II. Ultraviolet Spectral Studies 83
III. Polarography 92
IV. Animal Experiments 96
v


Page
DISCUSSION 100
I. Structure of Precipitated Cupric-Thiouracil
Complexes 100
II. Structure and Stability of Cupric-Thiouracil
Complexes in Homogeneous Solutions 105
III. Cadmium and Lead Complexes of Thiouracils.... Ill
IV. Complexation of Other Metal Ions with
Thiouracils 124
V. Biological Effects 129
SUMMARY AND CONCLUSIONS 135
APPENDIX
A. Tables 139
B. Figures..... 1?6
BIBLIOGRAPHY 235
BIOGRAPHICAL SKETCH 241
vi


LIST OF TABLES
Page
Table
I COMPOSITION OF SOLUTIONS AND ESTIMATED
STABILITY CONSTANTS AT VARIOUS TEMPERATURES
ON THE POSTULATE OF 1:1 METAL COMPLEXES OF
SUBSTITUTED THIOURACILS FROM POTENTIOMETRIC
TITRATIONS 140
II COMPOSITION OF SOLUTIONS AND ESTIMATED
LOGARITHMIC STABILITY CONSTANTS, K1 AND Kg,
AT VARIOUS TEMPERATURES ON THE POSTULATE OF
MIXED 1:1 AND 2:1 METAL COMPLEXES OF SUBSTI
TUTED THIOURACILS FROM POTENTIOMETRIC
TITRATIONS 144
III COMPOSITION OF AQUEOUS SOLUTIONS AND ULTRA
VIOLET ABSORBANCE OF AQUEOUS THIOURACIL-METAL
ION MIXTURES AT 25 149
IV INTRINSIC SOLUBILITIES OF LIGANDS AND pK*
St
VALUES DETERMINED FROM SOLUBILITY DATA AT 25.. 157
VEFFECTS OF CONCENTRATION OF CUPRIC ION AND PER
CHLORIC ACID ON DIFFUSION CURRENT, HALF-WAVE
POTENTIAL, E1/2 AND Ey^ VALUE FOR
POLAROGRAPHIC REDUCTION OF CUPRIC NITRATE 15
VIEFFECTS OF 6-n-PROPYL-2-THIOURACIL AND PER
CHLORIC ACID CONCENTRATIONS ON E^ AND
(Ei ) (E)_ VALUES OF 2.00 x lO"4 MOLAR CUPRIC
ION AT 21. 160
VIIEFFECT OF 5,6-DIMETHYL-2-THI0URACIL ON RAT
THYROID WEIGHT AFTER TWO WEEKS OF DRUG DIET.... 161
vii


Table Page
VIIIEFFECT OF 6-METHYL-N,N'-DIETHYL-2-THI0URACIL
ON THE RAT THYROID WEIGHT AFTER TWO WEEKS OF
DRUG DIET 164
IXEFFECT OF 6-n-PROPYL-2-THIOURACIL AND 5-METHYL-
-2-THIOURACIL ON THYROID WEIGHT AFTER TWO WEEKS
OF DRUG DIET 166
XNEGATIVE LOGARITHMS OF DISSOCIATION CONSTANTS
(pK*) OF LIGANDS AS A FUNCTION OF TEMPERATURE.. 170
81
XIENTROPY AND ENTHALPY VALUES FOR THE IONIZATION
OF THIOURACILS (HU) AND FOR THE FORMATION OF
MU+ BY COMPLEXATION WITH CADMIUM AND LEAD 171
XIICOMPOSITIONS AND ABSORBANCES OF 2-THIOURACIL-
- CUPRIC ION SOLUTIONS FOR JOB*S PLOTS 172
XIIINEGATIVE LOGARITHM OF THE RATIO OF THE DIS
SOCIATION CONSTANT OF THE OXIDIZED COMPLEX TO
THE DISSOCIATION CONSTANT OF THE REDUCED COM
PLEX OF CUPRIC-THIOURACILS (KQ/Kr) AS A FUNC
TION OF TEMPERATURE AND ACID CONCENTRATION 173
viii


LIST OF FIGURES
Figure Page
1 Potentiometric titration curves of filtered,
day old cupric nitrate-2-thiouracil mixtures
with ¡1 = 0.006 at 25.0 178
2 Ultraviolet spectra of mixtures of cupric
nitrate and 6-methyl-2-thiouracil in water
at 25.0 180
3 Ultraviolet spectra of mixtures of lead nitrate
and 2-thiouracil in water at 25*0 182
4 Ultraviolet spectra of ferric nitrate and
6-methyl-2-thiouracil in water at 25*0 184
5 Ultraviolet spectra of mixtures of cupric
nitrate and N,N -diethyl-6-methyl-2-thiouracil
in water at 25*0..... 186
6 Ultraviolet spectra of mixtures of cupric
nitrate and 2-ethylmercapto-4-hydroxypyrimldine
in water at 25*0 188
7 Plot of absorbance at 272 mp. versus the mole
fraction of cupric nitrate of aqueous solutions
of cupric nitrate and 2-thiouracil at 25.0 190
8 Jobs continuous variations plots of absorbance
of aqueous mixtures of cupric nitrate and
2-thiouracil 192
9 Plots of ultraviolet absorbance at 272 and 345
mp versus [Cu+2]/[2TU] 194
10 Ultraviolet spectra of aqueous 2-thiouracil
solutions as a function of pH at 25*0 196
11 Plots of absorbance of aqueous 2-thiouracil
solutions versus pH 198
12 Ultraviolet spectra of an aqueous solution
containing cupric nitrate and 2-thiouracil as
a function of pH at 25,0 200
ix


Page
Figure
13
14
15
16
17
18
19
20
21
22
23
Plots of absorbance of an aqueous solution
containing cupric nitrate and 2-thiouracil
versus pH.... 202
4
Polarogram of 2.00 x 10 M cupric nitrate in
the presence of 1.875 X 10"3 M 2-thiouracil and
0.0800 M HC104 at 21.0 204
Potentiometric titration curves of aqueous
mixtures of cupric nitrate and 2-thiouracll
containing precipitated complex with u = 0.006
at 25.0 206
Potentiometric titration curves of aqueous
solutions of lead nitrate and 2-thiouracil
with p = 0.006 at 25.0 208
Plot of log (l-/ against the negative logarithm
of 6-n-propyl-2-thiouracil anion concentration
obtained from lead nitrate -PTU mixture in
water with i = 0.006 at 25*8 210
Plot of log (1-J/ against the negative logarithm
of 6-n-propyl-2-thiouracil anion concentration
obtained from lead nitrate PTU mixture in water
with = 0.006 at 25*8 212
Plot of log (l-)/ against the negative
logarithm of 6-n-propyl-2-thiouracil anion
concentration obtained from lead nitrate PTU
mixture in water with ju = 0.006 at 25*8 214
Plot of log (l-)/ against the negative
logarithm of 6-n-propyl-2-thiouracil anion
concentration obtained from lead nitrate PTU
mixture in water with ^ = 0.006 at 25.8
Plots of /(1-)[2TU] against
[2TU~] from
aqueous mixtures of 2-thiouracil and lead
nitrate with p = 0.006 at 25*0
216
218
Potentiometric titration curves of aqueous
mixtures of cadmium nitrate and 6-n-propyl
-2-thiouracil with p = 0.006 at 250.
Potentiometric titration curves of aqueous
mixtures of nickel nitrate and 2-thiouracil
with p = 0.006 at 25.0 222
x


Page
Figure
24 Plot of the change in free energy (AF) versus
the absolute temperature calculated from the
stability constants of the 1:1 complexes (MU+)
for the cadmium and lead nitrate complexes of
2-thiouracil, 6-n-propyl-2-thiouracil,
5,6-dimethyl-2-thiouracil, 6-methyl-2-thiouracil,
5-methyl-2-thiouracil and 5-carboethoxy-2-
-thiouracil 224
25 Plot of the change in free energy (AF) versus
the absolute temperature calculated from the
acid dissociation constant of 2-thiouracil,
6-n-propyl-2-thiouracil, 51 6-dimethyl-2-
-thiouracil, 6-methyl-2-thiouracil, 5-methyl-
-2-thiouracil and 5-carboethoxy-2-thiouracil.... 226
26 Ultraviolet spectra of aqueous solutions at 250
containing 2-thiouracil and cupric nitrate 228
2? Plot of log (Ab AbH+)/AbH+ against pH for
aqueous, saturated 2-thiouracil solutions at 25*0
according to log (Ab AbH+)/AbH+ = pH pK^.... 230
28 Plots of -[(Ei)c (Ei)g] against the negative
logarithm of the 6-n-propyl-2-thiouracil
concentration 232
29 Plot of the logarithm of the Apparent Relative
Antithyroid Activity versus the negative
logarithm of the dissociation constant ratio
of the 2:1 cupric-ligand and 1:1 cuprous-ligand
complexes with thiouracils 234
xi


INTRODUCTION
The compound 2-thiouracil (I) is used in the treatment
of hyperthyroidism.
(I) H
In general, thiouracils are rapidly absorbed by the
gastrointestinal tract (1) and 50% of the administered
dose is excreted in the urine (2,3) At least l$% of
an oral dose is destroyed in the G. I. tract (4,5) and
is not available for absorption. The identity of poten
tial metabolites is unknown (4).
The probable mode of action of thiouracil is inter
ference with the incorporation of iodide into thyronine,
or its precursors, to produce thyroxine (6,7).
Thyroxine (V) is the hormone, elaborated by the
thyroid gland, which controls the rate of metabolism
and oxygen consumption of the body. High blood levels
1


2
of thyroxine, indicative of a hyperactive thyroid gland,
can cause loss of weight and high blood pressure (8). A
low blood concentration of thyroxine stimulates the ante
rior pituitary to secrete Thyroid Stimulating Hormone
(TSH) which causes an Increased rate of production of
thyroxine by the thyroid gland, a negative feedback mecha
nism. It has been suggested (8) that the production of
thyroxine proceeds by iodination of two molecules of the
amino acid tyrosine (II) which condense to form a molecule
of thyroxine (V) (Scheme I).
II Tyrosine
III Monoiodotyrosine
IV Diiodotyrosine
2 H
CH^HNH^OOH > H
ch2chnh2cooh
V Thyroxine
Scheme I


3
Goiter, the enlargement of the thyroid, gland, can
occur in two ways (8). Dietary deficiency of iodide will
result in a compensating hyperplasia of the thyroid to
trap as much of the available blood iodide as possible.
This condition, nontoxic goiter, usually results in normal
thyroxine blood levels. If the negative feedback mechanism
controlling the blood level of thyroxine fails, then the
anterior pituitary output of TSH escapes control by thyrox
ine and the thyroid gland is stimulated to produce thyrox
ine at an abnormal rate. This condition is termed toxic
goiter and can be controlled by antithyroid drugs such as
the thiouracils.
Alkyl substitution at the 5 r 6 position of thioura-
cil (I) usually enhances its antithyroid activity (9.11 )
For example, the 5-methyl, 6-methyl and 5.6-dimethyl sub
stituted thiouracils are active, with potencies relative
to 2-thiouracil, of 0.7, 1.0 and 1.2, respectively (9,10).
The present antithyroid derivative of choice, because of
its maximal activity and low toxicity in the intact animal
is 6-n-propyl-2-thiouracil (10,12,13). Alkylation of
thiouracil at the N-l, N-3 or sulfur positions greatly
reduces (10) and substitution by electronegative groups
at the 5 or 6 position nearly eliminates any antithyroid
activity (9-11).
The oxidation of thiouracil (I) by iodine has been
shown to occur with ease at physiological pH values (14).


4
Thiouracil is selectively oxidized by iodine even in the
presence of tyrosine (II) (14),the thyroxine (V) precursor
(15). In addition, no iodine was found in the recovered
tyrosine (14) so no iodinated tyrosine (111,1V) was formed.
The product of the iodine oxidation of 2-thiouracil is the
The disodium salt of VI is stable but the free acid readily
disproportionates to thiouracil and higher oxidation pro
ducts (14). The ease of oxidation of 2-thiouracil might
explain the degradation observed in the gastrointestinal
tract. This possibility is even more attractive when the
hydrolytic stability of thiouracil is considered (16).
It must be emphasized however, that ease of oxidation by
iodine cannot be the only determining factor since thioura
cil derivatives with no antithyroid activity are also
easily oxidized (14).


5
Cupric ion has been implicated in thyroid function
(17). The copper content of the normal and pathologic
thyroid has been determined (18) and verified by Kasanen
and Viitanen (19) who found elevated copper levels in
toxic and nontoxic goiters. The formation of diiodo-
tyrosine and thyroxine is increased when cupric ion is
added to homogenates of thyroid gland (20). Other evidence
that cupric ion aids in the formation of thyroxine, by
formation of iodine from iodide, has been presented (21,
22,23 ).
Since cupric ion and other heavy metals precipitate
thiouracil and its derivatives from aqueous solution (24,
25). Libermann conjectured that complexing ability and
antithyroid activity may be correlated. He assumed that
completeness of precipitation could be taken as a measure
of the stability of the complex. However, this assumption
is not always true. The extreme stability of some EDTA
complexes and their very high water solubility is a case
in point. He suggested a structure (VII) for the cupric
-thiouracil complex (25) which assumed a 1:1 stoichiometry
of metal to ligand and chelate binding of the cupric ion
by the sulfur at the 2 position and the oxygen at the 4
position. A consideration of the stereochemistry of
2-thiouracil and the square-planar nature of cupric ion
shows that the proposed structure is impossible because


6
(Vil)
H
the phenolic oxygen and thionyl sulfur are coplanar and
physically distant.
Cupric ion was presumed to react with thiouracil in
1,0 N NaOH in one days time to produce disulfide complexes
of cuprous ion (26). Elemental analyses were obtained on
the isolated complexes and a possible structure (VIII) was
proposed. It has been demonstrated that alkaline thioura
cil solutions are susceptible to air oxidation even in the
absence of cupric ion (16). Thus, the possibility exists
that the isolated complex contains cupric and not cuprous


7
ion as was suggested (26). To prove the existence of VIII
a solution of thiouracil disulfide in water was reacted
with cuprous chloride in concentrated hydrochloric acid
and a yellow product isolated (26). Weiss and Venner (26)
proposed, on the basis of nitrogen, copper and chloride
analyses, the substitution of chlorine for hydroxyl on
the copper in VIII. The extreme instability of thiouracil
disulfide (14) casts doubt on this assertion. It is pos
sible that the product may actually be IX or even X if the
ease of disproportionation of cuprous to cupric ion is
considered (27).
(IX)
Cu+Cl
Cl
The ease of the oxidation of 2-thiouracil by iodine
suggests that thiouracils mechanism of action may be the
reduction of iodine to the ineffective iodide. However,
since cupric ion has been implicated in thyroid function
at the level of iodine production, the removal of cupric
or cuprous ion by complexation with thiouracil could be


8
an alternate explanation for its mode of action. If the
complexation of cupric or cuprous copper is important in
the mechanism of action of the thiouracils then the sta
bility constant of the complex may be larger for copper
than with other physiological metal ions. Furthermore,
the stability constant of the copper complex may be cor
related with the biological activity of the particular
thiouracil derivative. The correlation can only be ex
pected under the conditions of equal concentrations at
the site of action. Any differences in the in vivo
solubility or stability of the thiouracil derivative
should be taken into consideration.
The principal objective of this work was to provide
quantitative information of the complexation of metal ions
with thiouracils. The types of metal ions which complex,
the effect of thiouracil substituents on the stability
constants and the structure of the complexes were de
termined in an attempt to predict substituted thiouracil
activity.


EXPERIMENTAL
I Materials
Purification of 2-thiouracil. 2-Thiouracil (Nutri
tional Biochemical Corp., Cleveland, Ohio) was recrystal
lized from hot water. The product was washed with water
and acetone and dried in a vacuum oven at 80, m.p. 322-323
dec. (all m.p. are uncorrected); literature value 310-312
dec. (28), ca. 340 (29)* Equivalent weight 130.3; calcu
lated for C^H^NgOS 128.1. I. R. spectrum (30), in cm.1
(Nujol mull): 3020 (NH); 1680 (0=0); 1280, 1240, 11??.
U. V. spectrum (28), (0.1 MHCIO^), Xmax> 273 ( 13,700),
W. 212 ( 16.600).
Purification of 6-n-propyl-2-thiouracll. The com
pound (Nutritional Biochemical Corp. and K & K Laboratories;
Plainview, New York) was recrystallized from hot water and
dried at 80, m.p. 219-221; literature value 219-221
(31), Equivalent weight 170.0; calculated for C^H^NgOS
170.2. I. R. spectrum, V in cm.1 (Nujol mull): 3100
(NH); 1650 (C=0); 1550, 1240, 1190. U. V. spectrum (0.1
MHC104), Xmax# 272 ( 15,840), XmaXi 214 (£15,840).
9


10
Purification of 6-methyl-2-thlouracil The compound
(Nutritional Biochemical Corp. and K & K Laboratories) was
recrystallized from hot water and dried in a vacuum oven
at 50. m.p. 331-332 dec.; literature value >300 (9).
Equivalent weight 142.1; calculated for C^H^NgOS 142.1.
I. R. spectrum, Qin cm.1 (Nujol mull): 3100 (NH); 1640
(C=0); 1195. 1165. U. V. spectrum (0.1 MHCIO^), Xmax#
274 (15.460), 213 ( 15.760).
Purification of 5.6-dlmethyl-2-thlouracll. This
material (K & K Laboratories) was recrystallized from hot
water and dried at 50 in a vacuum oven, m.p. 286-287
dec.; literature value 283-285 (9). Equivalent weight
156.5; calculated for C^HgNgOS 156.2. I. R. spectrum,U
in cm.1 (Nujol mull): 3210, 3110 (NH); 1660 (C=0); 1600
1210, 1130. U. V. spectrum (0.1 M HC10k), 276 (
17,340), Xmax# 215 ( 14,020).
Purification of 5-methyl-2-thlouracll. 5-Methyl-
-2-thiouracil (Sigma Chemical Co., St. Louis, Missouri)
was recrystallized from hot water, washed with water and
dried at 50 in a vacuum oven, m.p. 334 dec.; literature
value not available (32). Equivalent weight 141.2; calcu
lated for C^H^NgOS 142.1. I. R. spectrum, Din cm.1
(Nujol mull): 3090 (NH); 1640 (C=0); 1240, 1200, H65.
U. V. spectrum (0.1 M HCIO^), kmaXt 274 ( 15,450), XJBeLXu
213 ( 15,730).


11
Purification of 5-oarboethoxy-2-thlouracll. 5-Carbo-
ethoxy-2-thiouracil (Cyclo Chemical Corp., Los Angeles,
California) was used as received, m.p. 245-246; literature
value 245 (33) Equivalent weight 197*5; calculated for
C?H8N203S 200.2. U. V. spectrum (0.1 MHCIO^), Xmax# 310
(05,121), Xmax# 269 (C 8,991). A^. 213 ( 10,762).
Purification of 2-thio-6-aminouracll. 2-Thio-6-amino-
uracil was used as received, m.p. 330; literature value
295 (34). U. V. spectrum (0.1 MHCIO^), X^. 27$ (£
18,413), Amax# 202 (65,856).
Synthesis and purification of 6-methyl-N,N -dlethyl-
-2-thlouracll. This material was synthesized by the pro
cedure of Lacey (35). N.N'-diethylthiourea (032 gm.)
(Eastman Organic Chemical Co,, Rochester, New York) was
added to 20 ml. of glacial acetic acid and brought to a
boil in a round bottom flask fitted with a reflux condenser
and a dropping funnel containing 8.6 gm. of diketene (K & K
Laboratories, Plainview, New York). The diketene was added
over a half-hour period and the reaction was allowed to
cool overnight. The reaction was further heated for half
an hour and then cooled and the acetic acid removed by
vacuum evaporation. Twenty ml. of water was added to the
residue, the mixture was shaken to emulsify and put into
a refrigerator overnight. The precipitated contents were
recrystallized from hot water and dried in a vacuum oven


at 50, m.p. 97-98; literature value 97-98 (35) I R.
spectrum, Din cm.1 (Nujol mull): 1680 (C=0); 1250 1105
U. V. spectrum (HgO), lmax# 278 (6 13,100), XJnax# 222 (6
15250). Potentiometric titration with 0.1 N NaOH indi
cated no titratable acid function present. Yield is 73%,
Synthesis and purification of 2-ethylmercapto-4-hy-
droxypyrimldine. Eight gm. of 2-thiouracil (0.062 mole)
were added to 2.49 gm. NaOH (0.062 mole) in 100 ml. of
water and acetone added and the solution cooled overnight
in a refrigerator. A precipitate of the sodium salt of
2-thiouracil formed (8.3 gm., 0.055 mole) which was fil
tered and collected.
The sodium salt of 2-thiouracil (0.055 mole) and
ethyliodide (0.06 mole, 9*35 gm.) were added to 120 ml.
of 95$ ethanol in a round bottom flask and refluxed until
the sodium thiouracil had gone into solution. It was
necessary to add an additional 4 gm. (0.026 mole) of ethyl
iodide during the reaction to put the sodium thiouracil
into solution. The reaction mixture was cooled and the
ethanol removed by vacuum evaporation. A white residue
was left which was recrystallized from ethanol once and
then finally purified by sublimation, m.p. 152-153;
literature value 152 (38). I. R. spectrum, \j in cm.1
(Nujol mull): 1660 (C=0), 1270, U70, 1540. U. V. spec
trum (h2o), imax# 280 (6 5.500), xmax> 230 (6 11,750).


13
Equivalent weight 156.4; calculated for C^HgNgOS 156.2.
Anal.1 Caled, for CgHgNgOS: C, 46.13; H, 5.16; N.
17.94; S, 20.52.
Found; C, 46.40; H, 5*31; N, 18.04; S. 20.37.
Preparation and purification of disodlum salt of
2-thlouraoll disulfide monohydrate. This material was
prepared according to the procedure of Miller, et al. (14).
2-Thiouracil (2.56 gm.) was dissolved in 41 ml. of 1.0 N
NaOH and the solution cooled in an acetone-dry ice bath
until nearly frozen. Sodium iodide (4.57 gm.) and iodine
(2.54 gm.) were dissolved in 20 ml. of water and also
cooled in acetone-dry ice. The iodine-iodide solution
was added slowly, with stirring, to the alkaline 2-thioura-
cil solution in an 800 ml. beaker. About 500 ml. of ace
tone was added and the slow forming white precipitate was
collected. The precipitate was recrystallized from water
-acetone twice and dried at 70. Equivalent weight 156.6;
calculated for CgH^N^OgSgNagHjjO 158.1. U. V. spectrum (14)
(H20) Xmax. 270 ( i8,520), Xmax# 211 (£25,485); (pH 9-56
ammonia buffer) Anax# 273 (£ 12,469). Xmax>213 ( 31,537).
I. R. spectrum, in cm."1 (Nujol mull): 3250 (OH), 1570
1530, 1350, 1170, 1000, 828, 720. Molecular weight (ebul-
lioscopic in water) 96; calculated for CgH^N^OgSgNagHgO 105.
1The elemental analyses, molecular weight and weight
loss measurements in this investigation were performed by
Huffman Laboratories, Inc., Wheatrldge, Colorado.


14
Anal Caled, for CgH^N^OgSgNagHgOi C, 30.38; H, 1.91;
N, 17.71; S, 20.27; Na, 14.54.
Founds C, 29.87; H, 2.73? N, 14.38; S, 20.67; Na,
15.15.
Synthesis and purification of disodlum salt of 2-sul-
phlnyl-4-hydroxypyrlmidlne trihydrate. 2-Thiouracil (2.56
gm.) was dissolved in 26 ml. of 1.0 N NaOH and 2.1 ml. of
34# hydrogen peroxide was slowly added with stirring to
the acetone-dry ice chilled solution. Acetone was added
to the point of precipitation and the reaction mix main
tained cold overnight. The reaction mix was filtered and
the precipitate washed with acetone, the washings added
to the filtrate and the filtrate put hack into the refrig
erator overnight. A white precipitate came down which was
collected, recrystallized twice from acetone-water and
dried at 50 in a vacuum oven. Equivalent weight 257.3;
calculated for C^H2N20^SNa2OHgO 256.1. U. V. spectrum
(0.05 M HC104), Xmax> 257 ( 4,924), Xffiax# 213 ( 5.885),
Xmax. 270 ( 3.510), (0.1 MHC104), Xmax# 270 (3,631).
Anal. Caled, for C^HgNgO^SNags C, 23.8; H, 1.00;
N, 13.8; S, 15.87; Na, 22.8; weight loss for trihydrate,
21.15.
Found: C, 23.96; H, 3.16; N, 12.38; S, 16.27; Na,
21.03; weight loss, 21.53


15
II. Potentlometric Titrations
Description of titration apparatus. All potentio-
metric titrations were performed using a Sargent Model D
automatic titrator equipped with a 2.5 ml. capacity syringe
burette. The sample solutions were titrated in water-jack
eted titration cells and the temperature of each titration
was recorded. All sample solutions were flushed with
nitrogen gas, which was passed through alkaline pyrogallol
and then water to remove carbon dioxide, and a flow of the
gas into the solution during the titration was maintained.
A Beckman combination electrode with a Ag-AgCl reference
was used. The electrode was kept in water at all times
and a solution saturated with potassium chloride and silver
chloride added as needed. The pH scale of the titrator was
standardized before each titration using pH 4 and pH 7
standard buffers obtained from the Sargent Company. The
pH standardization was checked after each titration and in
no case was the pH drift larger than 0.05 pH units and
usually was less than 0.02 units.
The 1/10 normal sodium hydroxide titrant was prepared
carbonate-free by dissolving a weighed amount of sodium
hydroxide pellets in a minimum of water, filtering off the
precipitated sodium carbonate through a fritted glass fun
nel and diluting the filtrate with boiled, nitrogen purged
water. The titrant was stored in a pyrex bottle fitted


16
with ground glass connections directly to the burette on
the titrator. The titrant bottle was also fitted with an
Ascarite tube to avoid carbon dioxide absorption. The
sodium hydroxide titrant was standardized using primary
standard grade potassium acid phthalate.
Preparation of titration sample solutions. Sample
solutions for titration were prepared with constant con
centrations of thiouracil and varying concentrations of
metal ion (TablesI and II). Solution conditions for the
thiouracil-cupric ion titrations were: thiouracil deriva
tive, 0.00200 M; cupric ion, 0.00200, 0.00160, 0.00100 and
0.000400 M except for 2-thiouracil which varied from
0.00200 to 0.000200 M in 0.000200 M steps. A series of
2-thiouracil-cupric ion solutions, prepared as above, was
left standing overnight under nitrogen, filtered and a
4/5 fraction of the filtrate titrated in the usual way
(Fig. 1). No calculation of the stability constants was
possible due to the precipitation of the complex. All
thiouracil stock solutions were tested for stability by
ultraviolet analysis under working conditions of room
temperature and exposure to light. In all cases the
ultraviolet spectrum changed less than one percent over
a period of one month. Stock solutions older than one
month were discarded. The ionic strength was maintained


1?
constant at 0.006 by substituting 0.03 M sodium perchlorate
for equal volumes of 0.01 M divalent metal nitrate in ac
cordance with Eq. 1 where is the ionic strength, is
the molar concentration of the ion and is the charge
on the ion.
(Eq. 1)
The ratio of the molar concentrations of sodium perchlorate
to divalent metal nitrate is 1:3 for equal ionic strengths.
Standardization of metal ion solutions. Analytical
grade nitrate salts of cupric, cadmium, lead, nickel, fer
ric, cobalt, calcium, zinc and manganese were used to pre
pare stock solutions in distilled water. The stock solu
tions were standardized using the mercury, mercury-EDTA
electrode (Sargent Co., Chicago, Illinois) of Reilley (37,
38). The general procedure was to titrate potentiometri-
cally solutions containing 2 ml. aliquots of 0.01-0.04 M
metal nitrate, 25 0 ml. of buffer (38) and one drop of
1.0 x 10"3 M Hg-EDTA solution (38) with 0.1000 M EDTA
(EthyleneDiamineTetraAcetic acid). The reference electrode
was a saturated calomel electrode. In the case of cobalt,
an excess of standard calcium was added and the excess
determined by titration with EDTA according to the general
procedure outlined. Since Reilleys procedure cannot be


18
used to standardize ferric ion (38), the procedure of
Pribil, et al. (39) was used. The procedure was to titrate
a solution containing 5 ml of 0.01 M ferric ion, 250 ml.
of a pH 3 chloroacetic acid buffer (0.2 M) and one drop of
a 0.01 M ferrous perchlorate solution. The titrant was
0.1000 M EDTA and the electrode system consisted of a
platinum indicator electrode and a calomel reference.
Since ferrous ion is easily oxidized, fresh ferrous ion
stock solutions were prepared as needed from analytical
grade FeSO^^HgO (Matheson Coleman & Bell, Cincinnati,
Ohio) and boiled, nitrogen purged water.


19
III Ultraviolet Spectral Studies
Stability of thiouracll stock solutions. Aliquots
(2.0 ml.) of the aqueous 0.005 M stock solutions of 2-thio-
uracil, 6-n-propyl-2-thiouracil, 5.6-dimethyl-2-thiouracil,
I
2-ethylmercapto-4-hydroxypyrimidine and N,N -diethyl-6-meth-
ylthiouracil were diluted to 100 ml. with water and their
ultraviolet spectra recorded versus a water blank. Due to
the lower solubility of 5-methyl-2-thiouracil and 6-methyl-
-2-thiouracil, 0.0025 M stock solutions were prepared and
4.0 ml. aliquots were diluted to 100 ml. with water and
the spectra recorded versus a water blank. All ultraviolet
spectra were recorded on a Cary Model 15 dual beam record
ing spectrophotometer using 1 cm. matched cells. The
balancing of the cells was checked periodically by running
blank versus blank on the spectrophotometer and in all
cases no significant difference in absorbance was found.
The ultraviolet spectra of the stock solutions were fol
lowed for a period of at least one month and in all cases
the change in spectra was less than 1%.
Spectra of thlouracil-metal ion solutions. The ul
traviolet spectra of water solutions containing 2-thioura-
cil, 6-n-propyl-2-thiouracil, 6-methyl-2-thiouracil,
5-methyl-2-thiouracil, 5.6-dimethyl-2-thiouracil,
2-ethylmercapto-4-hydroxypyrimidine and 6-methyl-n,N'-di-
ethyl-2-thiouracil (1.00 x lO** M) and cupric, cadmium,


20
lead, ferrous and ferric ion (l.OOx 10^ and 7.00 x 10^)
were recorded on the Cary ultraviolet recording spectro
graph versus a water blank (Table III). Air versus air and
water versus water blank curves were run in all cases and
net absorbances were negligible. The general procedure
was to record the spectrum of the metal nitrate alone
(1.0 ml. of 0.0100 M metal nitrate diluted to 100.0 ml.),
the spectrum of the ligand alone (2.0 ml. of 0.005 M ligand
or 4.0 ml. of 0.0025 M ligand each diluted to 100.0 ml.)
and the spectrum of a solution containing 1.0 ml. of 0.01
M metal ion and 2.0 ml. of 0.005 M ligand (4.0 ml. of 0.0025
M ligand) diluted to 100.0 ml. with water (Fi^. 2-6)
Another solution containing a seven-fold excess of metal
ion over ligand was also run. Since ferric ion has a large
-4
ultraviolet absorption, the spectrum of a 7.0 x 10 M
ferric ion solution was run for comparison. Precipitates
formed when cupric ion was mixed with 6-n-propyl-2-thioura-
cil and 5*6-dimethyl-2-thiouracil, causing the absorbance
to drop below the true value. In the case of 2-thiouracil,
6-methyl-2-thiouracil and 5-niethyl-2-thiouracil no obvious
precipitate was present when cupric ion was added.
The ultraviolet spectrum of a cupric-2-thiouracil
solution as a function of hydrogen ion concentration was
also recorded. The five hydrogen ion concentrations ranged
-4
from 1.0 M to lx 10 M and the cupric ion concentration
-4 -4
was maintained at 2 x 10 M and 1 x 10 M in all cases.


21
The procedure was to add 10.0 ml. aliquots of the appro
priate HCIO^ concentration, 2.0 ml. aliquots of 0.005 M
2-thiouracil and 1.0 or 2.0 ml. aliquots of 0.01 M cupric
ion to a 100 ml. volumetric flask and dilute to the mark
with water. A solution containing HC10^ at the same con
centration as the sample was used as a blank.
Aqueous solubility of thlouraclls as a function of
hydrogen ion activity. Aqueous solutions (50 ml. in
100 ml. sealed flasks) of 2-thiouracil, 6-n-propyl-2-thi-
ouracil, 6-methyl-2-thiouracil, 5-methyl-2-thiouracil,
5,6-dimethyl-2-thiouracil, 5-carboethoxy-2-thiouracil and
6-amino-2-thiouracil were equilibrated with an excess of
solid ligand at 25.0 in a controlled temperature shaker
bath. Differing aliquots of 0.01 M NaOH and 0.1 M HCIO^
were included in the 50*0 ml. volume to vary the pH. The
pH values and the absorbances of the aliquots of the
equilibrated solutions were measured simultaneously as a
function of time to assure that equilibrium had been at
tained. Equilibrium conditions were established by repeat
ed sampling until no further change in absorbance was noted.
Filter sticks (Sargent #S-30417) were used to remove sam
ples from the equilibrated solutions. The first filtrate
was discarded and 2.0 ml. aliquots of the filtered, equi
librated solution were taken and diluted to 100.0 ml. with
0.1 M HC10^ and read on the Cary spectrophotometer


22
or Beckman DU versus a 0.1 M HCIO^ blank. Further dilu
tions were carried out as needed. The calculated solubil
ities are given in Table iv Since thiouracils are easily-
oxidized (14) the possible degradation of equilibrating
solutions of 2-thiouracil, 6-n-propyl-2-thiouracil, 5-meth-
yl-2-thiouracil and 6-methyl-2-thiouracil was checked.
Ultraviolet spectra of fresh solutions of the above com
pounds were recorded and the ratio of their absorbance to
absorbance of the equilibrated solutions calculated at
seven different wavelengths. The wavelengths chosen were
210, 220, 240, 260, 280, 300 and 320 mji. The lack of the
presence of appreciable oxidation products was indicated
by the constancy of the ratios for each curve.
Aqueous solubility of 2-thlouracil and 6-n-propyl-
-2-thlouracil as a function of pH in the presence of
cadmium ion. Sample solutions were prepared and treated
in exactly the same manner as in the previous section ex
cept that the total initial volume was carefully controlled
(50.00 ml.) and each flask (100 ml.) contained 5 ml* of
0.01 M or 25 ml. of 0.01 M cadmium nitrate. The final
total concentration of cadmium ion was 0.001 or 0.005 M.
The aliquots were suitably diluted with 0.1 M HCIO^ to
give an absorbance less than 1.00 on the Beckman DU when
measured at the
Amax


23
Determination of molar absorption coefficients (£ ).
Accurately weighed amounts of dried samples of 2-thiouracil,
6-n-propyl-2-thiouracil, 6-methyl-2-thiouracil, 5-methyl-
-2-thiouracil, 5.6-dimethyl-2-thiouracil, 2-ethylmercapto-
-4-hydroxypyrimidine, 6-methyl-N,N'-diethyl-2-thiouracil,
the disodium salt of 2-thiouracil disulfide trihydrate, the
disodium salt of 2-sulphinyl-4-hydroxypyrimidine, 5-carbo-
ethoxy-2-thiouracil and 6-amino-2-thiouracil were dissolved
(see section on materials for solvents used) in an accu
rately known volume and the spectra recorded on the Cary
versus a solvent blank. The molar absorption coefficient
(Q was calculated from Eq. 2 where A is the observed
= "Y^'c (E<1 2)
absorbance, 1 is the path length in cm. and C is the molar
concentration of the particular thiouracil. The results
are given in the section on materials.
Job's plots of 2-thlouracll and 6-n-propyl-2-thlouracil
in the presence of cupric ion. Varying aliquots of 1.00 x
10"3 M ligand were added to aliquots of 1.00 x 10 3 M
cupric nitrate so that the total number of moles of ligand
plus cupric ion were kept constant at 1.00 x 103. The
4
solutions were diluted to a final volume of 100 ml. (10
M in ligand plus cupric ion) with water in the case of
2-thiouracil and 0.05 M HC10^ in the case of 6-n-propyl-
-2-thiouracil. The ultraviolet spectrum of each solu
tion was recorded versus a water or 0.05 M acid blank,


24
respectively. The absorbance at 345 and. 272 mp (2-thioura-
cil) and 272 and 300 mp. (6-n-propyl-2-thiouracil) were
plotted versus the mole fraction of metal. A straight
line joining the absorbance at mole fraction unity to the
absorbance at mole fraction zero gives the absorbance, at
any mole fraction, that is expected under the assumption
of no interaction, i.e., no complexation (Fig. 7). The
difference in absorbance between the constructed straight
line and the experimental points is proportional to the
amount of complex formed. The concentration of complex
formed, and therefore the difference in absorbance between
constructed line and experimental curve, will be at a max
imum when the mole fraction of ligand in the test solution
corresponds to the stoichiometric fraction of ligand in
the complex. For example, a 1:1 complex would have the
largest deviation from linearity at a mole fraction of
0.5 while a 2:1 complex of ligand to metal would have the
largest deviation at a mole fraction of 0.66. Therefore,
the deviations from linearity were plotted versus mole
fraction of ligand and a smooth curve drawn through the
points (Fig. 8). If the complex absorbs at a wavelength
where the uncomplexed ligand does not, then it is not
necessary to construct a straight line and a plot of the
observed absorbances versus mole fraction will indicate
the stoichiometry as described above. This is the case
for 2-thiouracil at 345 mp. (Fig. 8).


25
In one case (2-thiouracil) the concentration of ligand
was held constant at 5*00 x 10'* molar and the cupric ion
concentration varied from 500 x 10-* M to zero. The
solutions were made up to 100 ml. with water and their
spectra recorded. The absorbances at 272 and 3^5 mjo. were
plotted versus the ratio of the molar concentrations of
cupric ion to 2-thiouracil. When the ratio of the molar
concentration of cupric ion to 2-thiouracil slightly ex
ceeds the ratio of metal ion to ligand in the complex then
the absorbance is constant for strong complexes. At ratios
of metal ion to ligand less than the stoichiometric ratio,
the observed absorbance will deviate from the constant
value (Fig. 9).
Spectrophotometric titration of 2-thlouracll. Thirty
ml. solutions of 0.921 x 10** and 1.00 x 10** M in 2-thio-
uracil and 8.00 x 10 J M in sodium perchlorate contained
in a water-jacketed titration cell (25*0) were titrated
with 1.00 N NaOH and 0.0499 N NaOH, respectively. Nitrogen
gas, free of carbon dioxide (passed through an Asearite
tube) and saturated with water (passed through a sparger
tube immersed in water) at 25*0 was passed into the cell
for each run. The cell solution was stirred with a mag
netic stirrer. A constant rate burette (Sargent Co.)
equipped with a 2.5 ml. syringe burette was used to
deliver the titrant to the cell. The volume delivered


26
could be read to the nearest 0.5 }il. A micro-aperture
flow cell (Beckman Co., catalog #97290) was connected by-
polyethylene tubing (Clay-Adams Co., #PE 200) to the
titration cell and to a 50 ml. gas-tight syringe (Hamilton
Co.). After each addition of standard alkali the gas-tight
syringe was actuated by hand to draw the titration cell
solution into the flow cell and the spectrum recorded
versus a water blank. The gas-tight syringe was actuated
several times before each spectrum was recorded to insure
thorough mixing of any solution that may have remained in
the flow cell or tubing. The pH of the solution was read
after mixing using a glass-calomel electrode system and a
Radiometer pH meter (#TTT 1). The total volume change
during the titration of the sample solution was less than
2% and was considered negligible. The absorbances at 230,
240, 270, 310, 320 and 285 m)i were plotted versus pH and
the apparent pK determined from the pH at half-neutraliza-
tion according to Eq. 3 (40) (Figs,ID, 11). The parameters
PK* = pH log A_ (Eq# 3)
a A Aalk.
f
in Eq. 3 have the following significance: the pK is the
8.
negative logarithm of the apparent dissociation constant;
the pH is the negative logarithm of the hydrogen ion activ
ity; Aacid is the absorbance of 2-thiouracil in dilute acid
(0.10 M HCIO^); A is the absorbance of 2-thiouracil solution


27
in the buffer region; Aa^k is the absorbance of 2-thioura-
cil solution in the alkaline region (pH 9)
Spectrophotometric titration of 2-thiouracil-cuprlc
ion mixtures. The procedure and apparatus used for the
titration of 2-thiouracil-cupric ion mixtures was exactly
as outlined in the previous section. The 30.0 ml. of test
A
solution was composed of 10.00 ml. of 3*00 x 10 M
-U
2-thiouracil, 2.00 ml. of 3*00 x 10 M cupric ion and
18.0 ml. of 0.008 M sodium perchlorate. Both the 2-thi-
ouracil and cupric ion solution were prepared using 0.008
M sodium perchlorate as diluent in order to maintain con
stant ionic strength. Concentrations of cupric ion
_k
0.5 x 10 and larger caused precipitation which frustrated
analysis. The recorded ultraviolet spectra (Fig.12 ) was
analyzed using Eq. 3 at 270, 285, 230, 240, 310 and 320 mji
(Fig. 13).
Ultraviolet spectra of 2-thiouracll and cupric ion
in dilute acid. The ultraviolet spectrum of a 1:50 dilu
tion of a solution containing 115 mg. of 2-thiouracil,
217 mg. of cupric nitrate trlhydrate and 20.0 ml. of 1.0
M HCIO^ diluted to 200.0 ml. was recorded versus a 0.05
M HC10^ blank. The final concentrations of 2-thiouracil,
cupric ion and perchloric acid were 9.0 x 10^, 9.0 x 10
and 0.05 molar, respectively. The 1:50 dilution was made
using 0.05 M HC10^ as diluent. Another solution was


28
prepared as above using 10.0 mg. of cupric nitrate trihy
drate, 115 mg. of 2-thiouracil and 20.0 ml. of 1,0 M HCIO^
and its ultraviolet spectrum recorded versus a 0.05 M HC10^
blank. The final concentration of 2-thiouracil, cupric ion
and perchloric acid were 9*0 x 10^, 4.3 x 10^ and 0.05
molar, respectively. The 2-thiouracil and cupric nitrate
trihydrate were weighed on a torsion balance and are accu
rate to + 5$.
Degradation of 2-thiouracil disulfide in absence and
presence of cupric ion. A solution (1.999 x 10^ M) of
the disodium salt of 2-thiouracil disulfide monohydrate
was prepared by dissolving 6.322 mg. in 100 ml. of water.
Aliquots (2.0 ml.) of the disulfide stock solution were
added to: A, 30 ml. HgO; B, 2.0 ml. HgO and 1.0 ml. of
4
4.00 x 10 M cupric nitrate; C, 1.0 ml. HgO and 2.0 ml.
, -4
of 4.00 x 10 M cupric nitrate. Each solution was pre
pared when needed and then 5*00 ml. of 0.1 M HCIO^ was
added very rapidly from a 5*00 ml. syringe (Hamilton),
mixed, and the change in absorbance recorded continuously
at 2?0 mp. The ultraviolet absorption was measured using
1 cm. cells and a 0.1 M HC10^ blank. A water jacketed
(25*0) Beckman DU, fitted with an automatic sample changer
(Gilford) and connected to a strip chart recorder (Sargent
SRL) was used to record the absorbance. The absorbance of
all three solutions was recorded until no further change
occurred


29
IV. Polarographic Studies
Apparatus. The polarographic studies were performed
on a Sargent Model XV Recording Polarograph equipped with
a water jacketed H cell (Sargent Co.) maintained at 21.0,
30.0, 350, 40.0 or 45O0. All potentials were measured
versus a saturated calomel electrode and were corrected
for potential error due to recorder current. The possible
sources of error in potential measurements using the Model
XV Polarograph are the recorder current measuring resistor,
the bridge resistance, cell resistance and the damping
resistor if used. The potential error due to recorder
current can be easily corrected for since the recorder
on the Model XV always requires 2.5 mv for full scale
deflection (41). The error in any potential reading was
corrected by simply reading the pen position at the desired
point, assuming that the scale is 0.0 to 2.5 mv and not
microamperes. The error due to the bridge resistance can
be found by multiplying the bridge resistance (200 ohms
at worst) by the observed current (1.0 to 2.0 )ia). This
gives a maximum error of 0.4 mv which is negligible. No
damping resistance was used in these studies and therefore,
potential errors due to damping do not exist. The cell
resistance was measured frequently using a Radiometer con
ductivity meter (type 2 d) operating at 300 cycles per
second. The reciprocal of the conductivity in mhos gives


30
the cell resistance in ohms. The resistance in all cases
was 300 ohms or less which is equivalent to 0.6 mv at 2 pa.
The 0.6 mv is a negligible error. The instrumental accu
racy for potential measurements, as quoted by the Sargent
Co. is +2.5 mv on a 1.0 volt scale (41). An agar plug
saturated with potassium nitrate (42) was used in the salt
bridge and was changed whenever the solid potassium nitrate
became depleted. The potassium nitrate was necessary since
the usual potassium chloride interfered durto the anodic
wave of chloride ion. All test solutions were deoxygenated
with water-saturated nitrogen gas (passed through a sparger
immersed in water) for at least ten minutes and were main
tained in a nitrogen atmosphere (nitrogen gas flowing over
the solution) during the analysis. The capillary constant
o /o *1 / zC
(m/J) t /D) was determined by running a supporting electro
lyte solution (0.2 M NaClO^) for a long, accurately mea
sured period of time (2,042.8 sec.), weighing the mercury
delivered (4.9185 gm.) and counting the number of drops
(determined from the number of pen oscillations. The val-
ues of t and m were calculated from Eqs. 4 and 5.
respectively.
(Eq. 4)
Total Sec.
(Eq. 5)


p /o a
The value of mc/J t for capillary number 2 is 2.215
The height of the mercury column was held constant at
31
58.5 cm. The current and voltage were always standardized
before each curve was run.
Polarography of cupric ion solutions. Aliquots
(3*00, 4.00 and 5*00 ml.) of a 0.0100 M cupric nitrate
solution (prepared in 0.2 M NaClO^) were added to 0.25
ml. of 0.2$ Triton X-100 (Rohm and Haas Co.) and 500 ml.
of standardized HC10^ in a 50 ml. volumetric flask and
diluted to the mark with 0.2 M NaClO^. These solutions
were purged with nitrogen and the polarograms recorded
on the Model XV at 21.0. The final perchloric acid
concentrations were 0.2648, 0.100, 0.0100, 0.00100 and
0.000100 molar. The final concentration of Triton maxi
mum suppressor (0.001$) is that recommended in the
literature (43-45) and was found to give reversible waves
as determined by Eq. 6 (46-49).
The parameters in Eq. 6 have the following significance;
E^/^ is the potential at 3A o* i^ where i^ is the observed
diffusion current in pA; Eis the potential at lA of
i^; R is the gas constant in cal./deg.-mole; T is the
absolute temperature; F is the faraday (23.060.3 cal./abs.


32
volt gram eq.); n is the number of electrons involved, in
the electrode reaction. The values of -2.303 R T log 9/F
at different temperatures are; 0.0557, (21); 0.0573 (30);
0.0583 (35); 0.0592 (40); 0.0602 (45). The derivation
of Eq. 6 is given in the section on equations.
Another series of solutions were prepared containing
1.00, 2.00, 3.00, 4.00 and 500 ml. aliquots of 0.0100 M
cupric nitrate solution (prepared in 0.2 M NaClO^), 0.5
ml. of 0.2$ Triton, 10.0 ml. aliquots of standardized HCIO^
and sufficient 0.2 M NaClO^ to dilute to 100 ml. The final
concentrations of HC10^ varied from 0.2648 to 0.000100 as
before. All solutions were purged with nitrogen gas to
remove oxygen as previously described and the polarograms
recorded on the Model XV at 21.0. All diffusion currents
were measured at the half-wave potential. The peak of the
pen oscillation was used to calculate all potentials and
currents (50) Details of the solution composition, half
-wave potentials, E^/^ Values an^- diffusion currents
are summarized in Table V.
Polarography of cupric ion-thiouracil mixtures.
Solutions containing 1.0 ml. of 0.01 M cupric ion, 0.25 ml.
of 0.2$ Triton, 5 ml. aliquots of standard HCIO^ (2.648,
2.00, 1.000, 0.800, 0.500, 0.100, 0.0100 and 0.00100 M)
and varying volumes of 0.005 M or 0.0025 M thiouracil
derivative (2-thiouracil, 6-n-propyl-2-thiouracil,


33
6-methyl-2-thiouracil, 5-dethyl-2-thiouracil, 5,6-dimethyl-
-2-thiouracil, 5-carboethoxy-2-thiouracil) were diluted to
50 ml. with 0.2 M NaClO^. The 0.005 M and 0.0025 M thi-
ouracil solutions were prepared in 0.2 M NaClO^ and as the
volume of thiouracil solution was decreased an equal volume
of 0.2 M NaClO^ was substituted to maintain the ionic
strength at 0.2 molar. The solutions were purged with
nitrogen as previously described and the polarograms re
corded at controlled temperatures (21.0, 30*0 35*0 40.0
45.0). A 1.00 ml. aliquot of 0.2 M NaClO^ was substituted
for the cupric ion in the blank solutions which were run
for each test solution. The curve for the blank was sub
tracted from the test solution curve and the half-wave
potential, E^/^-E^y^ value and the diffusion current
determined from the difference curve (Fig. 6 ). Each test
solution was run at least three times and the potential
and current measurements averaged. The peak of the pen
excursion was used to calculate all potentials and currents
(50). A typical polarogram is shown in Fig.14 and details
of typical solution composition, half-wave potentials,
E^/^-E^y^ values and diffusion currents are given in Table
VI


Polarography of solutions containing cuprous ion in
the presence of 2-thlouracll and 2-thiouracil disulfide.
Cuprous chloride was purified by dissolution of 5 gm. of
CuCl (light green in color) in 100 ml. of concentrated
hydrochloric acid to give a black solution containing
CuCl~. Addition of 300 ml. of water gave a light green
solution from which white crystals of pure CuCl precipi
tated (51) The following qualitative tests were run on
the precipitated CuCl. A solution of ammonium hydroxide
(28$) added to the CuCl gave a colorless solution which
+2
slowly turned dark blue (Cu(NH^)2 ). Concentrated HC1
added to the CuCl gave a colorless solution (CuClg) which
rapidly turned dark. Addition of potassium iodide to the
colorless CuClg solution produced a white precipitate
(Cul).
Polarograms of the following solutions were run as
previously described. Solution A contained 1.979 mg. CuCl
1.0 ml. of 0.2 M NaClO^, 0.25 ml. of 0.2$ Triton, 5*0 ml.
of 2.648 M HCIO^ and 43.75 ml. of 0.005 M 2-thiouracil
(in 0.2 M NaClO^). Solution B contained 1.55^ mg. 2-thi-
ouracil disulfide, 0.25 ml. of 0.2$ Triton, 5*0 ml. of
2.648 M HC104 and 44.75 ml. of 0.2 M NaClO^. Solution C
contained 1.979 mg. CuCl, 3*148 mg. 2-thiouracil disulfide
10 ml. of 2.648 M HCIO^, 0.5 ml. of 0.2$ Triton and 89.5
ml. of 0.2 M NaClO^.


35
Polarography of presumed cuprous complex of 2-thioura
cil. The cuprous chloride complex of 2-thiouracil (IX)
was synthesized according to the literature (26). The
procedure for the synthesis is the dissolution of 1.3 gm.
of 2-thiouracil in 150 ml. of hot water and 40 ml. of 1.0
M CuCl in concentrated HC1 solution was added. Yellow
crystals formed which were washed with water and then ace
tone and dried in a vacuum over at 50
Two solutions of IX were prepared. The first solution
contained I.883 mg. IX, 0.25 ml 0.2$ Triton, 50 ml. of
3.0 M HCIO^ and was diluted to 50 ml. with 0.2 M NaClO^.
The second solution contained I.876 mg. IX, 0.25 ml. 0.2$
Triton, 50 ml. of 3.0 M HCIO^ 14.75 ml. of 0.005 M
2-thiouracil (in 0.2 M NaClO^) and was diluted to 50.0
ml. with 0.2 M NaClO^. Both solutions were run on the
polarograph at 21.0. The solutions were purged with
nitrogen as usual and a blank solution containing every
thing but IX was run in both cases.
V. Synthesis of Complexes
Synthesis of bis(2-thiouracll)cadmlum(II). A solu
tion containing 0.06 mole of 2-thiouracil in 2 liters of
hot water was prepared. To the 2-thiouracil solution was
added slowly, with stirring, a solution containing 0.03
mole of cadmium nitrate in about 200 ml. of water. The
resulting mixture was allowed to stand 0.5 hour on low


36
heat (about ?0). The solution was allowed to cool to room
temperature and the pH adjusted to 6.5 with concentrated
NaOH. The resulting suspension was warmed again, cooled
to room temperature and filtered through a medium glass
fritted funnel. The product was washed with cold water,
acetone and dried in a vacuum oven at 60. I. R. spectrum,
Din cm.""1 (Nujol mull): 1620, 1550, 1530, 1300, 1210,
1180, 82?.
Anal. Caled, for CgHgN^OgSgCd: Cd, 30.65.
Found: Cd, 31.54, 30.96.
Synthesis of bis(2-thiouracll)lead(II). The pro
cedure for the synthesis of bis(2-thiouracil)lead(II) is
exactly the same as in the case of bis(2-thiouracil)cad-
mium(II) except, that lead nitrate was used in place of
cadmium nitrate. I. R. spectrum, Din cm.~^ (Nujol mull):
1660, 1630, 1560, 1500, 1280, 1000, 815.
Anal. Caled, for CgHgN^OgSgPb: Pb, 44.90.
Found: Pb, 44.51.
Synthesis of bis(2-thlouracil-cadmlum(II)). A solu
tion containing 0.06 mole of 2-thiouracil in a minimum of
hot water was added slowly, with stirring, to a one liter
solution containing 0.06 mole of cadmium nitrate heated
to the same temperature (about 80) as the hot thiouracil
solution. The resulting solution, which may contain a
suspension of product, was maintained at the initial


37
temperature for ^ hour and then allowed to cool to room
temperature. Concentrated sodium hydroxide was added to
pH 6.8 and the resulting suspension reheated. After \
hour the suspension was cooled to room temperature and
the product filtered through a medium glass fritted flannel.
The resulting product was washed with water and then ace
tone and dried in the vacuum oven at 60. I. R. spectrum,
Din cm."1 (Nujol mull): 3400, 1570, 1510, 1335, 1020.
Anal. Caled, for CgH^N^OgSgCdg: Cd, 4?.2.
Found: Cd, 45.8.
Synthesis of bis(2-thiouracll-lead(II)). The pro
cedure for bis(2-thiouracil-lead(II)) is exactly the same
as in the case of bis(2-thiouracil-cadmium(II)) with lead
nitrate substituted for cadmium nitrate. I. R. spectrum,
0 in cm."1 (Nujol mull): 1560, 1520, 1430, 1330, 1000,
820.
Anal. Caled, for CgH^N^OgSgPb: 62.15*
Found: Pb, 62.54.
Synthesis of bis(6-n-propyl-2-thlouracll)cadmium(II).
The preparation of this complex was performed in the
same manner as for bis(2-thiouracil)cadmium(II). The molar
amounts of 6-n-propyl-2-thiouracil and cadmium nitrate were
0.06 and 0.03, respectively. I. R. spectrum, Uin cm."1
(Nujol mull): 3100, 1630, 1500, 1270, 1220, 11?5, 1015,
970, 830.


38
Anal. Caled, for C^H^gN^OgSgCd: Cd, 24.9; C, 37 29;
H, 4.02; N,12.43; S, 14.22.
Found: Cd, 24.7; C, 37.94; H, 4.15; N, 11.95; S,
13.77.
Synthesis of bis(6-n-propyl-2-thiouracll-lead(II)).
This complex was prepared by the same procedure as for
bis(2-thiouracil-cadmium(II)). The molar amounts of
6-n-propyl-2-thiouracil and lead nitrate used were 0.06.
I. R. spectrum, Din cm.*"^ (Nujol mull): 1550. 1420, 1280,
1165, 1020, 825.
Anal. Caled, for C^H^N^O^Pb: Pb, 55*2.
Found: Pb, 55*4.
Synthesis of bis (2-thlouracll )-p-dlhydroxodicopper (II).
A solution containing 0.062 mole of cupric ion in 100 ml.
of water was added slowly, with stirring, to 400 ml. of a
warm, aqueous solution (about 1.0 liter) containing 0.062
mole of 2-thiouracil. Very fine, light-yellow crystals
slowly formed which were removed by filtration, washed
with water and acetone and dried at 120 overnight. The
elemental analysis was done on material dried to constant
weight. Ultraviolet assay of 2-thiouracil content gave
62.2$. Calculated value for CgHgO^N^SgCUg is 61.7$. The
procedure for the assay was to put an accurately weighed
amount of complex into hot 1.0 M HC10^ for one day and
then read the ultraviolet spectrum versus a 1.0 M HCIO^


39
blank. The concentration of 2-thiouracil was determined
from the molar absorption coefficient (£ 13,700) at 273
m}i. I. R. spectrum, Din cm.~^ (Nujol mull): 3090, 1640,
1600, 1540, 1280, 1160, 1070, 1015, 825. Weight loss when
dried at 100 under vacuum for two weeks was 4.57$. Cal
culated for CgHgO^S^Ug is 4.34^.
Anal. Caled, for CgHgN^OgS^Ug: C, 24.18; H, 1.52;
N, 14.10; S, 16.14; Cu, 31*98.
Found: C, 24.54, 24.06; H, 1.62, 1.66; N, 14.33;
S, 19.48, 17.51; Cu, 32.84.
Synthesis of bis(6-n-propyl-2-thlouracll)-p-oxodl-
copper(II). The procedure for the preparation of bis(6-
-n-propyl-2-thiouracil )-ji-oxodicopper (II) is exactly the
same as for the 2-thiouracil case. The yellow precipitate
was dried to constant weight at 120. The procedure for
the spectrophotometric analysis of the percent of 6-n-pro-
pyl-2-thiouracil is the same as for 2-thiouracil. I. R.
spectrum. Din cm.-* (Nujol mull): 3040, 1640, 1540, 1490,
1440, 1270, 1210, 1170, 1020, 955, 835.
Anal. Caled, for C^H^OgN^SgCUg; Cu, 26.39; 6-n-
-propyl-2-thiouracil, 70.7*
Found: Cu, 26.82; 6-n-propyl-2-thiouracil, 70.6.


40
Analytical procedure for cadmium content of complexes,
Accurately weighed (200 mg.) samples of the dried cad
mium complex were dissolved in 150 ml. of 0.01 M HgSO^ and
heated (70) until all of the sample had dissolved. Ap
proximately 150 mg. of sodium hydrogen sulfide in 15 ml.
of water was added and the resulting cadmium sulfide pre
cipitate digested for about two hours until the crystals
were large. The precipitate was filtered onto a tared
glass fritted funnel, washed with warm water and dried
in a vacuum oven at 5 overnight. The dried precipitate
was weighed in the funnel and the weight of cadmium sulfide
determined by subtraction. The weight of cadmium was cal
culated by multiplying the weight of cadmium sulfide by
the gravimetric factor O.778O. The percent cadmium was
calculated from the ratio of cadmium weight to sample weight
multiplied by 100. This analytical procedure is based on
the literature method (52).
Analytical procedure for lead content of complexes.
Accurately weighed samples (300 mg.) of the lead complex
was digested in 100 ml. of 1.0 M nitric acid until every
thing had dissolved. The volume was reduced to 25 ml by
evaporation and 100 ml. of a solution 2.0 M in sulfuric
acid and 1.0 M in sodium sulfate was added. The precipi
tate was digested to give large crystals and the volume
reduced by evaporation to about 50 ml. The precipitate


41
was collected on a tared, fritted funnel, washed with water
and dried at 120 overnight. The weight of lead was deter
mined as in the cadmium case. The gravimetric factor is
0.6832. This analytical procedure is based on the litera
ture method (53)
VI. Animal Experiments
Determination of antithyroid activity of 5.6-dlmethyl-
-2-thiouracll. Twenty white Carworth CFN rats were evenly
divided into four groups with five rats per group. One
group was used as a control and was fed Rockland Rat diet.
The other three groups were fed the same diet containing
0.025 0.050 and 0.100 weight-percent 5.6-dimethyl-2-thio-
uracil for a period of two weeks. Both diet and water
were supplied ad libitum. The doses were chosen with
the hope that a dose response relationship between drug
dose and thyroid weight could be obtained. Body weights
were measured just before sacrifice. At death weights
of thyroid glands were determined on a torsion balance. The
total amount of diet ingested was not measured but an esti
mate of 3*86 gm. of diet per 100 gm. of body weight per day
was made from the 5-Methyl-2-Thiouracil study. The effect
of the drug on thyroid weights is given in the section on
results and the thyroid weights are given in the results
and in Table VII


42
Determination of antithyroid activity of 6-methyl-
t
-N,N -dlethyl-2-thlouracil. Ten white, male Carworth
CFN rats (120-200 gm.) were divided into two groups of
five rats each. One group was used as a control and was
given the standard diet. The other group was given the
Rockland Rat diet containing 0.100$ 6-methyl-N,N -diethyl-
-2-thiouracil. The amount of drug ingested per day was
determined by weighing the diet supplied each day and
then weighing the amount uneaten at the next feeding
period. The rats were given the diets for a period of
two weeks and both diet and water were supplied daily
as needed. At the end of two weeks the animals were
sacrificed and the thyroid glands carefully removed and
weighed on a torsion balance. The thyroid weights are
given in the section on results and in Table VIII.
Determination of antithyroid activity of 6-n-propyl-
-2-thlouracil and 5-methyl-2-thiouracll. The procedure
for the determination of the antithyroid activity of 6-n-
-propyl-2-thiouracil and 5-.ethyl-2-thiouracil is exactly
the same as for 5*6-dimethyl-2-thiouracil except the


43
daily consumption of diet per 100 gm. of rat body-weight
was measured. The thyroid weights and consumption of
diet are given in Table IX.
VII. Magnetic Susceptibility Measurements

The corrected molar magnetic susceptibility,Xof
the cuprous-2-thiouracil complex was measured by the Chem
istry Department, University of Florida and gave a value
_L
of -1.10 x 10 a negative number. Since ji, the magnetic
dipole moment, is proportional to the square root ofX^, a
negative number, then the magnetic dipole moment is most
probably zero and the isolated complex is probably the
cuprous complex of 2-thiouracil. The absolute value of
X^ is rather large.


EQUATIONS
I Potentiometrlc Titrations
Derivation of equations for stabilillty constants
assuming the presence of only MU+ and MUp. The formation
of 1:1 and 2:1 complexes could occur as shown In Scheme II
for our specific compounds.


45
The derivation of the general equation for the calculation
of stability constants from potentiometric titration data
was first given by Bjerrum (54). An excellent summary of
computational techniques is also available (43). The
following derivations do not assume the formation of the
hydroxides of the free metal ion and require all species
to be in solution and at instantaneous equilibrium. The
highest pH at which usable data could be obtained in our
studies was about 6.5 since the precipitation which occurred
destroyed the equilibrium conditions. Hydrolysis constants
from the literature (55) permitted the estimation of values
for maximum hydrolysis of cadmium (0.2%) and lead (35$) at
pH 6.5. At pH values above 6.5 the loss of free cadmium
and lead could not be considered negligible. The apparent
acid dissociation constant of the thiouracils is:
K- ir] [H+j (Eq. 7)
a [HU]
where [u] and [HU] are the molar concentrations of
thiouracil anion and free acid respectively, [H+] ¡f + is
the hydrogen ion activity, and where ^+ is the mean activ
ity coefficient. The stability constant for the first
complex of metal ion with thiouracil is:
1 [K^2] [U-]
(Eq. 8)


46
where [Ml/] is the concentration of the first complex and
[M+2] is the concentration of free metal ion. The step
stability constant for the formation of the 1:2 complex
is*.
K = [MU2]
[MU+] [IT]
(Eq. 9)
where [MU2] is the concentration of the second complex.
The overall stability constant, the product of and K2,
is:
[MUj
fB o = K. K = r TT
2 1 2 [m+2] [u-]2
(Eq. 10)
The mass balance equations for thiouracil, metal ion and
sodium hydroxide titrant are:
[HU]0 = [U-] + [HU] + [MU+] + 2[MU2] (Eq. 11)
[M+2]0 = [M+2] + [MU+] + [MU2] (Eq. 12)
[NaOH] = [U~] + [MU+] + 2[MU2] + [OH"] [H+] (Eq. 13)
The initial stoichiometric concentrations of thiouracil and
metal ion are given by [hu]q and [m+2]q. The stoichio
metric concentration of alkali at any point in the titra
tion is given by [NaOH]. The concentration of hydroxyl
ion in Eq. 13 is the sum of the hydroxyl ion from the


47
titrant and from the dissociation of water. Since a hydro
gen ion is produced when a water molecule dissociates, the
hydrogen ion concentration corrects the hydroxyl ion con
centration for this phenomenon so that the resultant equa
tion accounts for the hydroxyl ion due to the titrant.
The degree of formation, is defined as the average
number of ligands bound to a metal ion.
[MU+] + 2[MUj
=
[m+2]0
(Eq. 14)
When the appropriately rearranged Eqs. 7 and 11 are sub
stituted into Eq. 14,
n
(Eq. 15)
When Eq. 13 is subtracted from Eq. 11, substitution of the
rearranged Eq. 7 for [HU] where the relatively small quan
tities [0H~] and [H+] are ignored,
[HU]0 [NaOH]
[H+] X/\
[u-] =
(Eq. 16 )


48
The right hand side of Eq. 16 contains only experimental
quantities and therefore, [U-] can be calculated. Sub
stitution of [u-] into Eq. 15 allows the calculation of
.
Substitution of the rearranged equilibrium expressions
for £MU+], [MUg] (Eqs. 8 and 10) and the mass balance (Eq.
12) for [M+2]0 into Eq. 14 gives, on simplification and
rearrangement, a relation between and 02,
n
d-)[u]
^-j e2[u-] + K:
(Eq. 17)
Equation 17 is linear with a slope of 02 and an Intercept
of If 02 is assumed to be zero Eq. 17 reduces to an
equation whose logarithmic transformation is:
log 1-n = pK. + p[lT] (Eq. 18)

where pK^ and p[U-] represent the negative logarithm of
and [U-], respectively. Equation 18 is a linear equa
tion with a slope of one and an intercept of pK^.
Derivation of equations for stability constants as
suming the presence of only MU+ and MUOH. The formation
of MUOH, the mixed ligand complex (XV), can conceivably
occur by reaction of XIII with 0H~.


^9
XV
OH
MU OH
This derivation of the equations for the calculation of
the stability constants of XV assumes that no significant
amounts of MOH+ and/or MUg are formed and all species are
in solution and at equilibrium.
The expressions for the apparent acid dissociation
t
constant, K and the equilibrium constant, K1 for the
a i.
first complex, MU+, have already been given (Eqs. 7, 8).
i
The step stability constant, Kg, of the mixed ligand
complex is:
, [mu oh]
Kg = (Eq. 19)
[MU+] [OH-] y +
where [mUOK] is the molar concentration of the mixed ligand
complex and [0H~] is the activity of hydroxyl ions. The
mean activity coefficient is given by Y+. The expression


50
for the autoprotolytic constant of water is:
Kw = [H+]J + [OH-]
(Eq. 20)
The overall stability constant, the product of Kw, and
f
Kg, for the mixed ligand complex is:
[MUOH] [H+]
[M+] [u]
(Eq. 21)
The value of [0H~] y+ has been substituted by Kw/[H+]
from Eq. 20. The mass balance equations for thiouracil,
metal ion and sodium hydroxide titrant are:
[HU]0 = [U] + [HU] + [MU+] + [MUOH] (Eq. 22)
[M+2]0 = [M+2] + [MU+] + [MUOH]
,+2
(Eq. 23)
[NaOH] = [U~] + [MU+] + 2[MU0H] + [0H] [H+] (Eq. 24)
The degree of formation for the mixed lagand complex is:
[MU+] + [MUOH]
n =
[m+2]0
(Eq. 25)


51
Substitution of rearranged Eqs.
suits in:
[HU]0 [U-] 1 +
=
7 and 22 into Eq.
[H+] Yt
t
0
25 re-
(Eq. 15)
Subtraction of Eq. 24 from Eq. 22, substitution of [HU]
from Eq. 7, [MUCH] from Eq. 21 and dropping the relatively
small quantities [0H] and [H+] gives:
[HU]0 [NaOH]
= [U~] (Eq. 26)
[H+]^+/k¡ 3U[M+2]/[H+]y+
The left hand side of Eq. 26 is a function of the free
metal ion concentration and therefore, the value of [U]
cannot be accurately calculated unless the free metal ion
concentration is determinable. If the assumption is made
that MUOH is not present in any significant amount during
some interval in the titration then Eq. 26 reduces to Eq.
16. This assumption would permit the observed data to
conform to Eq. 18.
The relation of to and 0^ for mixed ligand com
plexes is derived by substitution of the equilibrium ex
pressions for [MU+] (Eq. 8), [MUOH] (Eq. 21) and [M+2]
(Eq. 23) into Eq. 25 to give:


[p~]
[H+]
52
KjtU'] + Bn
n =
1 + KX[U] + en[u"]/[H+] x+
(Eq. 2?)
Rearrangement of Eq. 27 by multiplication of n by the
denominator and collection of similar terms gives:
n = K + (Eq. 28)
(1 5)[ir] 1 [H+]ir+
which is linear with a slope of g^ and an intercept of
. When it is assumed that the concentration of MUOH
is not significant, g^ approaches zero and the logarithmic
transformation of Eq. 28 reduces to Eq. 18.
Derivation of equations for stability constants as
suming the presence of MU+, MU and MgU^. The formation
of polynuclear complexes (MnUn) can conceivably occur as
shown in Scheme III#


53
XVI MU


54
Molecular models of XVII are easily formed and exhibit no
strain. It is also possible to rewrite XVII in the linear
form
XVIII KnUn
The assumptions for the derivation of the equations for
polynuclear complexes are the same as for simple and mixed
ligand complexes, that is, no significant amounts of MOH+,
MUg and/or MUOH are formed and all species are in solution
and at equilibrium.
The expressions for the apparent acid dissociation
constant of the ligand and the stability constant of the
first complex (MU+, XIII) have been given (Eqs. ?, 8). The
expressions for the acid dissociation constant of MU+ (XIII)
and the step stability constant of MgUg (^VII) are
. [MU] [H+] jfc
a2 [MU+]
(Eq. 29)


55
(Eq. 30)
Substitution of the equilibrium expression for [MU+]
(Eq. 8) into Eq. 29 and. collection of constants gives
Ka2Kl
[MU] [H+] X
[M+2] [IT]
(Eq. 31)
Substitution of the expression for [mu] from Eq. 31 into
Eq. 30 and collection of constants gives
P
22 =
[m2u2] [h+]22T+2
[m+2]2 [u]2
(Eq. 32)
The mass balance equations for thiouracil, metal ion and
sodium hydroxide titrant are
[HU]0 = [U-] + [HU] + [MU+] + [MU] + 2[M2U2] (Eq. 33)
[M+2]q = [M+2] + [MU+] + [MU] + 2[M2U2] (Eq. 34)
[NaOH] = [U] + [MU+] + 2[MU] + 4[M2U2] + [0H~] [H+]
(Eq. 35)
The degree of formation for the specific polynuclear com
plex M2U2 is
[MU+] + [mu] + 2[M2U2]
[+2]0
n
(Eq. 36)


56
Equation 36, when substituted by rearranged Eqs. 7 and 33
(Eq. 15)
[*+2]0
Subtraction of Eq. 35 from Eq. 33. substitution of Eq. 7,
31 and 32 for [HU], [mu] and [M2U2], respectively, and
dropping the relatively small quantities [CH~] and [H+]
gives
[HU]0 [O']
1 +
[H+]¡f +
K
n =
[HU]q [NaOH]
[H+] X+fii'a 3a2[M+2]/tH+] X 2g22[M+2]2[U-]/tH+]2^2
(Eq. 37)
The general equations for the ligand concentration in the
case of polynuclear complexes M^U^, where p must equal q
as in Scheme III, would be given by
[u-] =
[HU]n [NaOH]
chN
0 1
P gpq[M+2]q[U~]
-IP-1
[H+]P^+P
(Eq. 38)
When p = q = 1 Eq. 38 reduces to Eq. 26 and when p = 0
it reduces to Eq. 16. The left hand sides of Eqs, 37 and
38 are functions of the free metal ion concentration and
therefore, the value of U- cannot be accurately calculated


57
unless the free metal Ion concentration is known. If the
assumption is made that M^Up is not significant during
some interval in the titration then Eqs. 37 and 38 reduce
to Eq. 16 and permits the data to be plotted according to
Eq. 18.
The relation of to Kq, ga2 and @22 is derived by-
substitution of the equilibrium expressions for [MU+]
(Eq. 8), [MU] (Eq. 3D, [MgUg] (Eq* 32) and [M+2]0 into
Eq. 36 to give
KjU"] + 3a2[U~][H+]~1 y+"1 + 2g22[M+2][U~]2[H+]~2
1 + KX[U] + 3a2[U"][H+]"1 T+"1 + 2P22[M+2][U"]2[H+]"2 r+"2
(Eq. 39)
The general equation for in the case of MU+ in the pres
ence of MU is
q. P
Q P
^[U-] + 21 Y- pl3qp[M+2]q~1[UHH+]P'"P
1 1
Q P
Ki[u"] + Y_ q3qp[M+2]q"1[u"]p[H+]"p2r'p
' 1 0
(Eq. 40)
, rs ,
The presence of [M+ ] and [H j in Eq. 40 requires that the
free metal ion concentration be known for calculation of
gqp and Ki (56).


58
II. Polarography
Derivation of equation for polarographlc analysis of
copper complexes of thlouraclls. Kolthoff and Lingane
(57) have given derivations of the equations to be used
for the calculation of stability constants of complexes
from polarographic data. The equation for the reduction
of cupric-thiouracil complex to the cuprous complex at the
dropping mercury electrode (DME) is
Cu(U) + e" ^ CuU + U" (Eq. 41)
The anion of a thiouracil molecule is represented by U.
The difference between the polarographic half-wave poten
tials in the partial reduction of a simple (uncomplexed)
metal ion and in the reduction of a complex ion of the
same metal can be derived if we consider that the reaction,
for the purpose of discussion, occurs in two hypothetical
steps according to Eq. 41 and Eq. 42.
Cu+2 + e ^ Cu+ (Eq. 42)
The standard potentials for Eqs. 41 and 42 are E and
E, respectively. The overall reaction can be written
S
by subtraction of Eq. 42 from 41 since both are one
electron reductions.


59
CuU0 + Cu+ ^ CuU + U~ + Cu+2 (Eq. 43)
'C
The standard potential of Eq. 43 is the difference between
the standard potentials of Eqs.4l and 42.
E = E E (Eq. 44)
C s
The equilibrium constant for Eq. 43 is
[CuU] [U~] [Cu+2]
[Cu+] [CuU2]
(Eq. 45)
The dissociation constants of the oxidized and reduced
complexes are given, respectively, by
[Cu+2] [U"]2
[CuU2]
[Cu+] [U~]
[ CuU]
(Eq. 46)
(Eq. 4?)
The ratio of K to K is identical to K Since the free
o r ov
energy of a reaction is proportional to its potential (58)
we can write the equivalent equilibrium expression relating
potential and equilibrium constant.
RT
nF
ov
RT
nF
In
(Eq. 48)


60
The symbols in Eq. 48 have the following meanings: n,
electron change in the reaction; F, the faraday; R, the
molar gas constant; T, the absolute temperature. The
Nernst equations (58) for Eq*. 42 and 43, respectively,
are
[Cu+]*
T? F RT
^DME ~ W
In
[Cu+
2 ]jr
(Eq. 49)
f f RT
DME c nF
In
[CuU]0^
[CuU.]0^'
- -TT~ in [D-] t
(Eq. 50)
The values of P and Q represent the stoichiometric number
of ligands bound to the oxidized and reduced metal,
respectively, and is unity for the specific case of Eq.
43. The potential of the dropping mercury electrode is
represented by E^g. The concentrations of oxidized and
reduced species appearing in Ecp* 49 and 50 are those at
the electrode surface. The activity coefficients for the
oxidized and reduced states of the simple and complex ions
are represented by %Q, ^ and C K- respectively. Be
cause of the low concentration of reactive species at the
electrode surface the ratio of K/ *0and K' K may be
taken as unity and the activity of the thiouracil anion
is assumed to be equal to its concentration (59)


61
The relation of the concentration of the oxidized and
reduced species in the bulk solution may be derived in the
following manner. Since the initial concentration of re
duced species is zero at the electrode surface we can write
[Cu+] = i/kr
(Eq. 51)
which relates the electrode surface concentration to the
current at any point on the curve. The parameters in Eq.
51 have the following significance; i, the current in
microamps; [Cu+], the concentration of cuprous ion at
the electrode surface; kr, the proportionality constant
between current and concentration. The proportionality
factor, kr, is directly proportional to the square root
of the diffusion constant and comes from the Ilkovic
equation (60).
(Eq. 52)
The Ilkovic equation, which has been experimentally proven
(61), assumes that the current at any point on the current
-potential curve is directly proportional to the difference
between the concentration of the active species at the
electrode surface and the bulk solution (62). The para
meters in the Ilkovic equation have the following signifi
cance: i, current in microamperes at any point on the


62
curve; n, number of electrons Involved in the electrode
reaction; C, concentration of active species in millimoles
per liter; m, weight of mercury flowing from DME in mg.
sec. ; D, diffusion coefficient of active species; t,
drop time in seconds; 706, a constant equal to 4t7tt/3 F
where F is the faraday (96,50 coulombs). The proportion
ality between current and the concentration gradient from
bulk solution to electrode surface is expressed by
i = kQ([Cu+2] [Cu+2]0) (Eq. 53)
where kQ is the proportionality constant from the Ilkovic
equation and [Cu+2] and [Cu+2] are the bulk and electrode
surface concentrations of cupric ion, respectively. The
current will reach a maximum when the active species is
reduced as fast as it can diffuse to the electrode surface
causing the concentration at the surface to fall to zero.
This condition is expressed in Eq. 5^ by setting [Cu+2]^
in Eq. 53 equal to zero.
iD = kQ[cu+2] (Eq. 54)
Substitution of Eqs.51t 53 and 54 into Eq. 49 gives
T? T^o RT ko RT i
eDME Es W~ ln ~Tc F" ln i 55)
r D


63
When i = 1^/2, Eq. 55 becomes
r
Substitution of Eq. 56 into Eq. 55 gives
EDME = P~ ln in-i ^Eq*
The equations relating the diffusion current (i^) to the
bulk concentrations of oxidized and reduced complex are
CD) <[CUU2]
(Eq. 58)
-v; = <[cuu]
(Eq. 59)
Cathodic and anodic currents are designated by (in) and
D C
(in) respectively, and primes" are used for parameters
derived from complexes. The significance of the other
terms has already been mentioned. The electrode surface
concentration of the reduced comples, [CuU]^, is the sum
of its bulk concentration and any that is formed from re
duction of the oxidized complex.
i
[CuU] = [CuU] +
(Eq. 60)


64
In the particular case of Eq. 43 the bulk concentration of
reduced complex is negligible. The electrode surface con
centration of the oxidized complex is the difference be
tween its bulk concentration and the amount lost by reduc
tion at the electrode.
[CuU?] = [CuU2]
k
(Eq. 61)
o
Equations 60 and 61 are obtained by a rearrangement of
equations having the same form as Eq. 53 but written in
terms of the reduced and oxidized complex, respectively.
Substitution of Eqs. 58. 59. 60 and 6l into Eq. 50 and
collection of terms gives
k
k
r
o
(Eq. 62)
If [CuU] = 0 then -(i^)^ is also zero from Eq. 59 and Eq
62 becomes
k
In [U-]
o
k
r
(Eq. 63)


65
When i = i^/2, then E^g = (E|)c anc^ Eq. 63 is
c = E -
HT
nF
k
m -3- ^ -f- m [iT]
k
(Eq. 64)
Subtraction of Eq. 56 from Eq. 64 yields
kokr
k k'
o r
(Eq. 65)
which relates the change in the half-wave potential to the
change in the standard potential of a metal ion when it is
complexed.
The diffusion coefficient of a free cuprous ion is not
readily measurable because of its ease of disproportiona
tion to cupric ion and copper metal. Since the ionic radii
of cuprous and sodium ion are O.96 and 0.95 angstroms,
respectively, we can make the approximation that the charge
densities of the two ions are very similar. Because of the
similarity in charge densities the probable size of the
hydration spheres and hence the probable diffusion coef
ficients of the two ions would be expected to be the same,
i.e., 1.35 x 10 J cm. sec. as given for sodium ion (63).


66
Since cuprous ion is heavier than sodium ion however, the
value of may be overestimated. On the probably unwar
ranted assumption of the validity of the application of
the Stokes-Einstein Law to ionic species, it may be pre
dicted that the ratio of the diffusion coefficients of
cuprous to sodium ion is equal to the inverse ratio of
the cube roots of their atomic weights, i.e., (22.9) "V
(63*5 (64). This calculation gives a value of 0.96
x 10"5 for the diffusion coefficient of cuprous ion. Since
kr and kQ are proportionality constants which relate cur
rent and concentration, consideration of the Ilkovic equa
tion (Eq. 52), shows that ratios of the square roots of
the diffusion coefficients may be substituted for ratios
of the proportionality constants. The ratio of the square
roots of the diffusion coefficients of cuprous and cupric
ions (kr/kQ), using the value of 1.35 x 10^ for cuprous
ion, is 1.3 when the diffusion coefficient of cupric ion
is 0.72 x 10-5 cm.2 sec."^ (63). The ratio of the square
roots of the diffusion coefficients of the oxidized and
reduced complexes, (k^/k^), as estimated from the Stokes
-Einstein Law and their molecular weights (64) is approx
imately 0.92. Therefore, the calculated value of
RT/nF(ln kQkr/krko) is 4.6 mV which is approximately
experimental error. The true value will be less than
4.6 mV due to the overestimation of k^ as 1.35 x 10-^
since the hypothetical calculation given above indicates


67
that kr may be as low as O.96 x 10^. Therefore, the last
term on the right hand side of Eq. 65 may be considered
negligible and ignored. Substitution of Eq. 48 and Eq.
7 into Eq. 65 after dropping the last term on the right
gives
(Vo- In
Ko
*+l
11
+
w
1 1
Kr
/
< 1
(,p.-S) JgL in [HU]
T~
(Eq. 66 )
Eq. 66 is linear with a slope of (P-Q) RT/nF and an inter
cept from which the ratio of the stability constants can
be evaluated.
The value of (E|)g for the reduction of cupric to
cuprous ion is not experimentally determinable because
cuprous ion is reduced easier than cupric ion and a two
electron change always occurs. The half-wave potential
at any temperature can be calculated from Eq. 56 if the
values of E and k /k are known at the temperature in
sor
question. The temperature effect on diffusion coefficients
is only 2% per degree centigrade (65), is negligible and
kQ/kr may be considered constant for small temperature
changes. The standard potential for the reduction of
cupric to cuprous ion, as a function of temperature, can
be calculated from the standard free energy, entropy and
enthalpy at 25 using


68
i H2 + Cu+2 > H+ + Cu+ (Eq. 67)
AP = AH TAS = -nFE (Eq. 68)
assuming that the enthalpy and entropy are constant over
the temperature range in question. Sample calculations
and the required thermodynamic values are available in
the literature (66). The potential of the saturated calo
mel electrode as a function of temperature is also avail
able in the literature (67.68). The respective calculated
values of E and (EA) (Eq. $6) for the reduction of cupric
to cuprous ion at various temperatures are: 21, 0.1512,
-0.0878; 250, 0.1514, -0.0837; 35. 0.1521, -0.0806; 40,
0.1526, -0.0755; 45, 0.1530, -O.O744. The value of (E)g
at 25 agrees very well with the value calculated in the
literature (69), i.e., -0.079V., even though the literature
value was not corrected for the difference in diffusion
coefficients of cupric and cuprous ions. The literature
value of the half-wave potential for the reduction of
cuprous to free copper is +0.143 V versus the saturated
calomel electrode at 18.0 (70) so that, at potentials
more negative than +0.143 V, a two electron reduction
would be expected if free cuprous ion is produced by the
reduction of CuU


69
Derivation of equation for values. Sub
stitution of 1 = 0.75 Into Eq. 57 gives
e3/4 e1/2 -Hf- 1o 3 (Et>- 69 >
where E^/^ is the potential of the DME when 1 = 0.75 ijy
When 1 = 0.25 ijj then Eq. 57 gives
El/4 = El/2 log 0.333 (Eq. 70)
Subtraction of Eq. 69 from Eq. 71 gives the following
values of E^/^ in volts at the indicated tempera
tures: 21. -0.0556/n; 25, -0.0564/n; 30, -0.0573/n;
35, -0.0583/n; 40, -0.0592/n; 45, -0.062/n. Experi-
can be compared with the
calculated values as a means of establishing the reversi
bility and electron change of the electrode process (44-47).
mental values of E- E^y^
III. Solubility Analysis
Derivation of equations for solubility analysis of
thiouraclls and thlouracll complexes. The equation relat
ing the dissociation constant of a ligand to its intrinsic
solubility is
Ka = CH+^Ab AbH+)/AbH+
(Eq. 71)


70
where Ab is the absorbance in the buffer region and Abg+
is the absorbance of an acid solution saturated with
ligand. The negative logarithm of Eq. ?1 gives
pK^ = pH log (Ab AbH+)/AbH+ (Eq. 72)
The apparent total solubility of thiouracils as a
function of pH in the absence
S = [HU] + [IT] = SHU + [IT] (Eq. 73)
and presence of 1:1 and 2:1 complex-forming metal ions is
s' = SHU + [IT] + LMU+j + 2[MU2] (Eq. 74)
The respective total solubility in the absence and presence
t
of metal ion is represented by S and S while the intrinsic
solubility of HU is represented by S^. The experimental
conditions for solubility analysis, mainly constant pH and
temperature, are given in part III of the experimental sec
tion. Equation 74 is the same as the mass balance equation
(Eq. 11) for ligand in the case of 2:1 complexes except
that the solution is now saturated with the undissociated
1
ligand. Similarly, the equations for S for mixed ligand
and polynuclear complex formation can be shown to be simi
lar to their ligand mass balance equations (Eqs. 22, 33)


71
where the solutions are saturated with the undissociated
ligand. Subtraction of Eq. 73 from Eq. 74 gives, for
2:1 complexes
S* S = [MU+] + 2[MU2] (Eq. 75)
and subtraction of the appropriate expressions for the
cases of mixed ligand and polynuclear complexes gives,
respectively,
s' S = [MU+J + [MUOH] (Eq. ?6)
s' S = LMU+] + [MU] + 2[M2U2] (Eq. 77)
The definition of the degree of formation, H, is given by
Eqs. 14, 25 and 36 for 2:1, mixed ligand and polynuclear
complexes, respectively. Substitution of Eqs. 75* 76 and
77 into Eqs. 14, 25 and 36, respectively, gives the same
expression for in terms of solubility.
n =
(Eq. 78)
The expressions relating to the stability constants for
2:1, mixed ligand and polynuclear complexes are given by
Eqs. 17, 28 and 40, respectively.


RESULTS
I. Potentlometrlc Titrations
Titration of thiouracll-cuprlc ion mixtures. Solu
tions, 0.002 or 0.0016 M in thiouracil(2-thiouracil,
6-n-propyl-2-thioura.cil, 6-methyl-2-thiouracil, 5-raethyl-
-2-thiouracil and 56-dimethyl-2-thiouracil) and 0.002 to
0.0002 M in cupric nitrate, showed an immediate drop in pH
(ca. 55 to ca. 3*0) and the slow formation of precipitates.
Potentiometric titrations of these suspensions with stan
dard NaOH immediately produced more precipitate. The re
sultant titration curves (curves A, E and C, Figs. 1 and
15). using 2-thiouracil as an example, showed more alkali
consumption to pH 6 than the titration of the same amount
of the ligand alone (curve G, Figs. 1 and 15)* This large
alkali consumption was measured up to inflection 1 of curve
A (Fig. 15) and Inflection 2 of curve A (Fig. 1), a pH
region where the hydrolysis of free cupric ion does not
occur (curve H, Figs. 1 and 15)*
The titration curves indicated the presence of at
least three species. The two most acidic species (pK* 31
Si
and 4.4) are most easily seen in curves A E of Fig. 1.
72


73
A comparison of curves A E of Fig. 15 (titrated immedi
ately after preparation) with the same curves of Fig. 1
(titrated the day after preparation) show that the species
t f
with pK 3*1 increased in titer at the expense of the pK
a a
4.4 titer. The net titer to the pH 6 inflection remains
unchanged. The titer between inflection 1 and 2, curve A,
9
Fig. 15 has the same pK as curve H, the titration of cu-
a
pric ion alone and is assigned thereto.
The titer of excess, uncomplexed 2-thiouracil extends
from inflection 1 to 2, curves D, E and F, Fig. 15 and from
inflection 2 to 3. curves D, E and F, Fig. 1 as can be
ascertained from comparison with the pK 7.49 from the
a
titration of ligand alone (curve G, Figs. 1 and 15)*
If all the cupric ion added was complexed either with
thiouracll or hydroxyl ion then the total alkali consumed
in titrating the complexes (i.e., to pH ca. 6, curves A F,
Fig. 15, inflection 1 and Fig. 1, inflection 2) would be
expected to be twice the molar amount of cupric ion used
because cupric ion is divalent and consumes two equivalents
of alkali per mole of metal ion. This was found to be the
case. The total titer of alkali added to reach inflection
3 of curve A, Figs. 1 and 15, was only slightly less than
twice the titer of alkali consumed by the ligand alone
(curve G), indicating that the cupric-ligand complexes
still exist even at a relatively high pH and that the
higher hydroxide concentration does not cause disruption
of the complex and formation of Cu(0H)2.


Pages
are
misnumbered
following
this
insert


75
Continued titration past the point of precipitation of
the equi-molar ligand-metal solutions gave curve A, Fig. 16,
with an inflection at pH 8.3 corresponding to a total titer
of alkali equal to twice the alkali that would have been
consumed by the ligand alone. The amount of free, uncom-
plexed thiouracil at less than equimolar amounts of lead to
ligand, as estimated from the titer of alkali between in
flections 1 and 2 of curves B G of Fig. 16 (compare with
curve H, Fig. 16), also showed that the precipitating com
plex had a 1:1 stoichiometry. The titer of uncomplexed
thiouracil corresponded to that in excess of 1:1 stoichiom
etry. Infrared analysis of the precipitate, isolated at
pH 7.5 gave a curve identical to the infrared curve of
bis(2-thiouracil-cadmium(II)), MgU^.
The titration data, up to the point of precipitation,
were analyzed according to Eq. 18. Typical plots of log
1-/ versus p[U~] (Eq. 18) are given by Figs. 1?, 18, 19
and 20. The slopes for all plots of Eq. 18 are given in
Table I. In those cases where the metal and ligand con
centrations were equal, the slopes of plots of Eq. 18 were
unity. Therefore, the calculation of log for soluble
1:1, MU+, complexes was valid in these cases (Table I).
When the total metal ion concentration, [M+2]q, was less
than the total ligand concentration, [HU]q, the slope of
Eq. 18 was always between one and two which would be ex
pected if appreciable amounts of MUg were formed.


76
The plots of Eq. 18 are not linear over their entire
length (Figs. 17. 18, 19 and 20). Negative deviations
occur at high pH values due to precipitation. Positive
deviations at low pH values are attributed to error in
reading the very small amounts of initial alkali titer
on the recorder chart. These minor deviations were ignored
when the best straight lines were drawn by sight.
The data from the titration curves up to the point of
precipitation were also plotted according to Eq. 17.
((/l-)[u] versus te) [U-], Fig. 21) and values of log
and log K2 determined from the slopes and intercepts
(Table II). The values of log ranged from approximately
4.7 to 5*0 for the various thiouracils while log K2 was
approximately 3^ at 25 The magnitude of log and log
K2 were not dependent on the total metal ion concentration
used.
Titration of thiouracll-cadmium ion mixtures. Poten-
tiometric titration of all thiouracil-cadmlum mixtures,
except 6-n-propyl-2-thiouracil, gave very similar curves
as those found for the lead studies (Fig. 16). Precipita
tion occurred during the titration at about pH 6.5, limit
ing the calculation of stability constants to the region
of homogeneous solution. The alkali consumed to inflection
1 for the equimolar mixture, represented by curve A, Fig.
16, was twice the alkali consumed by the ligand alone


77
Indicating the precipitation of mixed ligand complexes,
MgUg. If the alkali had destroyed the precipitated com
plexes and formed M(0H)2 the total alkali titer would
have been three times the alkali consumed by the ligand
alone. The titer of excess thiouracil found at cadmium
concentrations less than equimolar, as calculated from
the alkali consumed between inflections 1 (near pH 7) and
2 (near pH 9) also indicated a complex being precipitated
with a 1:1 stoichiometry. The infrared curve of the
2-thiouracil precipitate, isolated at pH 10.5, was identi
cal to the infrared curve of bis(2-thiouracil-cadmium(II)),
M2U2, which was synthesized as outlined in the experimental
section.
The titration curves of cadmium-6-n-propyl-2-thiouracil
solutions at 25 and 35 (Fig* 22) differed from the curves
obtained with the other ligands and from the ligand-lead
mixtures. The titration curves showed precipitation near
pH 6 as before but the first inflection, near pH 8.0, for
all concentrations of cadmium ion equal or greater than
0.001 M (curve A and B, Fig. 22) occurred at an alkali
consumption equal to the titer of alkali consumed by the
ligand alone (the concentration of ligand was constant at
0.002 M). Only when the total cadmium concentration was
less than 0.001 M (curve C, Fig. 22), half of the ligand
concentration, was any uncomplexed ligand indicated. This
behavior requires a stoichiometry of MUg. However, at a
temperature of 45 cadmium-6-n-propyl-2-thiouracil solutions
gave curves as in Fig. 16, indicating a 1:1 stoichiometry.


78
That part of the titration data resulting from homo
geneous solution was analyzed by Eqs. 17 and 18 as explained
in the previous section. The slopes of Eq. 18 (Table I)
for less than equimolar concentrations of cadmium ion were
between 1 and 2 as was the case for lead and indicated
formation of MU2 complexes. When the slope of Eq. 18 was
unity the calculation of log was valid (Table I).
Typical plots of Eq. 18 (Fig. 17. 18, 19 and 20) are the
same as for lead ion. Typical plots of Eq. 17 are shown
in Fig. 21 and the values of log and log K2 are given
in Table II. The log and log K2 values were not depen
dent on the total metal ion concentrations used.
Titration of nickel-thiouracil and zinc-thiouracil
mixtures. Titration of 2-thiouracil, 6-n-propyl-2-thi-
ouracil, 6-methyl-2-thiouracil, 5*-methyl-2-thiouracil and
5,6-dimethyl-2-thiouracil in the presence of nickel gave
curves that showed an initial pH drop from 6.0 to 5*5 and
precipitation during the titration near pH 8 (Fig. 23
curves A D). Analysis of the part of the titration
curves corresponding to homogeneous solution by Eq. 18
gave slopes between 1 and 2 (Table I) indicating a stoi
chiometry of MUg. Analysis by Eq. 17 gave values of log
and log K2 (Table II) from the slope and intercept which
were about two units smaller than the log and log K2
values for the cadmium and lead complexes of the thioura-
cils


79
Titration of 6-n-propyl-2-thiouracil in the presence
of zinc gave an initial pH drop from 6.5 to 6.0 and pre
cipitation occurred near pH 7*3. The titration curves for
the zinc complexes were very similar to those found for
nickel (Fig. 23)* The data were plotted according to Eq.
17 and log and log Kg values are given in Table II.
The stability constants for both nickel and zinc are
smaller, by more than a factor of ten, than those found
for cadmium and lead.
The total alkali consumed to inflection 1 curve A,
Fig. 23, was greater than twice the titer of alkali con
sumed by the ligand alone (curve E) indicating formation
of M(0H)g complexes. The alkali consumed to inflection 2,
curves B D, Fig. 23. was greater than the alkali consumed
by the ligand alone again indicating formation of M(0H)g
complexes and probable disruption of the metal-thiouracil
complexes already formed.
Titration of 2-thlouracll and 6-n-propyl-2-thiouracil
in the presence of other metal ions. Potentiometric ti
trations of 2-thlouracil and 6-n-propyl-2-thiouracil in
the presence of ferric, ferrous, manganese, calcium and
cobaltous ions gave no indication of any complex formation.
The titration curves of the mixtures could be assigned to
uncomplexed ligand in the case of calcium and manganese and
to hydrolysis of the metal ion in the case of ferric, fer
rous and cobaltous ions.


80
Titrations of sterioally blocked thiouracll in the
presence of Cu(Il)t Cd(II) and Pb(Il). Solutions of
2-ethylmercapto-4-hydroxypyrimidine (2EM4HP) and N,N -di-
ethyl-6-methyl-2-thiouracil in the presence of cadmium,
lead and cupric ions were titrated with standard alkali.
The titration curve of 2EM4HP in the presence of cadmium
ion was the same as the titration of the ligand alone.
The titration curves of 2EM4HP in the presence of cupric
and lead ion could be assigned to a simple summation of
the metal hydrolysis and ligand titration curves indicating
no apparent complexation. The titration of solutions of
N,N -diethyl-6-methyl-2-thiouracil alone showed no titrat-
able group and the titrations of metal ion mixtures gave
curves showing only metal hydrolysis. In the case of
2EM4HP, the pK of the 4-hydroxy group was 7.01 at 25.
3
The fact that no complexation occurred at the 4-hydroxy
position even when it could form an anion more easily than
2-thiouracil itself (pK 7.49) argues for complexation at
3
the sulfur position in the sterically unblocked compounds.
t
Titration of aqueous solutions of ligands. The pK
3
of each ligand at 25*0 was determined by averaging at
least three determinations. This is important since small
f
changes in pK have a relatively large effect on the calcu-
Si
lations of the stability constants. The pic' at 35*0 and
3
45.0 were usually single determinations made at the same


81
time the complexation titrations were recorded to avoid
any possible error due to possible temperature effects on
1
the electrode or any other part of the system. The pK
£1
values were estimated from the pH of half-neutralization
when the ligands were titrated with standard alkali. The
I
solutions, temperatures and pK values are given in Table X.
a
Effect of temperature on the stability constants of
cadmium and lead complexes of thlouracils. The values of
log (Table II) at 25, 35 and 45 were used to calcu
late the change in free energy as a function of temperature
by
/\F = -RT In K1
(Eq. 79)
Since
AP = AH TAS (Eq. 80)
the values of AF, the change in free energy at temperature
"T", may be plotted versus the absolute temperature where
AH, the intercept, is the enthalpy change and the slope,
AS is the entropy change of the complexation reaction.
Relatively large error in AF results from small error in
log since the temperature range was purposely restricted
to maintain AH constant (71). Thus, only the general


82
trends in AH and /\S with structural variations in the
thiouracils may have significance. The /\S values (Table
XI) for the first complexation reaction tended to be
positive (apparent trend for negative slopes in Fig. 24)
except the value for lead-6-n-propyl-2-thiouracil which
was zero or slightly negative. The enthalpy change, AH
(Table XI), for the complexation of lead with the ligands
at 375 was negative and varied from -3*5 Kcal. for
6-methyl-2-thiouracil to -7.9 Kcal. for 6-n-propyl-2-thi-
ouracil with an estimated error of + 0.1 Kcal. The
enthalpy change at 375 for the cadmium complex of
5-carboethoxy-2-thiouracil was slightly positive. The
AH values for all other cadmium complexes were negative
except for that of 6-methyl-2-thiouracil which was approxi
mately zero.
The values of AF for the lead complexes, taken as
a group, were about one kilocalorie lower than for the
cadmium complexes.
Effect of temperature on the acid dissociation con
stants of the ligands. Plots of Eq. 80, where the changes
in free energy calculated from the acid dissociation con
stants, K of the ligands (Eq. 79) were plotted against
the absolute temperature (Eq. 80), seemed to give negative
entropy changes (tendency for positive slopes in Fig. 25)
except in the case of 6-n-propyl-2-thiouracil. No clear


83
-cut correlation of /\S with ligand is possible due to the
small temperature range required to maintain AH constant.
The enthalpy changes for the acid ionization were all
positive and varied from 3.5 to 13.2 Kcal. with an
estimated error of + 0.1 Kcal. and are given in Table XI.
II. Ultraviolet Spectral Studies
Jobts plots of 2-thiouracil and 6-n-propyl-2-thiouracil
in the presence of cupric ion. Aqueous solutions of a con
stant concentration of 2-thiouracll (5*00 x 10-*M) and equal
or lesser concentrations of cupric ion were run on the Cary
recording spectrophotometer. The spectra (Fig. 26) showed
a decrease in the absorbance of the 272 mp maximum and a
small rise in absorbance near 3^5 mj-i where 2-thiouracil
has no absorbance. Isosbestic points occurred at 306 and
254 which may indicate a single transformation. A plot
of absorbance at 272 and 3^5 nya versus the ratio of the
concentration of cupric ion to 2-thiouracil, at constant
concentration of 2-thiouracil (5 00 x 10^), showed a
deviation from linearity at 0.5 indicating a stoichiometry
(Cu:U = 1.2) of CuUg for the complex (Fig. 9) on the assump
tion of strong complexation. Jobs plots (72) of aqueous
cupric-2-thiouracil solutions showed rounded peaks with
maxima at a mole fraction of ligand of 0.66 which also
indicates a stoichiometry (Cu:U = 1.2) of CuUg. The com
position of the solutions used in Fig. 8 are given in


84
Table XII. The composition of the solutions used for Fig.
9 are outlined in the experimental section. A Job's plot
of 6-n-propyl-2-thiouracil and cupric ion in 0.05 M HCIO^
showed no discernible change in the absorbance that could
be assigned to complex formation.
Ultraviolet spectra of dilute acid solutions of
cupric-2-thiouracll mixtures. Since the cupric complex
of 2-thiouracil is so strong the equilibrium constant is
best measured in dilute acid so as to maintain significant
concentrations of complexed and free ligand. The ultra
violet spectra of acid solutions of the 2-thiouracil-cupric
ion mixtures listed in section III of the experimental were
run. Prepared 0.1 M HCIO^ solutions containing 4.5 x 10^
M cupric ion were diluted 1;50 and read on the Cary spec
trophotometer. The absorbances at 270 m)i were similar to
that for 2-thiouracil alone. The total change in absorb
ance was about 3 percent which is insufficient for good
analysis. The small change in absorbance, due to the
dilute solutions necessary for UV spectra, would cause
large errors in the calculated stability constants. Higher
pH values would cause precipitation of the complex. This
result indicates that the molar absorption coefficient of
the complex may be about the same as the value for the
uncomplexed ligand. Since the polarographic data indicated
a strong complex does exist then the small change in


85
absorbance is probably due to the low concentrat ions
necessary for UV spectra. The molar absorption coeffi
cient of the complex may be about the same as the uncom-
plexed ligand.
Furthermore, the data indicate that no sulfinic acid
(73) is formed and the proposed oxidation of 2-thiouracil
by cupric ion in alkaline solution does not occur in acid
solution (73) The absorbance of the diluted solutions
-3 _4
containing 4.5 x 10 J M and 2.15 x 10 M cupric ion was
1.305 and 1.250, respectively. The absorbance of 2-thi
ouracil alone is calculated to be 1.233* The absorbance
that was expected if oxidation occurred was 0.995 which
was based on formation of the disulfide and subsequent
disproportionation to the sulfinic acid and 2-thiouracil.
The calculated absorbance is based on the molar absorption
coefficients of sulfinic acid and 2-thiouracil given in the
experimental section and Eq. 2.
Spectrophotometric titration of 2-thiouracil. Anal-
ysis (Eq. 3) of the change in absorbance as a function of
increasing pH of aqueous solutions of 2-thiouracil gave a
pK near 7*6. Isosbestic points occur at 302, 258 and 219
3.
m;i. The spectra and plots of Eq. 3 are shown in Fig. 10
and 11.


86
Spectrophotometric titration of 2-thiouracil-cupric
Ion mixtures. The shapes of the spectra of highly dilute
aqueous mixtures of cupric ion and 2-thiouracil (Fig. 12)
as a function of increasing pH were nearly the same as for
the curves obtained during the spectrophotometric titration
in the absence of cupric ion (Fig. 10) except in the region
of 255 to 305 mp where the absorbance was reduced due to
the complex. A rise in absorbance near 3^5 mp was similar
to the rise in absorbance of spectra analyzed in the Job's
plots. The value of pK (7.5^) was only slightly changed
9
from the value of 7*3 found in the absence of metal ion.
These results indicate that the cupric complex of 2-thi
ouracil absorbs ultraviolet light less than the ligand
does in the pH region 6.0 to 10.0 between 255 and 305 mp*
The complex weakly absorbs ultraviolet light above 305 mp.
-4
Concentrations of cupric ion 0.500 x 10 and larger caused
precipitation which frustrated analysis. The useful data
4
collected at a cupric concentration below 0.500 x 10 were
analyzed using Eq. 3 at 270, 285, 230, 240, 310 and 320 mp.
The spectra and plots of Eq. 3 are given in Figs. 12 and
13.
Degradation of 2-thiouracil disulfide in the absence
and presence of cupric ion. It is known (73) that hydrol
ysis of disulfides proceeds by disproportionation according
to the following reaction.


87
2 RSSR + 3 H20 > 3 RSH + RSOgH (Eq. 81)
Cupric ion can oxidize 2-thiouracil to the disulfide in
dilute alkali (14). It was conceivable that the reaction
might also occur in dilute acid but would yield the sul-
finic acid and 2-thiouracil according to Eq. 81, since
it has been shown (14) that the disulfide of 2-thiouracil
is very unstable in dilute acid and produces 2-thiouracil
as one of the products. If cupric ion is present when N
moles of 2-thiouracil disulfide disproportionate the fol
lowing series of reactions can be expected.
N RSSR + 2N H20 > 3N/2 RSH + N/2 RSOgH
(Eq. 82)
3N/2 RSH + N Cu+2 > N/2 RSH + N/2 RSSR + N Cu+
(Eq. 83)
N/2 RSSR + N H20 > 3N/4 RSH + N/4 RSOgH
(Eq. 84)
The sum of the final concentrations of RSH in Eq. 83 and
84 is
^ RSH = 5N/4 RSH
(Eq. 85)


88
The sum of the final concentrations of RSOgH is
^ RS02H = 3N/4 RS02H (Eq. 86)
The final absorbance of the solution will be the sum of the
absorbances of RSH and RSOgH. By using Eq. 2, the molar
absorption coefficients of 2-thiouracil (^270 13 680) and
its sulfinic acid derivative (£ 2^q 3.510) and the original
concentration of 2-thiouracil disulfide (N = 3*99 x 10^ M),
the final absorbance can be calculated in the absence and
presence of cupric ion. The expected absorbance in the
absence of cupric ion, as a result of the disproportiona
tion expected in Eq. 82 is O.89. The observed absorbance
when equimolar and twice equimolar concentrations of cupric
ion (N and 2N) as of disulfide were mixed, was 0.83 This
was in reasonable agreement with the expected value of
O.89. The small difference may be due to partial formation
of the sulfonic acid derivative of 2-thiouracil (73).
The expected change in absorbance for the postulated
sequence of Eqs. 82 86 when N and 2N relative concentra
tions of cupric ion were used would be 12 percent and 23
percent, respectively. However, the final absorbance
values for disulfide in acid solution were the same in
the presence and absence of cupric ion (i.e., O.83). It
can be concluded that oxidation does not occur in acid


89
solution and further substantiates the spectral studies
which indicated that the molar absorption coefficient of
the cupric complex of 2-thiouracil in acid solution must
be nearly the same as for 2-thiouracil.
Aqueous solubility of thlouraclls as a function of
pH. The intrinsic solubility of 2-thiouracil, 6-n-propyl-
-2-thiouracil, 6-methyl-2-thiouracil, 5-methyl-2-thiouracil,
5,6-dimethyl-2-thiouracil, 5-carboethoxy-2-thiouracil and
6-amino-2-thiouracil varied from 7*97 x 10"^ to 1.79 x 10^
Ab-Ab+
M. Plots of log versus pH were prepared and a
AbH+
typical plot is given in Fig. 27* The calculated intrinsic
t
solubilities and pK values are given in Table IV for the
a
various ligands.
Aqueous solubility of 2-thiouracil and 6-n-propyl-2-
-thlouracll as a function of pH in the presence of cadmium
ion. Due to the fact that precipitation of complexes must
be avoided in solubility analyses, the pH range was re
stricted to values less than 6.5. This restriction
drastically reduced the concentration of 2-thiouracil anion
and suggested calculations which indicated that only a 1.3
percent increase in total solubility of 2-thiouracil could
be predicted from the stability constants obtained from
the potentiometric titrations data. This increase is less
than experimental error and thus no stability constants
could be obtained by this method.


Full Text
STABILITY CONSTANTS AND STRUCTURES
OF METAL COMPLEXES OF THE
ANTITHYROID THIOURACILS
By
DENNIS JOSEPH WEBER
A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
December, 1967

DEDICATION
To my wife,
Shirley,
whose encouragement, help and patience
made this meaningful and possible.

ACKNOWLEDGMENTS
The author gratefully acknowledges Dr. Edward R.
Garrett, chairman of the supervisory committee, for his
research guidance and his assistance in the preparation
of this manuscript.
The author is most grateful for the laboratory assis¬
tance of his wife, Shirley A. Weber, and for her patience
and understanding.
Sincere appreciation is expressed to Dr. Melvin J.
Fregly for his willing help in the animal studies.
The author extends his sincere thanks to the many
people of the College of Pharmacy for their time and
energy which were very helpful.
The author is indebted to Dr. Alan Agren of the
University of Uppsala, Sweden, for the helpful information
and discussions he so willingly gave.
This investigation was supported by a Public Health
Service Fellowship from the Division of General Medical
Sciences, Public Health Service. The author is deeply
grateful for this support.
The author wishes to thank his typist, Mrs. Arthur
Grant for a superb job.
iii

Appreciation is extended to Dr. Melvin J. Fregly,
Dr. Werner M. Lauter and Dr. R. Carl Stoufer for serving
on the supervisory committee.
iv

TABLE OF CONTENTS
Page
ACKNOWLEDGMENTS iii
LIST OF TABLES vii
LIST OF FIGURES ix
INTRODUCTION 1
EXPERIMENTAL 9
I. Materials 9
II. Potentlometric Titrations 15
III. Ultraviolet Spectral Studies 19
IV. Polarographic Studies 29
V. Synthesis of Complexes 35
VI. Animal Experiments 41
VII. Magnetic Susceptibility Measurements 43
EQUATIONS 44
I. Potentlometric Titrations 44
II. Polarography 5ti
III. Solubility Analysis 69
RESULTS ?2
I. Potentlometric Titrations 72
II. Ultraviolet Spectral Studies 83
III. Polarography 92
IV. Animal Experiments 96
v

Page
DISCUSSION 100
I. Structure of Precipitated Cupric-Thiouracil
Complexes 100
II. Structure and Stability of Cupric-Thiouracil
Complexes in Homogeneous Solutions 105
III. Cadmium and Lead Complexes of Thiouracils.... Ill
IV. Complexation of Other Metal Ions with
Thiouracils 124
V. Biological Effects 129
SUMMARY AND CONCLUSIONS 135
APPENDIX
A. Tables 139
B. Figures..... I76
BIBLIOGRAPHY 235
BIOGRAPHICAL SKETCH 241
vi

LIST OF TABLES
Page
Table
I COMPOSITION OF SOLUTIONS AND ESTIMATED
STABILITY CONSTANTS AT VARIOUS TEMPERATURES
ON THE POSTULATE OF 1:1 METAL COMPLEXES OF
SUBSTITUTED THIOURACILS FROM POTENTIOMETRIC
TITRATIONS 140
II COMPOSITION OF SOLUTIONS AND ESTIMATED
LOGARITHMIC STABILITY CONSTANTS, K1 AND Kg,
AT VARIOUS TEMPERATURES ON THE POSTULATE OF
MIXED 1:1 AND 2:1 METAL COMPLEXES OF SUBSTI¬
TUTED THIOURACILS FROM POTENTIOMETRIC
TITRATIONS 144
III COMPOSITION OF AQUEOUS SOLUTIONS AND ULTRA¬
VIOLET ABSORBANCE OF AQUEOUS THIOURACIL-METAL
ION MIXTURES AT 25° 149
IV INTRINSIC SOLUBILITIES OF LIGANDS AND pK*
Si _
VALUES DETERMINED FROM SOLUBILITY DATA AT 25°.. 157
VEFFECTS OF CONCENTRATION OF CUPRIC ION AND PER¬
CHLORIC ACID ON DIFFUSION CURRENT, HALF-WAVE
POTENTIAL, E1/2 AND E^ - E^ VALUE FOR
POLAROGRAPHIC REDUCTION OF CUPRIC NITRATE 158
VIEFFECTS OF 6-n-PROPYL-2-THIOURACIL AND PER¬
CHLORIC ACID CONCENTRATIONS ON E.- E^ AND
(Ei). - (E¿)_ VALUES OF 2.00 x lO"4 MOLAR CUPRIC
ION AT 21°. 160
VIIEFFECT OF 5.6-DIMETHYL-2-THI0URACIL ON RAT
THYROID WEIGHT AFTER TWO WEEKS OF DRUG DIET.... 161
vli

Table Page
VIIIEFFECT OF 6-METHYL-N,N*-DIETHYL-2-THI0URACIL
ON THE RAT THYROID WEIGHT AFTER TWO WEEKS OF
DRUG DIET 164
IXEFFECT OF 6-n-PROPYL-2-THIOURACIL AND 5-METHYL-
-2-THIOURACIL ON THYROID WEIGHT AFTER TWO WEEKS
OF DRUG DIET 166
XNEGATIVE LOGARITHMS OF DISSOCIATION CONSTANTS
(pK*) OF LIGANDS AS A FUNCTION OF TEMPERATURE.. 170
81
XIENTROPY AND ENTHALPY VALUES FOR THE IONIZATION
OF THIOURACILS (HU) AND FOR THE FORMATION OF
MU+ BY COMPLEXATION WITH CADMIUM AND LEAD 171
XIICOMPOSITIONS AND ABSORBANCES OF 2-THIOURACIL-
- CUPRIC ION SOLUTIONS FOR JOB*S PLOTS 172
XIIINEGATIVE LOGARITHM OF THE RATIO OF THE DIS¬
SOCIATION CONSTANT OF THE OXIDIZED COMPLEX TO
THE DISSOCIATION CONSTANT OF THE REDUCED COM¬
PLEX OF CUPRIC-THIOURACILS (KQ/Kr) AS A FUNC¬
TION OF TEMPERATURE AND ACID CONCENTRATION 173
viii

LIST OF FIGURES
Figure Page
1 Potentiometric titration curves of filtered,
day old cupric nitrate-2-thiouracil mixtures
with ]x = 0.00b at 25.0° 1?8
2 Ultraviolet spectra of mixtures of cupric
nitrate and 6-methyl-2-thiouracil in water
at 25*0° 180
3 Ultraviolet spectra of mixtures of lead nitrate
and 2-thiouracil in water at 25*0° 182
4 Ultraviolet spectra of ferric nitrate and
6-methyl-2-thiouracil in water at 25*0° 184
5 Ultraviolet spectra of mixtures of cupric
nitrate and N,N -diethyl-6-methyl-2-thiouracil
in water at 25*0° 186
6 Ultraviolet spectra of mixtures of cupric
nitrate and 2-ethylmercapto-4-hydroxypyrimldine
in water at 25*0° 188
7 Plot of absorbance at 2?2 mp versus the mole
fraction of cupric nitrate of aqueous solutions
of cupric nitrate and 2-thiouracil at 25.0° 190
8 Job’s continuous variations plots of absorbance
of aqueous mixtures of cupric nitrate and
2-thiouracil 192
9 Plots of ultraviolet absorbance at 272 and 345
mp versus [ Cu+2]/[ 2TU] 194
10 Ultraviolet spectra of aqueous 2-thiouracil
solutions as a function of pH at 25*0° 196
11 Plots of absorbance of aqueous 2-thiouracil
solutions versus pH 198
12 Ultraviolet spectra of an aqueous solution
containing cupric nitrate and 2-thiouracil as
a function of pH at 25*0° 200
ix

Page
Figure
13
14
15
16
17
18
19
20
21
22
23
Plots of absorbance of an aqueous solution
containing cupric nitrate and 2-thiouracil
versus pH...*» 202
_k
Polarogram of 2.00 x 10 M cupric nitrate in
the presence of 1.875 X 10"3 m 2-thiouracil and
0.0800 M HC104 at 21.0° 204
Potentiometric titration curves of aqueous
mixtures of cupric nitrate and 2-thiouracll
containing precipitated complex with u = 0.006
at 25.0° 206
Potentiometric titration curves of aqueous
solutions of lead nitrate and 2-thiouracil
with p = 0.006 at 25.0 208
Plot of log 0--ñ/ñ against the negative logarithm
of 6-n-propyl-2-thiouracil anion concentration
obtained from lead nitrate -PTU mixture in
water with p = 0.006 at 25*8° 210
Plot of log (1-ñ/ñ against the negative logarithm
of 6-n-propyl-2-thiouracil anion concentration
obtained from lead nitrate - PTU mixture in water
with p = 0.006 at 25*8° 212
Plot of log (l-ñ)/ñ against the negative
logarithm of 6-n-propyl-2-thiouracil anion
concentration obtained from lead nitrate - PTU
mixture in water with p = 0.006 at 25*8° 214
Plot of log (l-ñ)/ñ against the negative
logarithm of 6-n-propyl-2-thiouracil anion
concentration obtained from lead nitrate - PTU
mixture in water with p = 0.006 at 25.8°
Plots of H/(1-H)[2TU~] against
n-2
ñ-1,
[2TU-] from
aqueous mixtures of 2-thiouracil and lead
nitrate with p = 0.006 at 25*0°
216
218
Potentiometric titration curves of aqueous
mixtures of cadmium nitrate and 6-n-propyl
-2-thiouracil with p = 0.006 at 25«0°
Potentiometric titration curves of aqueous
mixtures of nickel nitrate and 2-thiouracil
with p = 0.006 at 25.0° 222
x

Page
Figure
24 Plot of the change in free energy (AF) versus
the absolute temperature calculated from the
stability constants of the 1:1 complexes (Mu )
for the cadmium and lead nitrate complexes of
2-thiouracil, 6-n-propyl-2-thiouracil,
5,6-dimethyl-2-thiouracil, 6-methyl-2-thiouracil,
5-methyl-2-thiouracil and 5-carboethoxy-2-
-thiouracil 224
25 Plot of the change in free energy (AF) versus
the absolute temperature calculated from the
acid dissociation constant of 2-thiouracil,
6-n-propyl-2-thiouracil, 5.6-dimethyl-2-
-thiouracil, 6-methyl-2-thiouracil, 5-methyl-
-2-thiouracil and 5-carboethoxy-2-thiouracil.... 226
26 Ultraviolet spectra of aqueous solutions at 25*0°
containing 2-thiouracil and cupric nitrate 228
27 Plot of log (Ab - AbH+)/AbH+ against pH for
aqueous, saturated 2-thiouracil solutions at 25*0°
according to log (Ab - Ab^+)/Ab^+ = pH - pK^.... 230
28 Plots of -[(Ex)_ - (Ei)_] against the negative
2 ^ 2 S
logarithm of the 6-n-propyl-2-thiouracil
concentration 232
29 Plot of the logarithm of the Apparent Relative
Antithyroid Activity versus the negative
logarithm of the dissociation constant ratio
of the 2:1 cupric-ligand and 1:1 cuprous-ligand
complexes with thiouracils 234
xi

INTRODUCTION
The compound. 2-thiouracil (I) is used in the treatment
of hyperthyroidism.
(I) H
In general, thiouracils are rapidly absorbed by the
gastrointestinal tract (1) and 50% of the administered
dose is excreted in the urine (2,3)» At least 15% of
an oral dose is destroyed in the G. I. tract (4,5) and
is not available for absorption. The identity of poten¬
tial metabolites is unknown (4).
The probable mode of action of thiouracil is inter¬
ference with the incorporation of iodide into thyronine,
or its precursors, to produce thyroxine (6,7).
Thyroxine (V) is the hormone, elaborated by the
thyroid gland, which controls the rate of metabolism
and oxygen consumption of the body. High blood levels
1

2
of thyroxine, indicative of a hyperactive thyroid gland,
can cause loss of weight and high blood pressure (8), A
low blood concentration of thyroxine stimulates the ante¬
rior pituitary to secrete Thyroid Stimulating Hormone
(TSH) which causes an Increased rate of production of
thyroxine by the thyroid gland, a negative feedback mecha¬
nism. It has been suggested (8) that the production of
thyroxine proceeds by iodination of two molecules of the
amino acid tyrosine (II) which condense to form a molecule
of thyroxine (V) (Scheme I).
II Tyrosine
III Monoiodotyrosine
I
IV Diiodotyrosine
2 H
•CH^HNH^OOH —> H
ch2chnh2cooh
V Thyroxine
Scheme I

3
Goiter, the enlargement of the thyroid gland, can
occur in two ways (8). Dietary deficiency of iodide will
result in a compensating hyperplasia of the thyroid to
trap as much of the available blood iodide as possible.
This condition, nontoxic goiter, usually results in normal
thyroxine blood levels. If the negative feedback mechanism
controlling the blood level of thyroxine fails, then the
anterior pituitary output of TSH escapes control by thyrox¬
ine and the thyroid gland is stimulated to produce thyrox¬
ine at an abnormal rate. This condition is termed toxic
goiter and can be controlled by antithyroid drugs such as
the thiouracils.
Alkyl substitution at the 5 or 6 position of thioura-
cil (I) usually enhances its antithyroid activity (9.11 )•
For example, the 5-methyl, 6-methyl and 5.6-dimethyl sub¬
stituted thiouracils are active, with potencies relative
to 2-thiouracil, of 0.7, 1.0 and 1.2, respectively (9,10).
The present antithyroid derivative of choice, because of
its maximal activity and low toxicity in the intact animal
is 6-n-propyl-2-thiouracil (10,12,13). Alkylation of
thiouracil at the N-l, N-3 or sulfur positions greatly
reduces (10) and substitution by electronegative groups
at the 5 or 6 position nearly eliminates any antithyroid
activity (9-11).
The oxidation of thiouracil (I) by iodine has been
shown to occur with ease at physiological pH values (14).

4
Thiouracil is selectively oxidized by iodine even in the
presence of tyrosine (II) (14),the thyroxine (V) precursor
(15). In addition, no iodine was found in the recovered
tyrosine (14) so no iodinated tyrosine (111,1V) was formed.
The product of the iodine oxidation of 2-thiouracil is the
The disodium salt of VI is stable but the free acid readily
disproportionates to thiouracil and higher oxidation pro¬
ducts (14). The ease of oxidation of 2-thiouracil might
explain the degradation observed in the gastrointestinal
tract. This possibility is even more attractive when the
hydrolytic stability of thiouracil is considered (16).
It must be emphasized however, that ease of oxidation by
iodine cannot be the only determining factor since thioura¬
cil derivatives with no antithyroid activity are also
easily oxidized (14).

5
Cupric ion has been implicated, in thyroid function
(17). The copper content of the normal and pathologic
thyroid has been determined (18) and verified by Kasanen
and Viitanen (19) who found elevated copper levels in
toxic and nontoxic goiters. The formation of diiodo-
tyrosine and thyroxine is Increased when cupric ion is
added to homogenates of thyroid gland (20). Other evidence
that cupric ion aids in the formation of thyroxine, by
formation of iodine from iodide, has been presented (21,
22,23 ).
Since cupric ion and other heavy metals precipitate
thiouracil and its derivatives from aqueous solution (24,
25), Libermann conjectured that complexing ability and
antithyroid activity may be correlated. He assumed that
completeness of precipitation could be taken as a measure
of the stability of the complex. However, this assumption
is not always true. The extreme stability of some EDTA
complexes and their very high water solubility is a case
in point. He suggested a structure (VII) for the cupric
-thiouracil complex (25) which assumed a 1:1 stoichiometry
of metal to ligand and chelate binding of the cupric ion
by the sulfur at the 2 position and the oxygen at the 4
position. A consideration of the stereochemistry of
2-thiouracil and the square-planar nature of cupric ion
shows that the proposed structure is impossible because

6
(Vil)
H
the phenolic oxygen and thionyl sulfur are coplanar and
physically distant.
Cupric ion was presumed to react with thiouracil in
1.0 N NaOH in one day’s time to produce disulfide complexes
of cuprous ion (26). Elemental analyses were obtained on
the isolated complexes and a possible structure (VIII) was
proposed. It has been demonstrated that alkaline thioura¬
cil solutions are susceptible to air oxidation even in the
absence of cupric ion (16). Thus, the possibility exists
that the isolated complex contains cupric and not cuprous

7
ion as was suggested (26). To prove the existence of VIII
a solution of thiouracil disulfide in water was reacted
with cuprous chloride in concentrated hydrochloric acid
and a yellow product isolated (26). Weiss and Venner (26)
proposed, on the basis of nitrogen, copper and chloride
analyses, the substitution of chlorine for hydroxyl on
the copper in VIII. The extreme instability of thiouracil
disulfide (14) casts doubt on this assertion. It is pos¬
sible that the product may actually be IX or even X if the
ease of disproportionation of cuprous to cupric ion is
considered (27)»
The ease of the oxidation of 2-thiouracil by iodine
suggests that thiouracil*s mechanism of action may be the
reduction of iodine to the ineffective iodide. However,
since cupric ion has been implicated in thyroid function
at the level of iodine production, the removal of cupric
or cuprous ion by complexation with thiouracil could be

8
an alternate explanation for its mode of action. If the
complexation of cupric or cuprous copper is important in
the mechanism of action of the thiouracils then the sta¬
bility constant of the complex may be larger for copper
than with other physiological metal ions. Furthermore,
the stability constant of the copper complex may be cor¬
related with the biological activity of the particular
thiouracil derivative. The correlation can only be ex¬
pected under the conditions of equal concentrations at
the site of action. Any differences in the in vivo
solubility or stability of the thiouracil derivative
should be taken into consideration.
The principal objective of this work was to provide
quantitative Information of the complexation of metal ions
with thiouracils. The types of metal ions which complex,
the effect of thiouracil substituents on the stability
constants and the structure of the complexes were de¬
termined in an attempt to predict substituted thiouracil
activity.

EXPERIMENTAL
I> Materials
Purification of 2-thlouracil. 2-Thiouracil (Nutri¬
tional Biochemical Corp., Cleveland, Ohio) was recrystal¬
lized from hot water. The product was washed with water
sind acetone and dried in a vacuum oven at 80°, m.p. 322-323°
dec. (all m.p. are uncorrected); literature value 310-312°
dec. (28), ca. 340° (29)* Equivalent weight 130.3; calcu¬
lated for C^H^NgOS 128.1. I. R. spectrum (30), 0 in cm.-1
(Nujol mull): 3020 (NH); 1680 (C=0); 1280, 1240, 11??.
U. V. spectrum (28), (0.1 MHCIO^), Xmax> 273 (€ 13,700),
Xmax. 212 (€ 16.600).
Purification of 6-n-propyl-2-thiouracll. The com¬
pound (Nutritional Biochemical Corp. and K & K Laboratories
Plainview, New York) was recrystallized from hot water and
dried at 80°, m.p. 219-221°; literature value 219-221°
(31), Equivalent weight 170.0; calculated for C^H^NgOS
170.2. I. R. spectrum, 0 in cm.“* (Nujol mull): 3100
(NH); 1650 (C=0); 1550, 1240, 1190. U. V. spectrum (0.1
MHC104), \max# 272 (€ 15,840), Xmax< 214 (£15,840).
9

10
Purification of 6-methyl-2-thlouracil. The compound
(Nutritional Biochemical Corp. and K & K Laboratories) was
recrystallized from hot water and dried in a vacuum oven
at 50°, m.p. 331-332° dec.; literature value >300° (9)»
Equivalent weight 142.1; calculated for C^H^NgOS 142.1.
I. R. spectrum, Qin cm.-1 (Nujol mull): 3100 (NH); 1640
(C=0); 1195. 1165. U. V. spectrum (0.1 MHCIO^), Xmax#
274 (€15.460), Xmax< 213 (€ 15.760).
Purification of 5.6-dlmethyl-2-thlouracll. This
material (K & K Laboratories) was recrystallized from hot
water and dried at 50° in a vacuum oven, m.p. 286-287°
dec.; literature value 283-285° (9)» Equivalent weight
156.5; calculated for C^HgNgOS 156.2. I. R. spectrum,V
in cm.”1 (Nujol mull): 3210, 3110 (NH); 1660 (C=0); 1600
1210, 1130. U. V. spectrum (0.1 M HC10k), Xmov 276 (€
17.340), Xmax> 215 (€ 14,020).
Purification of 5-methyl-2-thlouracll. 5-Methyl-
-2-thiouracil (Sigma Chemical Co., St. Louis, Missouri)
was recrystallized from hot water, washed with water and
dried at 50° in a vacuum oven, m.p. 334° dec.; literature
value not available (32). Equivalent weight 141.2; calcu¬
lated for C^H^NgOS 142.1. I. R. spectrum, Din cm."1
(Nujol mull): 3090 (NH); 1640 (C=0); 1240, 1200, H65.
U. V. spectrum (0.1 MHCIO^), Xmax# 274 (€ 15,450), Xmax#
213 (€ 15,730).

11
Purification of 5-carboethoxy-2-thlouracll. 5-Carbo-
ethoxy-2-thiouracil (Cyclo Chemical Corp., Los Angeles,
California) was used as received, m.p. 245-246°; literature
value 245° (33)» Equivalent weight 197*5; calculated for
CyHgNgO^S 200.2. U. V. spectrum (0.1 MHCIO^), lmax> 310
(615.121), Anax# 269 (68.991). Xmax# 213 (610,762).
Purification of 2-thio-6-amlnouracll. 2-Thio-6-amino-
uracil was used as received, m.p. 330°; literature value
295° (34). U. V. spectrum (0.1 MHCIO^), 275 (6
18,413). XmsiXa 202 (65.856).
. T
Synthesis and purification of 6-methyl-N,N -dlethyl-
-2-thiouracll. This material was synthesized by the pro¬
cedure of Lacey (35)* N,N*-diethylthiourea (63*2 gm.)
(Eastman Organic Chemical Co,, Rochester, New York) was
added to 20 ml. of glacial acetic acid and brought to a
boil in a round bottom flask fitted with a reflux condenser
and a dropping funnel containing 8.6 gm. of diketene (K & K
Laboratories, Plainview, New York). The diketene was added
over a half-hour period and the reaction was allowed to
cool overnight. The reaction was further heated for half
an hour and then cooled and the acetic acid removed by
vacuum evaporation. Twenty ml. of water was added to the
residue, the mixture was shaken to emulsify and put into
a refrigerator overnight. The precipitated contents were
recrystallized from hot water and dried in a vacuum oven

at 50°, m.p. 97-98°; literature value 97-98° (35). I. H.
spectrum, Din cm.”* (Nujol mull): 1680 (C=0); 1250, 1105
U. V. spectrum (HgO), 278 (£13,100), Xmax# 222 (6
15.250). Potentiometric titration with 0.1 N NaOH indi¬
cated no titratable acid function present. Yield is 73%.
Synthesis and purification of 2-ethylmercapto-4-hy-
droxypyrimldine. Eight gm. of 2-thiouracil (0.062 mole)
were added to 2.49 gm. NaOH (0.062 mole) in 100 ml. of
water and acetone added and the solution cooled overnight
in a refrigerator. A precipitate of the sodium salt of
2-thiouracil formed (8.3 gm., 0.055 mole) which was fil¬
tered and collected.
The sodium salt of 2-thiouracil (O.O55 mole) and
ethyliodide (0.06 mole, 9*35 gm.) were added to 120 ml.
of 95$ ethanol in a round bottom flask and refluxed until
the sodium thiouracil had gone into solution. It was
necessary to add an additional 4 gm. (0.026 mole) of ethyl
iodide during the reaction to put the sodium thiouracil
into solution. The reaction mixture was cooled and the
ethanol removed by vacuum evaporation. A white residue
was left which was recrystallized from ethanol once and
then finally purified by sublimation, m.p. 152-153°;
literature value 152° (36). I. R. spectrum, \j in cm.”*
(Nujol mull): 1660 (C=0), 1270, 1170, 1540. U. V. spec¬
trum (H20), lmax> 280 (6 5.500), Xmax> 230 (6 11,750).

13
Equivalent weight 156.4; calculated for C^HgNgOS 156.2.
Anal.1 Caled, for CgHgNgOS: C, 46.13; H, 5*16; N.
17.94; S, 20.52.
Found: C, 46.40; H, 5»31; N, 18.04; S. 20.37.
Preparation and purification of disodlum salt of
2-thlouraoll disulfide monohydrate. This material was
prepared according to the procedure of Miller, et al. (14).
2-Thiouracil (2.56 gm.) was dissolved in 41 ml. of 1.0 N
NaOH and the solution cooled in an acetone-dry ice bath
until nearly frozen. Sodium iodide (4.57 gm.) and iodine
(2.54 gm.) were dissolved in 20 ml. of water and also
cooled in acetone-dry ice. The iodine-iodide solution
was added slowly, with stirring, to the alkaline 2-thioura-
cll solution in an 800 ml. beaker. About 500 ml. of ace¬
tone was added and the slow forming white precipitate was
collected. The precipitate was recrystallized from water
-acetone twice and dried at 70°. Equivalent weight 156.6;
calculated for CgH^N^OgSgNagHgO 158.1. U. V. spectrum (14)
(H20)’ Xmax. 270 (€ 18*520), Xmax> 211 (£ 25,485); (pH 9-56
ammonia buffer) 273 (£ 12,469). *max#213 (€ 31,537).
I. R. spectrum, in cm."1 (Nujol mull): 3250 (OH), 1570
1530, 1350, 1170, 1000, 828, 720. Molecular weight (ebul-
lioscopic in water) 96; calculated for CgH^N^OgSgNagHgO 105.
1The elemental analyses, molecular weight and weight
loss measurements in this investigation were performed by
Huffman Laboratories, Inc., Wheatrldge, Colorado.

14
Anal» Caled, for CgH^N^SgNagHgO: C, 30.38; H, 1.91;
N, 17.71; S, 20.27; Na, 14.54.
Founds C, 29.87; H, 2.73? N, 14.38; S, 20.67; Na,
15.15.
Synthesis and purification of disodlum salt of 2-sul-
phlnyl-4-hydroxypyrlmidlne trihydrate. 2-Thiouracil (2.56
gm.) was dissolved in 26 ml. of 1.0 N NaOH and 2.1 ml. of
34# hydrogen peroxide was slowly added with stirring to
the acetone-dry ice chilled solution. Acetone was added
to the point of precipitation and the reaction mix main¬
tained cold overnight. The reaction mix was filtered and
the precipitate washed with acetone, the washings added
to the filtrate and the filtrate put hack into the refrig¬
erator overnight. A white precipitate came down which was
collected, recrystallized twice from acetone-water and
dried at 50° in a vacuum oven. Equivalent weight 257.3;
calculated for C^H2N20^SNa2*3H20 256.1. U. V. spectrum
(0.05 M HC104), Xmax# 257 (€ 4,924), Xffiax# 213 (€ 5.885),
Xmax. 270 (€3.510), (0.1 MHC104), Xffiax> 270 (€3.631).
Anal. Caled, for CjjHgNgO^SNag: C, 23.8; H, 1.00;
N, 13.8; S, 15.87; Na, 22.8; weight loss for trihydrate,
21.15.
Found: C, 23.96; H, 3*16; N, 12.38; S, 16.27; Na,
21.03; weight loss, 21.53*

15
II» Potentlometric Titrations
Description of titration apparatus, All potentio-
metric titrations were performed using a Sargent Model D
automatic titrator equipped with a 2.5 ml. capacity syringe
burette. The sample solutions were titrated in water-jack¬
eted titration cells and the temperature of each titration
was recorded. All sample solutions were flushed with
nitrogen gas, which was passed through alkaline pyrogallol
and then water to remove carbon dioxide, and a flow of the
gas into the solution during the titration was maintained.
A Beckman combination electrode with a Ag-AgCl reference
was used. The electrode was kept in water at all times
and a solution saturated with potassium chloride and silver
chloride added as needed. The pH scale of the titrator was
standardized before each titration using pH 4 and pH 7
standard buffers obtained from the Sargent Company. The
pH standardization was checked after each titration and in
no case was the pH drift larger than 0.05 pH units and
usually was less than 0.02 units.
The 1/10 normal sodium hydroxide titrant was prepared
carbonate-free by dissolving a weighed amount of sodium
hydroxide pellets in a minimum of water, filtering off the
precipitated sodium carbonate through a fritted glass fun¬
nel and diluting the filtrate with boiled, nitrogen purged
water. The titrant was stored in a pyrex bottle fitted

16
with ground glass connections directly to the burette on
the titrator. The titrant bottle was also fitted with an
Ascarite tube to avoid carbon dioxide absorption. The
sodium hydroxide titrant was standardized using primary
standard grade potassium acid phthalate.
Preparation of titration sample solutions. Sample
solutions for titration were prepared with constant con¬
centrations of thiouracil and varying concentrations of
metal ion (TablesI and II). Solution conditions for the
thiouracil-cupric ion titrations were: thiouracil deriva¬
tive, 0.00200 M; cupric ion, 0.00200, 0.00160, 0.00100 and
0.000400 M except for 2-thiouracil which varied from
0.00200 to 0.000200 M in 0.000200 M steps. A series of
2-thiouracil-cupric ion solutions, prepared as above, was
left standing overnight under nitrogen, filtered and a
4/5 fraction of the filtrate titrated in the usual way
(Fig. 1). No calculation of the stability constants was
possible due to the precipitation of the complex. All
thiouracil stock solutions were tested for stability by
ultraviolet analysis under working conditions of room
temperature and exposure to light. In all cases the
ultraviolet spectrum changed less than one percent over
a period of one month. Stock solutions older than one
month were discarded. The ionic strength was maintained

17
constant at 0.006 by substituting 0.03 M sodium perchlorate
for equal volumes of 0.01 M divalent metal nitrate in ac¬
cordance with Eq. 1 where ja is the ionic strength, is
the molar concentration of the ion and is the charge
on the ion.
(Eq. 1)
The ratio of the molar concentrations of sodium perchlorate
to divalent metal nitrate is 1:3 for equal ionic strengths.
Standardization of metal ion solutions. Analytical
grade nitrate salts of cupric, cadmium, lead, nickel, fer¬
ric, cobalt, calcium, zinc and manganese were used to pre¬
pare stock solutions in distilled water. The stock solu¬
tions were standardized using the mercury, mercury-EDTA
electrode (Sargent Co., Chicago, Illinois) of Reilley (37.
38). The general procedure was to titrate potentiometri-
cally solutions containing 2 ml. aliquots of 0.01-0.04 M
metal nitrate, 25*0 ml. of buffer (38) and one drop of
1.0 x 10“3 M Hg-EDTA solution (38) with 0.1000 M EDTA
(EthyleneDiamineTetraAcetic acid). The reference electrode
was a saturated calomel electrode. In the case of cobalt,
an excess of standard calcium was added and the excess
determined by titration with EDTA according to the general
procedure outlined. Since Reilley’s procedure cannot be

18
used to standardize ferric ion (38), the procedure of
Pribil, et al. (39) was used. The procedure was to titrate
a solution containing 5 ml» of 0.01 M ferric ion, 25 »0 ml.
of a pH 3 chloroacetic acid buffer (0.2 M) and one drop of
a 0.01 M ferrous perchlorate solution. The titrant was
0.1000 M EDTA and the electrode system consisted of a
platinum indicator electrode and a calomel reference.
Since ferrous ion is easily oxidized, fresh ferrous ion
stock solutions were prepared as needed from analytical
grade FeS0^*7H20 (Matheson Coleman & Bell, Cincinnati,
Ohio) and boiled, nitrogen purged water.

19
III» Ultraviolet Spectral Studies
Stability of thiouracll stock solutions» Aliquots
(2.0 ml.) of the aqueous 0.005 M stock solutions of 2-thio-
uracil, 6-n-propyl-2-thiouracil, 5.6-dimethyl-2-thiouracil,
2-ethylmercapto-4-hydroxypyrimidine and N,N -diethyl-6-meth-
ylthiouracil were diluted to 100 ml. with water and their
ultraviolet spectra recorded versus a water blank. Due to
the lower solubility of 5-nie'thyl-2-thiouracil and 6-methyl-
-2-thiouracil, 0.0025 M stock solutions were prepared and
4.0 ml. aliquots were diluted to 100 ml. with water and
the spectra recorded versus a water blank. All ultraviolet
spectra were recorded on a Cary Model 15 dual beam record¬
ing spectrophotometer using 1 cm. matched cells. The
balancing of the cells was checked periodically by running
blank versus blank on the spectrophotometer and in all
cases no significant difference in absorbance was found.
The ultraviolet spectra of the stock solutions were fol¬
lowed for a period of at least one month and in all cases
the change in spectra was less than 1%.
Spectra of thlouracil-metal ion solutions. The ul¬
traviolet spectra of water solutions containing 2-thioura-
cil, 6-n-propyl-2-thiouracil, 6-methyl-2-thiouracil,
5-methyl-2-thiouracil, 5.6-dimethyl-2-thiouracil,
i
2-ethylmercapto-4-hydroxypyrimidine and 6-methyl-n,N -di-
ethyl-2-thiouracil (1.00 x 10”^ M) and cupric, cadmium,

20
lead, ferrous and ferric ion (l.OOx 10“^ and 7,00x 10“^)
were recorded on the Cary ultraviolet recording spectro¬
graph versus a water blank (Table III). Air versus air and
water versus water blank curves were run in all cases and
net absorbances were negligible. The general procedure
was to record the spectrum of the metal nitrate alone
(1.0 ml. of 0.0100 M metal nitrate diluted to 100.0 ml.),
the spectrum of the ligand alone (2.0 ml. of 0.005 M ligand
or 4.0 ml. of 0.0025 M ligand each diluted to 100.0 ml.)
and the spectrum of a solution containing 1.0 ml. of 0.01
M metal ion and 2.0 ml. of 0.005 M ligand (4.0 ml. of 0.0025
M ligand) diluted to 100.0 ml. with water (Fi^. 2-6)
Another solution containing a seven-fold excess of metal
ion over ligand was also run. Since ferric ion has a large
-4
ultraviolet absorption, the spectrum of a 7.0 x 10 M
ferric ion solution was run for comparison. Precipitates
formed when cupric ion was mixed with 6-n-propyl-2-thioura-
cil and 5*6-dimethyl-2-thiouracil, causing the absorbance
to drop below the true value. In the case of 2-thiouracil,
6-methyl-2-thiouracil and 5-methyl-2-thiouracil no obvious
precipitate was present when cupric ion was added.
The ultraviolet spectrum of a cupric-2-thiouracil
solution as a function of hydrogen ion concentration was
also recorded. The five hydrogen ion concentrations ranged
—4
from 1.0 M to 1 x 10 M and the cupric ion concentration
-4 -4
was maintained at 2 x 10 M and 1 x 10 M in all cases.

21
The procedure was to add 10.0 ml. aliquots of the appro¬
priate HCIO^ concentration, 2.0 ml. aliquots of 0.005 M
2-thiouracil and 1.0 or 2.0 ml. aliquots of 0.01 M cupric
ion to a 100 ml. volumetric flask and dilute to the mark
with water. A solution containing HC10^ at the same con¬
centration as the sample was used as a blank.
Aqueous solubility of thiouraclls as a function of
hydrogen ion activity. Aqueous solutions (50 ml. in
100 ml. sealed flasks) of 2-thiouracil, 6-n-propyl-2-thi-
ouracil, 6-methyl-2-thiouracil, 5-methyl-2-thiouracil,
5,6-dimethyl-2-thiouracil, 5-carboethoxy-2-thiouracil and
6-amino-2-thiouracil were equilibrated with an excess of
solid ligand at 25.0° in a controlled temperature shaker
bath. Differing aliquots of 0.01 M NaOH and 0.1 M HCIO^
were included in the 50*0 ml. volume to vary the pH. The
pH values and the absorbances of the aliquots of the
equilibrated solutions were measured simultaneously as a
function of time to assure that equilibrium had been at¬
tained. Equilibrium conditions were established by repeat¬
ed sampling until no further change in absorbance was noted.
Filter sticks (Sargent #S-30417) were used to remove sam¬
ples from the equilibrated solutions. The first filtrate
was discarded and 2.0 ml. aliquots of the filtered, equi¬
librated solution were taken and diluted to 100.0 ml. with
0.1 M HC10^ and read on the Cary spectrophotometer

22
or Beckman DU versus a 0.1 M HCIO^ blank. Further dilu¬
tions were carried out as needed. The calculated solubil¬
ities are given in Table IV» Since thiouracils are easily
oxidized (14) the possible degradation of equilibrating
solutions of 2-thiouracil, 6-n-propyl-2-thiouracil, 5-meth-
yl-2-thlouracil and 6-methyl-2-thiouracil was checked.
Ultraviolet spectra of fresh solutions of the above com¬
pounds were recorded and the ratio of their absorbance to
absorbance of the equilibrated solutions calculated at
seven different wavelengths. The wavelengths chosen were
210, 220, 240, 260, 280, 300 and 320 mji. The lack of the
presence of appreciable oxidation products was indicated
by the constancy of the ratios for each curve.
Aqueous solubility of 2-thlouracil and 6-n-propyl-
-2-thiouracil as a function of pH in the presence of
cadmium ion. Sample solutions were prepared and treated
in exactly the same manner as in the previous section ex¬
cept that the total initial volume was carefully controlled
(50.00 ml.) and each flask (100 ml.) contained 5 ml• of
0.01 M or 25 ml. of 0.01 M cadmium nitrate. The final
total concentration of cadmium ion was 0.001 or 0.005 M.
The aliquots were suitably diluted with 0.1 M HCIO^ to
give an absorbance less than 1.00 on the Beckman DU when
measured at the .
Amax

23
Determination of molar absorption coefficients (£ ).
Accurately weighed amounts of dried samples of 2-thiouracil,
6-n-propyl-2-thiouracil, 6-methyl-2-thiouracil, 5-methyl-
-2-thiouracil, 5»6-dimethyl-2-thiouracil, 2-ethylmercapto-
-4-hydroxypyrimidine, 6-methyl-N,N *-diethy1-2-thiouracil,
the disodium salt of 2-thiouracil disulfide trihydrate, the
disodium salt of 2-sulphinyl-4-hydroxypyrimidine, 5-carbo-
ethoxy-2-thiouracil and 6-amino-2-thiouracil were dissolved
(see section on materials for solvents used) in an accu¬
rately known volume and the spectra recorded on the Cary
versus a solvent blank. The molar absorption coefficient
(Q was calculated from Eq. 2 where A is the observed
€ = (E<1* 2)
absorbance, 1 is the path length in cm. and C is the molar
concentration of the particular thiouracil. The results
are given in the section on materials.
Job^ plots of 2-thlouracil and 6-n-propyl-2-thlouracil
in the presence of cupric ion. Varying aliquots of 1.00 x
10’3 M ligand were added to aliquots of 1.00 x 10-3 M
cupric nitrate so that the total number of moles of ligand
plus cupric ion were kept constant at 1.00 x 10“3. The
—4
solutions were diluted to a final volume of 100 ml. (10
M in ligand plus cupric ion) with water in the case of
2-thiouracil and 0.05 M HC10^ in the case of 6-n-propyl-
-2-thiouracil. The ultraviolet spectrum of each solu¬
tion was recorded versus a water or 0.05 M acid blank,

24
respectively. The absorbance at 345 and. 272 mp (2-thioura-
cil) and 272 and 300 mji (6-n-propyl-2-thiouracil) were
plotted versus the mole fraction of metal. A straight
line joining the absorbance at mole fraction unity to the
absorbance at mole fraction zero gives the absorbance, at
any mole fraction, that is expected under the assumption
of no interaction, i.e., no complexation (Fig. 7). The
difference in absorbance between the constructed straight
line and the experimental points is proportional to the
amount of complex formed. The concentration of complex
formed, and therefore the difference in absorbance between
constructed line and experimental curve, will be at a max¬
imum when the mole fraction of ligand in the test solution
corresponds to the stoichiometric fraction of ligand in
the complex. For example, a 1:1 complex would have the
largest deviation from linearity at a mole fraction of
0.5 while a 2:1 complex of ligand to metal would have the
largest deviation at a mole fraction of 0.66. Therefore,
the deviations from linearity were plotted versus mole
fraction of ligand and a smooth curve drawn through the
points (Fig. 8). If the complex absorbs at a wavelength
where the uncomplexed ligand does not, then it is not
necessary to construct a straight line and a plot of the
observed absorbances versus mole fraction will indicate
the stoichiometry as described above. This is the case
for 2-thiouracil at 345 mp. (Fig. 8).

25
In one case (2-thiouracil) the concentration of ligand
was held constant at 5«00 x 10“^ molar and the cupric ion
concentration varied from 5*00 x 10“^ M to zero. The
solutions were made up to 100 ml. with water and their
spectra recorded. The absorbances at 272 and 3^5 mp were
plotted versus the ratio of the molar concentrations of
cupric ion to 2-thiouracil. When the ratio of the molar
concentration of cupric ion to 2-thiouracil slightly ex¬
ceeds the ratio of metal ion to ligand in the complex then
the absorbance is constant for strong complexes. At ratios
of metal ion to ligand less than the stoichiometric ratio,
the observed absorbance will deviate from the constant
value (Fig. 9)«
Spectrophotometric titration of 2-thlouracil. Thirty
ml. solutions of 0.921 x 10“^ and 1.00 x 10“^ M in 2-thio-
uracil and 8.00 x 10 J M in sodium perchlorate contained
in a water-jacketed titration cell (25»0°) were titrated
with 1.00 N NaOH and 0.0499 N NaOH, respectively. Nitrogen
gas, free of carbon dioxide (passed through an Asearite
tube) and saturated with water (passed through a sparger
tube immersed in water) at 25*0° was passed into the cell
for each run. The cell solution was stirred with a mag¬
netic stirrer. A constant rate burette (Sargent Co.)
equipped with a 2.5 ml. syringe burette was used to
deliver the titrant to the cell. The volume delivered

26
could be read to the nearest 0. 5 )il. A micro-aperture
flow cell (Beckman Co., catalog #97290) was connected by
polyethylene tubing (Clay-Adams Co., #PE 200) to the
titration cell and to a 5»0 ml. gas-tight syringe (Hamilton
Co.). After each addition of standard alkali the gas-tight
syringe was actuated by hand to draw the titration cell
solution into the flow cell and the spectrum recorded
versus a water blank. The gas-tight syringe was actuated
several times before each spectrum was recorded to insure
thorough mixing of any solution that may have remained in
the flow cell or tubing. The pH of the solution was read
after mixing using a glass-calomel electrode system and a
Radiometer pH meter (#TTT 1). The total volume change
during the titration of the sample solution was less than
2% and was considered negligible. The absorbances at 230.
240, 2?0, 310, 320 and 285 m)i were plotted versus pH and
the apparent pK determined from the pH at half-neutraliza-
tion according to Eq. 3 (40) (Figs,ID, 11). The parameters
PK* = pH - log A- (Eq. 3)
a A ~ Aalk.
f
in Eq. 3 have the following significance; the pK is the
St
negative logarithm of the apparent dissociation constant;
the pH is the negative logarithm of the hydrogen ion activ¬
ity; Aacid is the absorbance of 2-thiouracil in dilute acid
(0.10 M HCIO^); A is the absorbance of 2-thiouracil solution

27
in the buffer region; A is the absorbance of 2-thioura-
cll solution in the alkaline region (pH 9)»
Spectrophotometric titration of 2-thiouracil-cuprlc
ion mixtures» The procedure and apparatus used for the
titration of 2-thiouracil-cupric ion mixtures was exactly
as outlined in the previous section. The 30*0 ml. of test
_ k
solution was composed of 10.00 ml. of 3*00 x 10 M
_k
2-thiouracil, 2.00 ml. of 3*00 x 10 M cupric ion and
18.0 ml. of 0.008 M sodium perchlorate. Both the 2-thi¬
ouracil and cupric ion solution were prepared using 0.008
M sodium perchlorate as diluent in order to maintain con¬
stant ionic strength. Concentrations of cupric ion
_k
0.5 x 10 and larger caused precipitation which frustrated
analysis. The recorded ultraviolet spectra (Fig.12 ) was
analyzed using Eq. 3 at 270. 285, 230, 240, 310 and 320 mji
(Fig. 13).
Ultraviolet spectra of 2-thiouracll and cupric ion
in dilute acid. The ultraviolet spectrum of a 1:50 dilu¬
tion of a solution containing 115 mg. of 2-thiouracil,
217 mg. of cupric nitrate trlhydrate and 20.0 ml. of 1.0
M HCIO^ diluted to 200.0 ml. was recorded versus a 0.05
M HC10^ blank. The final concentrations of 2-thiouracil,
cupric ion and perchloric acid were 9.0 x 10“^, 9.0 x 10“^
and 0.05 molar, respectively. The 1:50 dilution was made
using 0.05 M HCIO^ as diluent. Another solution was

28
prepared as above using 10.0 mg. of cupric nitrate trihy¬
drate, 115 mg. of 2-thiouracil and 20.0 ml. of 1.0 M HCIO^
and its ultraviolet spectrum recorded versus a 0.05 M HC10^
blank. The final concentration of 2-thiouracil, cupric ion
and perchloric acid were 9*0 x lO-^, 4.3 x 10”^ and 0.05
molar, respectively. The 2-thiouracil and cupric nitrate
trihydrate were weighed on a torsion balance and are accu¬
rate to + $%.
Degradation of 2-thiouracll disulfide in absence and
presence of cupric ion. A solution (1.999 x 10“24' M) of
the disodium salt of 2-thiouracil disulfide monohydrate
was prepared by dissolving 6.322 mg. in 100 ml. of water.
Aliquots (2.0 ml.) of the disulfide stock solution were
added to: A, 3»0 ml. HgO; B, 2.0 ml. HgO and 1.0 ml. of
—4
4.00 x 10 M cupric nitrate; C, 1.0 ml. HgO and 2.0 ml.
-4
of 4.00 x 10 M cupric nitrate. Each solution was pre¬
pared when needed and then 5»00 ml. of 0.1 M HC10^ was
added very rapidly from a $.00 ml. syringe (Hamilton),
mixed, and the change in absorbance recorded continuously
at 2?0 mp. The ultraviolet absorption was measured using
1 cm. cells and a 0.1 M HC10^ blank. A water jacketed
(25*0°) Beckman DU, fitted with an automatic sample changer
(Gilford) and connected to a strip chart recorder (Sargent
SRL) was used to record the absorbance. The absorbance of
all three solutions was recorded until no further change
occurred

29
IV. Polarographic Studies
Apparatus. The polarographic studies were performed
on a Sargent Model XV Recording Polarograph equipped with
a water jacketed H cell (Sargent Co.) maintained at 21.0,
30.0, 35«0, 40.0 or 45.0°. All potentials were measured
versus a saturated calomel electrode and were corrected
for potential error due to recorder current. The possible
sources of error in potential measurements using the Model
XV Polarograph are the recorder current measuring resistor,
the bridge resistance, cell resistance and the damping
resistor if used. The potential error due to recorder
current can be easily corrected for since the recorder
on the Model XV always requires 2.5 mv for full scale
deflection (41). The error in any potential reading was
corrected by simply reading the pen position at the desired
point, assuming that the scale is 0.0 to 2.5 mv and not
microamperes. The error due to the bridge resistance can
be found by multiplying the bridge resistance (200 ohms
at worst) by the observed current (1.0 to 2.0 ya). This
gives a maximum error of 0.4 mv which is negligible. No
damping resistance was used in these studies and therefore,
potential errors due to damping do not exist. The cell
resistance was measured frequently using a Radiometer con¬
ductivity meter (type 2 d) operating at 3000 cycles per
second. The reciprocal of the conductivity in mhos gives

30
the cell resistance in ohms. The resistance in all cases
was 300 ohms or less which is equivalent to 0.6 mv at 2 pa.
The 0.6 mv is a negligible error. The instrumental accu¬
racy for potential measurements, as quoted by the Sargent
Co. is +2.5 mv on a 1.0 volt scale (41). An agar plug
saturated with potassium nitrate (42) was used in the salt
bridge and was changed whenever the solid potassium nitrate
became depleted. The potassium nitrate was necessary since
the usual potassium chloride interfered dvurto the anodic
wave of chloride ion. All test solutions were deoxygenated
with water-saturated nitrogen gas (passed through a sparger
immersed in water) for at least ten minutes and were main¬
tained in a nitrogen atmosphere (nitrogen gas flowing over
the solution) during the analysis. The capillary constant
(m¿/^ t /D) was determined by running a supporting electro¬
lyte solution (0.2 M NaClO^) for a long, accurately mea¬
sured period of time (2,042.8 sec.), weighing the mercury
delivered (4.9185 gm.) and counting the number of drops
(determined from the number of pen oscillations. The val-
1 /6 2 /3
ues of t ' and m were calculated from Eqs. 4 and 5.
respectively.
(Eq. 4)
Total Sec.
(Eq. 5)

The value of m2/^ t1^ for capillary number 2 is 2.215
The height of the mercury column was held constant at
31
58.5 cm. The current and voltage were always standardized
before each curve was run.
Polarography of cupric ion solutions. Aliquots
(3*00, 4.00 and 5*00 ml.) of a 0.0100 M cupric nitrate
solution (prepared in 0.2 M NaClO^) were added to 0.25
ml. of 0.2$ Triton X-100 (Rohm and Haas Co.) and 5»00 ml.
of standardized HC10^ in a 50 ml. volumetric flask and
diluted to the mark with 0.2 M NaClO^. These solutions
were purged with nitrogen and the polarograms recorded
on the Model XV at 21.0°. The final perchloric acid
concentrations were 0.2648, 0.100, 0.0100, 0.00100 and
0.000100 molar. The final concentration of Triton maxi¬
mum suppressor (0.001$) is that recommended in the
literature (43-45) and was found to give reversible waves
as determined by Eq. 6 (46-49).
log (Eq. 6)
The parameters in Eq. 6 have the following significance;
Eis the potential at 3A of* i¿ where i^ is the observed
diffusion current in pA; E^y^ is the potential at lA of
i^; R is the gas constant in cal./deg.-mole; T is the
absolute temperature; F is the faraday (23.060.3 cal./abs.

32
volt gram eq.); n is the number of electrons involved, in
the electrode reaction. The values of -2.303 R T log 9/F
at different temperatures are; 0.0557, (21°); 0.0573 (30°);
0.0583 (35°); 0.0592 (40°); 0.0602 (45°). The derivation
of Eq. 6 is given in the section on equations.
Another series of solutions were prepared containing
1.00, 2.00, 3.00, 4.00 and 5*00 ml. aliquots of 0.0100 M
cupric nitrate solution (prepared in 0.2 M NaClO^), 0.5
ml. of 0.2$ Triton, 10.0 ml. aliquots of standardized HCIO^
and sufficient 0.2 M NaC10¿^ to dilute to 100 ml. The final
concentrations of HCIO^ varied from 0.2648 to 0.000100 as
before. All solutions were purged with nitrogen gas to
remove oxygen as previously described and the polarograms
recorded on the Model XV at 21.0°. All diffusion currents
were measured at the half-wave potential. The peak of the
pen oscillation was used to calculate all potentials and
currents (50)» Details of the solution composition, half
-wave potentials, values and diffusion currents
are summarized in Table V.
Polarography of cupric ion-thiouracil mixtures.
Solutions containing 1.0 ml. of 0.01 M cupric ion, 0.25 ml.
of 0.2$ Triton, 5 ml. aliquots of standard HCIO^ (2.648,
2.00, 1.000, 0.800, 0.500, 0.100, 0.0100 and 0.00100 M)
and varying volumes of 0.005 M or 0.0025 M thiouracil
derivative (2-thiouracil, 6-n-propyl-2-thiouracil,

33
6-methyl-2-thiouracil, 5-Bie'fchyl-2-thiouracil, 5.6-dimethyl-
-2-thiouracil, 5-carboethoxy-2-thiouracil) were diluted to
50 ml. with 0.2 M NaClO^. The 0.005 M and 0.0025 M thi-
ouracil solutions were prepared in 0.2 M NaClO^ and as the
volume of thiouracil solution was decreased an equal volume
of 0.2 M NaClO^ was substituted to maintain the ionic
strength at 0.2 molar. The solutions were purged with
nitrogen as previously described and the polarograms re¬
corded at controlled temperatures (21.0, 30.0, 35*0, 40.0
45.0°). A 1.00 ml. aliquot of 0.2 M NaClO^ was substituted
for the cupric ion in the blank solutions which were run
for each test solution. The curve for the blank was sub¬
tracted from the test solution curve and the half-wave
potential, E^/^-E^y^ value and the diffusion current
determined from the difference curve (Fig. 6 ). Each test
solution was run at least three times and the potential
and current measurements averaged. The peak of the pen
excursion was used to calculate all potentials and currents
(50). A typical polarogram is shown in Fig.14 and details
of typical solution composition, half-wave potentials,
values and diffusion currents are given in Table
VI

34
Polarography of solutions containing cuprous ion in
the presence of 2-thlouracll and 2-thiouracll disulfide.
Cuprous chloride was purified by dissolution of 5 gm* of
CuCl (light green in color) in 100 ml. of concentrated
hydrochloric acid to give a black solution containing
CuCl". Addition of 300 ml. of water gave a light green
solution from which white crystals of pure CuCl precipi¬
tated (51)» The following qualitative tests were run on
the precipitated CuCl. A solution of ammonium hydroxide
(28$) added to the CuCl gave a colorless solution which
+2
slowly turned dark blue (Cu(NH^)2 ). Concentrated HC1
added to the CuCl gave a colorless solution (CuClg) which
rapidly turned dark. Addition of potassium iodide to the
colorless CuClg solution produced a white precipitate
(Cul).
Polarograms of the following solutions were run as
previously described. Solution A contained 1.979 mg. CuCl,
1.0 ml. of 0.2 M NaClO^, 0.25 ml. of 0.2$ Triton, 5.0 ml.
of 2.648 M HCIO^ and 43*75 ml. of 0.005 M 2-thiouracil
(in 0.2 M NaClO^). Solution B contained 1.554 mg. 2-thi¬
ouracil disulfide, 0.25 ml. of 0.2$ Triton, 5*0 ml. of
2.648 M HC104 and 44.75 ml. of 0.2 M NaClO^. Solution C
contained 1.979 mg. CuCl, 3*148 mg. 2-thiouracil disulfide,
10 ml. of 2.648 M HCIO^, 0.5 ml. of 0.2$ Triton and 89.5
ml. of 0.2 M NaClO^.

35
Polarography of presumed cuprous complex of 2-thioura¬
cil. The cuprous chloride complex of 2-thiouracil (IX)
was synthesized according to the literature (26). The
procedure for the synthesis is the dissolution of 1.3 gm.
of 2-thiouracil in 150 ml. of hot water and 40 ml. of 1.0
M CuCl in concentrated HC1 solution was added. Yellow
crystals formed which were washed with water and then ace¬
tone and dried in a vacuum over at 50°•
Two solutions of IX were prepared. The first solution
contained I.883 mg. IX, 0.25 ml. 0.2$ Triton, 5*0 ml. of
3.0 M HCIO^ and was diluted to 50 ml. with 0.2 M NaClO^.
The second solution contained I.876 mg. IX, 0.25 ml. 0.2$
Triton, 5.0 ml. of 3,0 M HCIO^ 14.75 ml. of 0.005 M
2-thiouracil (in 0.2 M NaClO^) and was diluted to 50.0
ml. with 0.2 M NaClOj^. Both solutions were run on the
polarograph at 21.0°. The solutions were purged with
nitrogen as usual and a blank solution containing every¬
thing but IX was run in both cases.
V. Synthesis of Complexes
Synthesis of bis(2-thiouracll)cadmlum(II). A solu¬
tion containing 0.06 mole of 2-thiouracil in 2 liters of
hot water was prepared. To the 2-thiouracil solution was
added slowly, with stirring, a solution containing 0.03
mole of cadmium nitrate in about 200 ml. of water. The
resulting mixture was allowed to stand 0.5 hour on low

36
heat (about ?0°). The solution was allowed to cool to room
temperature and the pH adjusted to 6.5 with concentrated
NaOH. The resulting suspension was warmed again, cooled
to room temperature and filtered through a medium glass
fritted funnel. The product was washed with cold water,
acetone and dried in a vacuum oven at 60°. I. R. spectrum,
Din cm.”1 (Nujol mull): 1620, 1550, 1530, 1300, 1210,
1180, 82?.
Anal. Caled, for CgH^N^OgSgCd: Cd, 30.65.
Found: Cd, 31.54, 30.96.
Synthesis of bis(2-thiouracll)lead(II). The pro¬
cedure for the synthesis of bis(2-thiouracil)lead(II) is
exactly the same as in the case of bis(2-thiouracil)cad-
mium(II) except, that lead nitrate was used in place of
cadmium nitrate. I. R. spectrum, Din cm.”1 (Nujol mull):
1660, 1630, 1560, 1500, 1280, 1000, 815.
Anal. Caled, for CgHgN^OgSgPb: Pb, 44.90.
Found: Pb, 44.51.
Synthesis of bis(2-thlouracil-cadmlum(II)). A solu¬
tion containing 0.06 mole of 2-thiouracil in a minimum of
hot water was added slowly, with stirring, to a one liter
solution containing 0.06 mole of cadmium nitrate heated
to the same temperature (about 80°) as the hot thiouracil
solution. The resulting solution, which may contain a
suspension of product, was maintained at the initial

37
temperature for \ hour and then allowed to cool to room
temperature. Concentrated sodium hydroxide was added to
pH 6.8 and the resulting suspension reheated. After \
hour the suspension was cooled to room temperature and
the product filtered through a medium glass fritted flannel.
The resulting product was washed with water and then ace¬
tone and dried in the vacuum oven at 60°. I. R. spectrum,
Din cm."1 (Nujol mull): 3400, 1570, 1510, 1335, 1020.
Anal. Caled, for CgH^N^OgSgCdg: Cd, 47«2.
Found: Cd, 45.8.
Synthesis of bis(2-thiouracll-lead(II)). The pro¬
cedure for bis(2-thiouracil-lead(II)) is exactly the same
as in the case of bis(2-thiouracil-cadmium(II)) with lead
nitrate substituted for cadmium nitrate. I. R. spectrum,
0 in cm."1 (Nujol mull): I56O, 1520, 1430, 1330, 1000,
820.
Anal. Caled, for CgH^N^OgSgPb: 62.15*
Found: Pb, 62.54.
Synthesis of bis(6-n-propyl-2-thlouracll)cadmium(II).
The preparation of this complex was performed in the
same manner as for bis(2-thiouracil)cadmium(II). The molar
amounts of 6-n-propyl-2-thiouracil and cadmium nitrate were
0.06 and 0.03, respectively. I. R. spectrum, Uin cm."1
(Nujol mull): 3100, I63O, 1500, 1270, 1220, 1175, 1015,
970, 830.

38
Anal. Caled, for C^H^N^O^Cd: Cd, 24.9; C, 37.29;
H, 4.02; N,12.43; S, 14.22.
Found: Cd, 24.7; C, 37-94; H, 4.15; N, 11.95; S,
13.77.
Synthesis of bis(6-n-propyl-2-thiouracll-lead(II)).
This complex was prepared by the same procedure as for
bis(2-thiouracil-cadmium(II)). The molar amounts of
6-n-propyl-2-thiouracil and lead nitrate used were 0.06.
I. R. spectrum, Din cm.'"^ (Nujol mull): 1550* 1420, 1280,
1165, 1020, 825.
Anal. Caled, for C^H^N^O^Pb: Pb, 55*2.
Found: Pb, 55*4.
Synthesis of bis (2-thlouraoll )-p-dlhydroxodicopper (II).
A solution containing 0.062 mole of cupric ion in 100 ml.
of water was added slowly, with stirring, to 400 ml. of a
warm, aqueous solution (about 1.0 liter) containing 0.062
mole of 2-thiouracil. Very fine, light-yellow crystals
slowly formed which were removed by filtration, washed
with water and acetone and dried at 120° overnight. The
elemental analysis was done on material dried to constant
weight. Ultraviolet assay of 2-thiouracil content gave
62.2$. Calculated value for c8H8°4N4S2Cu2 is 61.7$. The
procedure for the assay was to put an accurately weighed
amount of complex into hot 1.0 M HC10^ for one day and
then read the ultraviolet spectrum versus a 1.0 M HCIO^

39
blank. The concentration of 2-thiouracll was determined
from the molar absorption coefficient (£ 13.700) at 273
m^i. I. R. spectrum, 0 in cm.~^ (Nujol mull): 3090, 1640,
1600, 1540, 1280, 1160, 1070, 1015, 825. Weight loss when
dried at 100° under vacuum for two weeks was 4.57$. Cal¬
culated for CgHgO^SgCUg
Anal. Caled, for CgHgN^OgS^Ug: C, 24.18; H, 1.52;
N, 14.10; S, 16.14; Cu, 31*98.
Found: C, 24.54, 24.06; H, 1.62, 1.66; N, 14.33;
S, 19*48, 17.51; Cu, 32.84.
Synthesis of bis(6-n-propyl-2-thlouracll)-p-oxodl-
copper(II). The procedure for the preparation of bis(6-
-n-propyl-2-thiouracil )-ji-oxodicopper (II) is exactly the
same as for the 2-thiouracil case. The yellow precipitate
was dried to constant weight at 120°. The procedure for
the spectrophotometric analysis of the percent of 6-n-pro-
pyl-2-thiouracil is the same as for 2-thiouracil. I. R.
spectrum, Üin cm.“^ (Nujol mull): 3040, 1640, 1540, 1490,
1440, 1270, 1210, 1170, 1020, 955. 835.
Anal. Caled, for ci4%8°3N4S2Cu2: Cu, 26.39; 6-n-
-propyl-2-thiouracil, 70.7*
Found: Cu, 26.82; 6-n-propyl-2-thiouracil, 70.6.

40
Analytical procedure for cadmium content of complexes»
Accurately weighed (200 mg.) samples of the dried cad¬
mium complex were dissolved In 150 ml. of 0.01 M HgSO^ and
heated (70°) until all of the sample had dissolved. Ap¬
proximately 150 mg. of sodium hydrogen sulfide in 15 ml.
of water was added and the resulting cadmium sulfide pre¬
cipitate digested for about two hours until the crystals
were large. The precipitate was filtered onto a tared
glass fritted funnel, washed with warm water and dried
in a vacuum oven at 5°° overnight. The dried precipitate
was weighed in the funnel and the weight of cadmium sulfide
determined by subtraction. The weight of cadmium was cal¬
culated by multiplying the weight of cadmium sulfide by
the gravimetric factor O.778O. The percent cadmium was
calculated from the ratio of cadmium weight to sample weight
multiplied by 100. This analytical procedure is based on
the literature method (52).
Analytical procedure for lead content of complexes.
Accurately weighed samples (300 mg.) of the lead complex
was digested in 100 ml. of 1.0 M nitric acid until every¬
thing had dissolved. The volume was reduced to 25 ml* by
evaporation and 100 ml. of a solution 2.0 M in sulfuric
acid and 1.0 M in sodium sulfate was added. The precipi¬
tate was digested to give large crystals and the volume
reduced by evaporation to about 50 ml. The precipitate

41
was collected on a tared, fritted funnel, washed with water
and dried at 120° overnight. The weight of lead was deter¬
mined as in the cadmium case. The gravimetric factor is
0.6832. This analytical procedure is based on the litera¬
ture method (53)»
VI. Animal Experiments
Determination of antithyroid activity of 5.6-dlmethyl-
-2-thiouracil. Twenty white Carworth CFN rats were evenly
divided into four groups with five rats per group. One
group was used as a control and was fed Rockland Rat diet.
The other three groups were fed the same diet containing
0.025» 0.050 and 0.100 weight-percent 5»6-dimethyl-2-thio-
uracil for a period of two weeks. Both diet and water
were supplied ad libitum. The doses were chosen with
the hope that a dose response relationship between drug
dose and thyroid weight could be obtained. Body weights
were measured just before sacrifice. At death weights
of thyroid glands were determined on a torsion balance. The
total amount of diet ingested was not measured but an esti¬
mate of 3*86 gm. of diet per 100 gm. of body weight per day
was made from the 5-Methyl-2-Thiouracil study. The effect
of the drug on thyroid weights is given in the section on
results and the thyroid weights are given in the results
and in Table VII

42
Determination of antithyroid activity of 6-methyl-
t
-N,N -dlethyl-2-thiouracil. Ten white, male Carworth
CFN rats (120-200 gm.) were divided into two groups of
five rats each. One group was used as a control and was
given the standard diet. The other group was given the
Rockland Rat diet containing 0.100$ 6-methyl-N,N -diethyl-
-2-thiouracil. The amount of drug ingested per day was
determined by weighing the diet supplied each day and
then weighing the amount uneaten at the next feeding
period. The rats were given the diets for a period of
two weeks and both diet and water were supplied daily
as needed. At the end of two weeks the animals were
sacrificed and the thyroid glands carefully removed and
weighed on a torsion balance. The thyroid weights are
given in the section on results and in Table VIII.
Determination of antithyroid activity of 6-n-propyl-
-2-thlouracll and 5-methyl-2-thiouracll> The procedure
for the determination of the antithyroid activity of 6-n-
-propyl-2-thiouracil and 5-niethyl-2-thiouracil is exactly
the same as for 5.6-dimethyl-2-thiouracil except the

43
daily consumption of diet per 100 gm. of rat body-weight
was measured. The thyroid weights and consumption of
diet are given in Table IX.
VII. Magnetic Susceptibility Measurements
«
The corrected molar magnetic susceptibility,XM» of
the cuprous-2-thiouracil complex was measured by the Chem¬
istry Department, University of Florida and gave a value
_ ¿L
of -1.10 x 10 , a negative number. Since ji, the magnetic
1
dipole moment, is proportional to the square root ofX^, a
negative number, then the magnetic dipole moment is most
probably zero and the isolated complex is probably the
cuprous complex of 2-thiouracil. The absolute value of
is rather large.

EQUATIONS
I» Potentlometrlc Titrations
Derivation of equations for stablllllty constants
assuming the presence of only MU+ and MUg. The formation
of 1:1 and 2:1 complexes could occur as shown In Scheme II
for our specific compounds.

45
The derivation of the general equation for the calculation
of stability constants from potentiometric titration data
was first given by Bjerrum (54). An excellent summary of
computational techniques is also available (43). The
following derivations do not assume the formation of the
hydroxides of the free metal ion and require all species
to be in solution and at instantaneous equilibrium. The
highest pH at which usable data could be obtained in our
studies was about 6.5 since the precipitation which occurred
destroyed the equilibrium conditions. Hydrolysis constants
from the literature (55) permitted the estimation of values
for maximum hydrolysis of cadmium (0.2$) and lead (3*5$) at
pH 6.5. At pH values above 6.5 the loss of free cadmium
and lead could not be considered negligible. The apparent
acid dissociation constant of the thiouraclls is:
K' = .,[»-] [h+3 £+ (E 7)
3 [HU]
where [u“] and [HU] are the molar concentrations of
thiouracil anion and free acid respectively, [H+] + is
the hydrogen ion activity, and where ){+ is the mean activ¬
ity coefficient. The stability constant for the first
complex of metal ion with thiouracil is:
r [MU+]
Kl~ [K+2] [U-]
(Eq. 8)

where [MU'] is the concentration of the first complex and
[M+2] is the concentration of free metal ion. The step
stability constant for the formation of the 1:2 complex
46
is:
K = CMU2]
[MU+] [IT]
(Eq. 9)
where [MUg] is the concentration of the second complex.
The overall stability constant, the product of and Kg,
is:
[MUg]
= K. K„ = —x 7$
2 1 2 [m+2] [u-]2
(Eq. 10)
The mass balance equations for thiouracil, metal ion and
sodium hydroxide titrant are:
[HU]0 = [U-] + [HU] + [MU+] + 2[MUg] (Eq. 11)
[M+2]0 = [M+2] + [MU+] + [MUg] (Eq. 12)
[NaOH] = [U~] + [MU+] + 2[MUg] + [0H~] - [H+] (Eq. 13)
The initial stoichiometric concentrations of thiouracil and
metal ion are given by [hu]q and [m+2]q. The stoichio¬
metric concentration of alkali at any point in the titra¬
tion is given by [NaOH]. The concentration of hydroxyl
ion in Eq. 13 is the sum of the hydroxyl ion from the

47
titrant and from the dissociation of water. Since a hydro¬
gen ion is produced when a water molecule dissociates, the
hydrogen ion concentration corrects the hydroxyl ion con¬
centration for this phenomenon so that the resultant equa¬
tion accounts for the hydroxyl ion due to the titrant.
The degree of formation, ñ, is defined as the average
number of ligands bound to a metal ion.
[MU+] + 2[MUj
ñ = —
[m+2]0
(Eq. 14)
When the appropriately rearranged Eqs. 7 and 11 are sub¬
stituted into Eq. 14,
n
(Eq. 15)
When Eq. 13 is subtracted from Eq. 11, substitution of the
rearranged Eq. 7 for [HU] where the relatively small quan¬
tities [OH""] and [H+] are ignored,
[HU]q - [NaOH]
[H+] X'+ZK^
[U-] =
(Eq. 16 )

48
The right hand side of Eq. 16 contains only experimental
quantities and therefore, [U-] can be calculated. Sub¬
stitution of [U-] into Eq. 15 allows the calculation of
ñ.
Substitution of the rearranged equilibrium expressions
for [MU+], [MU2] (Eqs. 8 and 10) and the mass balance (Eq.
12) for [M+2]0 into Eq. 14 gives, on simplification and
rearrangement, a relation between ñ, and 0^,
n
(l-H)[u“]
?-j e2[u-] + K:
(Eq. 17)
Equation 17 is linear with a slope of 02 and an Intercept
of . If 02 is assumed to be zero Eq. 17 reduces to an
equation whose logarithmic transformation is:
log 1-n = pK. + p[tT] (Eq. 18)
ñ
where pK^ and p[U-] represent the negative logarithm of
and [U-], respectively. Equation 18 is a linear equa¬
tion with a slope of one and an intercept of pK^.
Derivation of equations for stability constants as¬
suming the presence of only MU+ and MUOH. The formation
of MUOH, the mixed ligand complex (XV), can conceivably
occur by reaction of XIII with OH-.

^9
XV
OH
MU OH
This derivation of the equations for the calculation of
the stability constants of XV assumes that no significant
amounts of MOH+ and/or MUg are formed and all species are
in solution and at equilibrium.
The expressions for the apparent acid dissociation
t
constant, K and the equilibrium constant, K1 , for the
a
first complex, MU+, have already been given (Eqs. 7, 8).
The step stability constant, Kg, of the mixed ligand
complex is:
, [mu oh]
Kg = (Eq. 19)
[MU+] [OH-] y +
where [MUOH] is the molar concentration of the mixed ligand
complex and [0H“] %+ is the activity of hydroxyl ions. The
mean activity coefficient is given by Y+. The expression

50
for the autoprotolytic constant of water is:
Kw = [H+]%+ [OH-] K+
(Eq. 20)
The overall stability constant, the product of Kw, and
Kg, for the mixed ligand complex is:
[MUOH] [H+]
[M+¿] [u“]
(Eq. 21)
The value of [0H~] y+ has been substituted by Kw/[H+]
from Eq. 20. The mass balance equations for thiouracil,
metal ion and sodium hydroxide titrant are:
[HU]0 = [U“] + [HU] + [MU+] + [MUOH] (Eq. 22)
[M+2]0 = [M+2] + [MU+] + [MUOH]
.+2
(Eq. 23)
[NaOH] = [IT] + [MU+] + 2[MU0H] + [0H~] - [H+] (Eq. 24)
The degree of formation for the mixed lagand complex is:
[MU+] + [muoh]
n =
[m+2]0
(Eq. 25)

51
Substitution of rearranged Eqs. 7 and 22 into Eq.
suits in:
o - CD'J
[m+2]0
/ [h+] y+
i + 1—
25 re-
(Eq. 15)
Subtraction of Eq. 24 from Eq. 22, substitution of [HU]
from Eq. 7, [MUCH] from Eq. 21 and dropping the relatively
small quantities [0H~] and [H+] gives:
[HU]0 - [NaOH]
= [IT] (Eq. 26)
[H+]^+/k¡ - 311[M+2]/fH+]y+
The left hand side of Eq. 26 is a function of the free
metal ion concentration and therefore, the value of [u“]
cannot be accurately calculated unless the free metal ion
concentration is determinable. If the assumption is made
that MUOH is not present in any significant amount during
some interval in the titration then Eq. 26 reduces to Eq.
16. This assumption would permit the observed data to
conform to Eq. 18.
The relation of ñ to and for mixed ligand com¬
plexes is derived by substitution of the equilibrium ex¬
pressions for [MU+] (Eq. 8), [MUOH] (Eq. 21) and [M+2]
(Eq. 23) into Eq. 25 to give:

52
Kl[u-] + hi —
[H+] +
n = (Eq. 2?)
1 + K^IT] + Pn[U"]/[H+] x±
Rearrangement of Eq. 27 by multiplication of n by the
denominator and collection of similar terms gives:
n = K + — (Eq. 281
(1 - S)[ir] 1 [H+]ir+
which is linear with a slope of g^ and an intercept of
. When it is assumed that the concentration of MUOH
is not significant, g^ approaches zero and the logarithmic
transformation of Eq. 28 reduces to Eq. 18.
Derivation of equations for stability constants as¬
suming the presence of MU+, MU and ftUU,-,. The formation
of polynuclear complexes (MnUn) can conceivably occur as
shown in Scheme III*

53
XVI MU
XVII m9u,
Scheme III ¿ '

Molecular models of XVII are easily formed and exhibit no
strain. It is also possible to rewrite XVII in the linear
form
54
XVIII KnUn
The assumptions for the derivation of the equations for
polynuclear complexes are the same as for simple and mixed
ligand complexes, that is, no significant amounts of MOH+,
MU2 and/or MUOH are formed and all species are in solution
and at equilibrium.
The expressions for the apparent acid dissociation
constant of the ligand and the stability constant of the
first complex (MU+, XIII) have been given (Eqs. ?, 8). The
expressions for the acid dissociation constant of MU+ (XIII)
and the step stability constant of MgU2 (XVII) are
[mu] [h+] y+
[MU+]
(Eq. 29)

55
(Eq. 30)
Substitution of the equilibrium expression for [MU+]
(Eq. 8) into Eq. 29 and. collection of constants gives
Ka2Kl
[MU] [H+]¿r+
[M+2] [U~]
(Eq. 3D
Substitution of the expression for [MU] from Eq. 31 into
Eq. 30 and collection of constants gives
P
22 =
[m2u2] [h+]22T+2
[m+2]2 [u~]2
(Eq. 32)
The mass balance equations for thiouracil, metal ion and
sodium hydroxide titrant are
[HU]0 = [U-] + [HU] + [MU+] + [MU] + 2[M2U2] (Eq. 33)
[M+2]0 = [M+2] + [MU+] + [MU] + 2[M2U2] (Eq. 34)
[NaOH] = [U“] + [MU+] + 2[MU] + 4[M2U2] + [0H~] - [H+]
(Eq. 35)
The degree of formation for the specific polynuclear com¬
plex M0U0 is
[MU+] + [mu] + 2[M2U2]
C«+2]0
n
(Eq. 36)

56
Equation 36, when substituted by rearranged Eqs. 7 and 33
(Sq. 15)
C«+2]0
Subtraction of Eq. 35 from Eq. 33» substitution of Eq. 7,
31 and 32 for [HU], [mu] and [M2U2], respectively, and
dropping the relatively small quantities [CH~] and [H+]
gives
[HU]0 - [IT]
1 +
[H+]
K
n =
[HU]q - [NaOH]
[H+] X+fc'a - 0a2[M+2]/tH+] - 2P22[M+2]2[u']/tH+]2^+2
(Eq. 37)
The general equations for the ligand concentration in the
case of polynuclear complexes MpUp, where p must equal q
as in Scheme III, would be given by
[u-] =
[HU]n - [NaOH]
P Q
[h+]
0 1
P gpq[M+2]q[U~]
-IP-1
[H+]P ^±P
(Eq. 38)
When p = q = 1 Eq. 38 reduces to Eq. 26 and when p = 0
it reduces to Eq. 16. The left hand sides of Eqs, 37 and
38 are functions of the free metal ion concentration and
therefore, the value of U~ cannot be accurately calculated

57
unless the free metal Ion concentration is known. If the
assumption is made that MpUp is not significant during
some Interval in the titration then Eqs. 37 and 38 reduce
to Eq. 16 and permits the data to be plotted according to
Eq. 18.
The relation of H to , ga2 and @22 is derived by-
substitution of the equilibrium expressions for [MU+]
(Eq. 8), [MU] (Eq. 3D, [M^] (Eq. 32) and [M+2]q into
Eq. 36 to give
K^U"] + ga2[U~][H+]~1 yjf1 + 2g22[M+2][U~]2[H+]~2
1 + k1[u~] + ea2[u"][H+]"1 ¿f+_1 + 2322[m+2][u~]2[h+]"2 ¿r+"2
(Eq. 39)
The general equation for ñ in the case of MU+ in the pres¬
ence of MU is
q. P
Q P
Ki[u~*] + YL Y. ppqpCM+2^q"1Cu"]p[H+]"p¿r+_p
D4
kiCu_] + Y- «ipqp[M+2]q“1[^]pCH+]"p2r±'p
1 0
(Eq. 40)
The presence of [M+ ] and [H+] in Eq. 40 requires that the
free metal ion concentration be known for calculation of
@qp and Ki (56).

58
II. Polarography
Derivation of equation for polarographlc analysis of
copper complexes of thlouraclls. Kolthoff and Llngane
(57) have given derivations of the equations to be used
for the calculation of stability constants of complexes
from polarographic data. The equation for the reduction
of cupric-thiouracil complex to the cuprous complex at the
dropping mercury electrode (DME) is
Cu(U)„ + e" ^ CuU + U" (Eq. 41)
The anion of a thiouracil molecule is represented by U”.
The difference between the polarographic half-wave poten¬
tials in the partial reduction of a simple (uncomplexed)
metal ion and in the reduction of a complex ion of the
same metal can be derived if we consider that the reaction,
for the purpose of discussion, occurs in two hypothetical
steps according to Eq. 41 and Eq. 42.
Cu+2 + e~ ^ Cu+ (Eq. 42)
^
The standard potentials for Eqs. 41 and 42 are E° and
E°, respectively. The overall reaction can be written
s
by subtraction of Eq. 42 from 41 since both are one
electron reductions.

59
CuU0 + Cu'
¿ •'C
CuU + U“ + Cu+2
(Eq. 43)
The standard potential of Eq. 43 is the difference between
the standard potentials of Eqs.4l and 42.
E° = E° - E°
c s
(Eq. 44)
The equilibrium constant for Eq. 43 is
[CuU] [U~] [Cu+2]
[Cu+] [CuU,,]
K
ov
(Eq. 45)
The dissociation constants of the oxidized and reduced
complexes are given, respectively, by
[Cu+2] [IT]2
[CuU2]
[Cu+] [U~]
[CuU]
(Eq. 46)
(Eq. 4?)
The ratio of K to K is identical to K . Since the free
o r ov
energy of a reaction is proportional to its potential (58)
we can write the equivalent equilibrium expression relating
potential and equilibrium constant.
RT
nF
ln K„
ov
RT
nF
In
(Eq. 48)

60
The symbols in Eq. 48 have the following meanings: n,
electron change in the reaction; F, the faraday; R, the
molar gas constant; T, the absolute temperature. The
Nernst equations (58) for Eqs-. 42 and 43, respectively,
are
0+]°*
edme = Es - -fir ln
[Cu+
2]°¿r
(Eq. 49)
F - F° RT
DME c " nF
ln
[CuU]0^
[CuU,]0^'
- ' ~r~ ln Cu_] ÍT±
(Eq. 50)
The values of P and Q represent the stoichiometric number
of ligands bound to the oxidized and reduced metal,
respectively, and is unity for the specific case of Eq.
43. The potential of the dropping mercury electrode is
represented by E^g. The concentrations of oxidized and
reduced species appearing in Eqs* 49 and 50 are those at
the electrode surface. The activity coefficients for the
oxidized and reduced states of the simple and complex ions
are represented by Q, and C K- respectively. Be¬
cause of the low concentration of reactive species at the
electrode surface the ratio of X'Q and Yr/ XQ may be
taken as unity and the activity of the thiouracil anion
is assumed to be equal to its concentration (59)»

61
The relation of the concentration of the oxidized and
reduced species in the bulk solution may be derived in the
following manner. Since the initial concentration of re¬
duced species is zero at the electrode surface we can write
[Cu+]° = i/kr
(Eq. 51)
which relates the electrode surface concentration to the
current at any point on the curve. The parameters in Eq.
51 have the following significance; i, the current in
microamps; [Cu+]°, the concentration of cuprous ion at
the electrode surface; kr, the proportionality constant
between current and concentration. The proportionality
factor, kr, is directly proportional to the square root
of the diffusion constant and comes from the Ilkovic
equation (60).
(Eq. 52)
The Ilkovic equation, which has been experimentally proven
(61), assumes that the current at any point on the current
-potential curve is directly proportional to the difference
between the concentration of the active species at the
electrode surface and the bulk solution (62). The para¬
meters in the Ilkovic equation have the following signifi¬
cance: i, current in microamperes at any point on the

62
curve; n, number of electrons Involved in the electrode
reaction; C, concentration of active species in millimoles
per liter; m, weight of mercury flowing from DME in mg.
sec. ; D, diffusion coefficient of active species; t,
drop time in seconds; 706, a constant equal to 4Y7tt/3 F
where F is the faraday (96,500 coulombs). The proportion¬
ality between current and the concentration gradient from
bulk solution to electrode surface is expressed by
i = kQ([Cu+2] - [Cu+2]0) (Eq. 53)
where kQ is the proportionality constant from the Ilkovic
equation and [Cu+2] and [Cu+2]^ are the bulk and electrode
surface concentrations of cupric ion, respectively. The
current will reach a maximum when the active species is
reduced as fast as it can diffuse to the electrode surface
causing the concentration at the surface to fall to zero.
This condition is expressed in Eq. 54 by setting [|Cu+2]^
in Eq. 53 equal to zero.
iD = kQ[cu+2] (Eq. 54)
Substitution of Eqs.51t 53 and 54 into Eq. 49 gives
T? _ T7° RT ko RT i
edme ” es ñF~ ln 1c W~ ln ~T'- I (EC1* 55)
— r — D

63
When i = 1^/2, Eq. 55 becomes
¿ — r
Substitution of Eq. 56 into Eq. 55 gives
EDME = ñP~ ln i -i ^Eq* 57)
The equations relating the diffusion current (i^) to the
bulk concentrations of oxidized and reduced complex are
Cd’6 ' k0C°uU2]
(Eq. 58)
- *;ccuu]
(Eq. 59)
Cathodic and anodic currents are designated by (in) and
u c
(in) , respectively, and "primes" are used for parameters
derived from complexes. The significance of the other
terms has already been mentioned. The electrode surface
concentration of the reduced comples, [CuU]^, is the sum
of its bulk concentration and any that is formed from re¬
duction of the oxidized complex.
i
[CuU]° = [CuU] +
(Eq. 60)

64
In the particular case of Eq. 43 the bulk concentration of
reduced complex is negligible. The electrode surface con¬
centration of the oxidized complex is the difference be¬
tween its bulk concentration and the amount lost by reduc¬
tion at the electrode.
[CuU ]° = [CuU2] - —
k
(Eq. 61)
o
Equations 60 and 6l are obtained by a rearrangement of
equations having the same form as Eq. 53 but written in
terms of the reduced and oxidized complex, respectively.
Substitution of Eqs. 58, 59. 60 and 6l into Eq. 50 and
collection of terms gives
k
k
r
o
(Eq. 62)
If [CuU] = 0 then -(i^)^ is also zero from Eq. 59 and Eq
62 becomes
k
In [U-]
o
k
r
(Eq. 63)

65
When i = 1^/2, then E^g = (E|)c anc^ Eq. 63 is
(E ) = F° -M- ln k°
^Ef'c Ec “ nF 1 k
f- - ^ -f - 1» CiT]
(Eq. 64)
Subtraction of Eq. 56 from Eq. 64 yields
k’k
o r
k k'
o r
(Eq. 65)
which relates the change in the half-wave potential to the
change ln the standard potential of a metal ion when it is
complexed.
The diffusion coefficient of a free cuprous ion is not
readily measurable because of its ease of disproportiona¬
tion to cupric ion and copper metal. Since the ionic radii
of cuprous and sodium ion are O.96 and 0.95 angstroms,
respectively, we can make the approximation that the charge
densities of the two ions are very similar. Because of the
similarity in charge densities the probable size of the
hydration spheres and hence the probable diffusion coef¬
ficients of the two ions would be expected to be the same,
i.e., 1.35 x 10 J cm. sec. as given for sodium ion (63).

66
Since cuprous ion is heavier than sodium ion however, the
value of may be overestimated. On the probably unwar¬
ranted assumption of the validity of the application of
the Stokes-Einstein Law to ionic species, it may be pre¬
dicted that the ratio of the diffusion coefficients of
cuprous to sodium ion is equal to the inverse ratio of
the cube roots of their atomic weights, i.e., (22.9) ' "V
(63*5 (64). This calculation gives a value of 0.96
x 10“5 for the diffusion coefficient of cuprous ion. Since
kr and kQ are proportionality constants which relate cur¬
rent and concentration, consideration of the Ilkovic equa¬
tion (Eq. 52), shows that ratios of the square roots of
the diffusion coefficients may be substituted for ratios
of the proportionality constants. The ratio of the square
roots of the diffusion coefficients of cuprous and cupric
ions (kr/ko), using the value of 1.35 x 10“^ for cuprous
ion, is 1.3 when the diffusion coefficient of cupric ion
is 0.72 x 10“5 cm.2 sec.-'*" (63). The ratio of the square
roots of the diffusion coefficients of the oxidized and
/ * / 1 V
reduced complexes, (ko/kr), as estimated from the Stokes
-Einstein Law and their molecular weights (64) is approx¬
imately 0.92. Therefore, the calculated value of
RT/nF(ln kokr/krkQ) is 4.6 mV which is approximately
experimental error. The true value will be less than
4.6 mV due to the overestimation of kr as 1.35 x 10“
since the hypothetical calculation given above indicates

67
that may be as low as O.96 x 10“-*. Therefore, the last
term on the right hand side of Eq. 65 may be considered
negligible and ignored. Substitution of Eq. 48 and Eq.
7 into Eq. 65 after dropping the last term on the right
gives
In
Ko
[H+]ár+v
Kr
<)
. m. ln [HD]
~Y~
(Eq. 66 )
Eq. 66 is linear with a slope of (P-Q) RT/nF and an inter¬
cept from which the ratio of the stability constants can
be evaluated.
The value of (E^)s for the reduction of cupric to
cuprous ion is not experimentally determinable because
cuprous ion is reduced easier than cupric ion and a two
electron change always occurs. The half-wave potential
at any temperature can be calculated from Eq. 56 if the
values of E° and k /k are known at the temperature in
sor
question. The temperature effect on diffusion coefficients
is only 2% per degree centigrade (65), is negligible and
kQ/kr may be considered constant for small temperature
changes. The standard potential for the reduction of
cupric to cuprous ion, as a function of temperature, can
be calculated from the standard free energy, entropy and
enthalpy at 25° using

68
| H2 + Cu+2 > H+ + Cu+ (Eq. 67)
AP° = AH° - TAS° = -nFE° (Eq. 68)
assuming that the enthalpy and entropy are constant over
the temperature range in question. Sample calculations
and the required thermodynamic values are available in
the literature (66). The potential of the saturated calo¬
mel electrode as a function of temperature is also avail¬
able in the literature (67*68). The respective calculated
values of E° and (Ei)_ (Eq. 56) for the reduction of cupric
2 s
to cuprous ion at various temperatures are: 21°, 0.1512,
-O.O878; 25°, 0.1514, -0.0837; 35°. 0.1521, -0.0806; 40°,
0.1526, -0.0755; 45°, 0.1530, -0.0744. The value of (E¿)g
at 25° agrees very well with the value calculated in the
literature (69), i.e., -0.079V., even though the literature
value was not corrected for the difference in diffusion
coefficients of cupric and cuprous ions. The literature
value of the half-wave potential for the reduction of
cuprous to free copper is +0.143 V versus the saturated
calomel electrode at 18.0° (70) so that, at potentials
more negative than +0.143 V, a two electron reduction
would be expected if free cuprous ion is produced by the
reduction of CuU

69
Derivation of equation for E-w^-E^ /t] values. Sub¬
stitution of 1 = 0.75 iD Into Eq. 57 gives
E.
3/4
- El/2 - 1o« 3
(Eq. 69)
where E^/^ is the potential of the DME when 1 = 0.75 i^»
When 1 = 0.25 in then Eq. 57 gives
= E1 /o - ——- log 0*333
Jl/4 ^1/2
(Eq. 70)
Subtraction of Eq. 69 from Eq. 71 gives the following
values of E^/^ - Ein volts at the indicated tempera¬
tures: 21°, -0.0556/n; 25°, -0.0564/n; 30°, -0.0573/n;
35°, -0.0583/n; 40°, -0.0592/n; 45°, -0.062/n. Experi¬
mental values of E^/^ " El/4 can be comPare calculated values as a means of establishing the reversi¬
bility and electron change of the electrode process (44-47).
III. Solubility Analysis
Derivation of equations for solubility analysis of
thiouraclls and thlouracll complexes. The equation relat¬
ing the dissociation constant of a ligand to its intrinsic
solubility is
Ka = CH+^Ab “ AbH+)/AbH+
(Eq. 71)

70
where Ab is the absorbance in the buffer region and Abg+
is the absorbance of an acid solution saturated with
ligand. The negative logarithm of Eq. ?1 gives
pK^ = pH - log (Ab - AbH+)/AbH+ (Eq. 72)
The apparent total solubility of thiouracils as a
function of pH in the absence
S = [HU] + [IT] = SHU + [IT] (Eq. 73)
and presence of 1:1 and 2:1 complex-forming metal ions is
s' = SHU + [IT] + LMU+] + 2[MU2] (Eq. 74)
The respective total solubility in the absence and presence
t
of metal ion is represented by S and S while the intrinsic
solubility of HU is represented by SH^. The experimental
conditions for solubility analysis, mainly constant pH and
temperature, are given in part III of the experimental sec¬
tion. Equation 74 is the same as the mass balance equation
(Eq. 11) for ligand in the case of 2:1 complexes except
that the solution is now saturated with the undissociated
t
ligand. Similarly, the equations for S for mixed ligand
and polynuclear complex formation can be shown to be simi¬
lar to their ligand mass balance equations (Eqs. 22, 33)

71
where the solutions are saturated with the undissociated
ligand. Subtraction of Eq. 73 from Eq. 74 gives, for
2:1 complexes
S* - S = [MU+] + 2[MU2] (Eq. 75)
and subtraction of the appropriate expressions for the
cases of mixed ligand and polynuclear complexes gives,
respectively,
s' - S = [MU+J + [MUOH] (Eq. ?6)
s' - S = LMU+] + [MU] + 2[M2U2] (Eq. 77)
The definition of the degree of formation, ñ, is given by
Eqs. 14, 25 and 36 for 2:1, mixed ligand and polynuclear
complexes, respectively. Substitution of Eqs. 75. 76 and
77 into Eqs. 14, 25 and 36, respectively, gives the same
expression for ñ in terms of solubility.
n =
(Eq. 78)
The expressions relating ñ to the stability constants for
2:1, mixed ligand and polynuclear complexes are given by
Eqs. 17, 28 and 40, respectively.

RESULTS
I. Potentlometrlc Titrations
Titration of thiouracll-cuprlc ion mixtures. Solu¬
tions, 0.002 or 0.0016 M in thiouracil(2-thiouracil,
6-n-propyl-2-thiouracil, 6-methyl-2-thiouracil, 5-raethyl-
-2-thiouracil and 5»6-dimethyl-2-thioura.cil) and 0.002 to
0.0002 M in cupric nitrate, shotted an immediate drop in pH
(ca. 5«5 to ca. 3.0) and the slow formation of precipitates.
Potentlometrlc titrations of these suspensions with stan¬
dard NaOH immediately produced more precipitate. The re¬
sultant titration curves (curves A, E and C, Figs. 1 and
15). using 2-thiouracil as an example, showed more alkali
consumption to pH 6 than the titration of the same amount
of the ligand alone (curve G, Figs. 1 and 15)* This large
alkali consumption was measured up to inflection 1 of curve
A (Fig. 15) and inflection 2 of curve A (Fig. 1), a pH
region where the hydrolysis of free cupric ion does not
occur (curve H, Figs. 1 and 15)»
The titration curves indicated the presence of at
least three species. The two most acidic species (pK* 3«1
&L
and 4.4) are most easily seen in curves A - E of Fig. 1.
72

73
A comparison of curves A - E of Fig. 15 (titrated immedi¬
ately after preparation) with the same curves of Fig. 1
(titrated the day after preparation) show that the species
V f
with pK 3*1 increased in titer at the expense of the pK
a a
4.4 titer. The net titer to the pH 6 inflection remains
unchanged. The titer "between inflection 1 and 2, curve A,
9
Fig. 15 has the same pK as curve H, the titration of cu-
a
pric ion alone and is assigned thereto.
The titer of excess, uncomplexed 2-thioura.cil extends
from inflection 1 to 2, curves D, E and F, Fig. 15 and from
inflection 2 to 3. curves D, E and F, Fig. 1 as can be
ascertained from comparison with the pK 7.49 from the
a
titration of ligand alone (curve G, Figs. 1 and 15)*
If all the cupric ion added was complexed either with
thiouracil or hydroxyl ion then the total alkali consumed
in titrating the complexes (i.e., to pH ca. 6, curves A - F,
Fig. 15, inflection 1 and Fig. 1, inflection 2) would be
expected to be twice the molar amount of cupric ion used
because cupric ion is divalent and consumes two equivalents
of alkali per mole of metal ion. This was found to be the
case. The total titer of alkali added to reach inflection
3 of curve A, Figs. 1 and 15, was only slightly less than
twice the titer of alkali consumed by the ligand alone
(curve G), indicating that the cupric-ligand complexes
still exist even at a relatively high pH and that the
higher hydroxide concentration does not cause disruption
of the complex and formation of Cu(0H)2.

Pages
are
misnumbered
following
this
insert

75
Continued titration past the point of precipitation of
the equi-molar ligand-metal solutions gave curve A, Fig. 16,
with an inflection at pH 8.3 corresponding to a total titer
of alkali equal to twice the alkali that would have been
consumed by the ligand alone. The amount of free, uncom-
plexed thiouracil at less than equimolar amounts of lead to
ligand, as estimated from the titer of alkali between in¬
flections 1 and 2 of curves B - G of Fig. 16 (compare with
curve H, Fig. 16), also showed that the precipitating com¬
plex had a 1:1 stoichiometry. The titer of uncomplexed
thiouracil corresponded to that in excess of 1:1 stoichiom¬
etry. Infrared analysis of the precipitate, isolated at
pH 7.5 gave a curve identical to the infrared curve of
bis(2-thiouracil-cadmium(II)), .
The titration data, up to the point of precipitation,
were analyzed according to Eq. 18. Typical plots of log
1-ñ/ñ versus p[u~] (Eq. 18) are given by Figs. 1?, 18, 19
and 20. The slopes for all plots of Eq. 18 are given in
Table I. In those cases where the metal and ligand con¬
centrations were equal, the slopes of plots of Eq. 18 were
unity. Therefore, the calculation of log for soluble
1:1, MU+, complexes was valid in these cases (Table I).
When the total metal ion concentration, [M+2]q, was less
than the total ligand concentration, [HU]q, the slope of
Eq. 18 was always between one and two which would be ex¬
pected if appreciable amounts of MU2 were formed.

76
The plots of Eq. 18 are not linear over their entire
length (Figs. 1?. 18, 19 and 20). Negative deviations
occur at high pH values due to precipitation. Positive
deviations at low pH values are attributed to error in
reading the very small amounts of initial alkali titer
on the recorder chart. These minor deviations were ignored
when the best straight lines were drawn by sight.
The data from the titration curves up to the point of
precipitation were also plotted according to Eq. 17.
((ñ/l-ñ)[u“] versus te) [U-], Fig. 21) and values of log
and log K2 determined from the slopes and intercepts
(Table II). The values of log ranged from approximately
4.7 to 5.0 for the various thiouracils while log K2 was
approximately 3»^ at 25° • The magnitude of log and log
K2 were not dependent on the total metal ion concentration
used.
Titration of thiouracil-cadmium ion mixtures. Poten-
tiometric titration of all thiouracil-cadmium mixtures,
except 6-n-propyl-2-thiouracil, gave very similar curves
as those found for the lead studies (Fig. 16). Precipita¬
tion occurred during the titration at about pH 6.5, limit¬
ing the calculation of stability constants to the region
of homogeneous solution. The alkali consumed to inflection
1 for the equimolar mixture, represented by curve A, Fig.
16, was twice the alkali consumed by the ligand alone

77
indicating the precipitation of mixed ligand complexes,
M^Ug. If the alkali had destroyed the precipitated com¬
plexes and formed M(0H)2* the total alkali titer would
have been three times the alkali consumed by the ligand
alone. The titer of excess thiouracil found at cadmium
concentrations less than equimolar, as calculated from
the alkali consumed between inflections 1 (near pH 7) and
2 (near pH 9) also indicated a complex being precipitated
with a 1:1 stoichiometry. The infrared curve of the
2-thiouracil precipitate, isolated at pH 10.5, was identi¬
cal to the infrared curve of bis(2-thiouracil-cadmium(II)),
M2U2, which was synthesized as outlined in the experimental
section.
The titration curves of cadmium-6-n-propyl-2-thiouracil
solutions at 25° and 35° (Fig* 22) differed from the curves
obtained with the other ligands and from the ligand-lead
mixtures. The titration curves showed precipitation near
pH 6 as before but the first inflection, near pH 8.0, for
all concentrations of cadmium ion equal or greater than
0.001 M (curve A and B, Fig. 22) occurred at an alkali
consumption equal to the titer of alkali consumed by the
ligand alone (the concentration of ligand was constant at
0.002 M). Only when the total cadmium concentration was
less than 0.001 M (curve C, Fig. 22), half of the ligand
concentration, was any uncomplexed ligand indicated. This
behavior requires a stoichiometry of MUg. However, at a
temperature of 45° cadmium-6-n-propyl-2-thiouracil solutions
gave curves as in Pig. 16, indicating a 1:1 stoichiometry.

78
That part of the titration data resulting from homo¬
geneous solution was analyzed by Eqs. 17 and 18 as explained
in the previous section. The slopes of Eq. 18 (Table I)
for less than equimolar concentrations of cadmium ion were
between 1 and 2 as was the case for lead and indicated
formation of MU2 complexes. When the slope of Eq. 18 was
unity the calculation of log was valid (Table I).
Typical plots of Eq. 18 (Pig. 17. 18, 19 and 20) are the
same as for lead ion. Typical plots of Eq. 17 are shown
in Fig. 21 and the values of log and log Kg are given
in Table II. The log and log K2 values were not depen¬
dent on the total metal ion concentrations used.
Titration of nickel-thiouracil and zinc-thiouracil
mixtures. Titration of 2-thiouracil, 6-n-propyl-2-thi-
ouracil, 6-methyl-2-thiouracil, 5*-methyl-2-thiouracil and
5,6-dimethyl-2-thiouracil in the presence of nickel gave
curves that showed an initial pH drop from 6.0 to 5*5 and
precipitation during the titration near pH 8 (Fig. 23
curves A - D). Analysis of the part of the titration
curves corresponding to homogeneous solution by Eq. 18
gave slopes between 1 and 2 (Table I) indicating a stoi¬
chiometry of MUg. Analysis by Eq. 17 gave values of log
and log Kg (Table II) from the slope and intercept which
were about two units smaller than the log and log Kg
values for the cadmium and lead complexes of the thioura-
cils

79
Titration of 6-n-propyl-2-thiouracil in the presence
of zinc gave an initial pH drop from 6.5 to 6.0 and pre¬
cipitation occurred near pH 7.3» The titration curves for
the zinc complexes were very similar to those found for
nickel (Fig. 23)* The data were plotted according to Eq.
17 and log and log Kg values are given in Table II.
The stability constants for both nickel and zinc are
smaller, by more than a factor of ten, than those found
for cadmium and lead.
The total alkali consumed to inflection 1 curve A,
Fig. 23, was greater than twice the titer of alkali con¬
sumed by the ligand alone (curve E) indicating formation
of M(0H)g complexes. The alkali consumed to inflection 2,
curves B - D, Fig. 23. was greater than the alkali consumed
by the ligand alone again indicating formation of M(0H)g
complexes and probable disruption of the metal-thiouracil
complexes already formed.
Titration of 2-thlouracil and 6-n-propyl-2-thiouracil
in the presence of other metal ions. Potentiometric ti¬
trations of 2-thlouracil and 6-n-propyl-2-thiouracil in
the presence of ferric, ferrous, manganese, calcium and
cobaltous ions gave no indication of any complex formation.
The titration curves of the mixtures could be assigned to
uncomplexed ligand in the case of calcium and manganese and
to hydrolysis of the metal ion in the case of ferric, fer¬
rous and cobaltous ions.

80
Titrations of sterically blocked thlouracll in the
presence of Cu(Il), Cd(II) and Pb(Il). Solutions of
2-ethylmercapto-4-hydroxypyrimidine (2EM4HP) and N,N -di-
ethyl-6-methyl-2-thiouracil in the presence of cadmium,
lead and cupric ions were titrated with standard alkali.
The titration curve of 2EM4HP in the presence of cadmium
ion was the same as the titration of the ligand alone.
The titration curves of 2EM4HP in the presence of cupric
and lead ion could be assigned to a simple summation of
the metal hydrolysis and ligand titration curves indicating
no apparent complexation. The titration of solutions of
N,N -diethyl-6-methyl-2-thiouracil alone showed no titrat-
able group and the titrations of metal ion mixtures gave
curves showing only metal hydrolysis. In the case of
2EM4HP, the pK of the 4-hydroxy group was 7.01 at 25°.
3
The fact that no complexation occurred at the 4-hydroxy
position even when it could form an anion more easily than
2-thiouracil itself (pK 7.49) argues for complexation at
3
the sulfur position in the sterically unblocked compounds.
t
Titration of aqueous solutions of ligands. The pK
3
of each ligand at 25*0° was determined by averaging at
least three determinations. This is important since small
f
changes in pK have a relatively large effect on the calcu-
3
lations of the stability constants. The pit' at 35.0° and
3
45.0° were usually single determinations made at the same

81
time the complexation titrations were recorded to avoid
any possible error due to possible temperature effects on
I
the electrode or any other part of the system. The pK
St
values were estimated from the pH of half-neutralization
when the ligands were titrated with standard alkali. The
I
solutions, temperatures and pK values are given in Table X*
a
Effect of temperature on the stability constants of
cadmium and lead complexes of thiouracils. The values of
log (Table II) at 25°. 35° and 45° were used to calcu¬
late the change in free energy as a function of temperature
by
/\F = -RT In K1
(Eq. 79)
Since
AF = AH - TAS (Eq. 80)
the values of AF, the change in free energy at temperature
"T", may be plotted versus the absolute temperature where
AH, the intercept, is the enthalpy change and the slope,
AS is the entropy change of the complexation reaction.
Relatively large error in AF results from small error in
log since the temperature range was purposely restricted
to maintain AH constant (71). Thus, only the general

82
trends in AH and AS with structural variations in the
thiouracils may have significance. The AS values (Table
XI) for the first complexation reaction tended to be
positive (apparent trend for negative slopes in Fig. 24)
except the value for lead-6-n-propyl-2-thiouracil which
was zero or slightly negative. The enthalpy change, AH
(Table XI), for the complexation of lead with the ligands
at 37«5° was negative and varied from -3*5 Kcal. for
6-methyl-2-thiouracil to -7.9 Kcal. for 6-n-propyl-2-thi-
ouracil with an estimated error of + 0.1 Kcal. The
enthalpy change at 37»5° for the cadmium complex of
5-carboethoxy-2-thiouracil was slightly positive. The
AH values for all other cadmium complexes were negative
except for that of 6-methyl-2-thiouracil which was approxi¬
mately zero.
The values of AF for the lead complexes, taken as
a group, were about one kilocalorie lower than for the
cadmium complexes.
Effect of temperature on the acid dissociation con¬
stants of the ligands. Plots of Eq. 80, where the changes
in free energy calculated from the acid dissociation con¬
stants, K of the ligands (Eq. 79) were plotted against
the absolute temperature (Eq. 80), seemed to give negative
entropy changes (tendency for positive slopes in Fig. 25)
except in the case of 6-n-propyl-2-thiouracil. No clear

83
-cut correlation of /\S with ligand is possible due to the
small temperature range required to maintain AH constant.
The enthalpy changes for the acid ionization were all
positive and varied from 3.5 to 13.2 Kcal. with an
estimated error of + 0.1 Kcal. and are given in Table XI.
II. Ultraviolet Spectral Studies
Jobts plots of 2-thiouracil and 6-n-propyl-2-thiouracll
in the presence of cupric ion. Aqueous solutions of a con¬
stant concentration of 2-thiouracll (5*00 x 10“^M) and equal
or lesser concentrations of cupric ion were run on the Cary
recording spectrophotometer. The spectra (Fig. 26) showed
a decrease in the absorbance of the 272 mp maximum and a
small rise in absorbance near 3^5 mji where 2-thiouracil
has no absorbance. Isosbestic points occurred at 306 and
25^ mja which may indicate a single transformation. A plot
of absorbance at 272 and 3^5 versus the ratio of the
concentration of cupric ion to 2-thiouracil, at constant
concentration of 2-thiouracil (5«00 x 10“^), showed a
deviation from linearity at 0.5 indicating a stoichiometry
(Cu:U = 1.2) of CuU2 for the complex (Fig. 9) on the assump¬
tion of strong complexation. Job’s plots (72) of aqueous
cupric-2-thiouracil solutions showed rounded peaks with
maxima at a mole fraction of ligand of 0.66 which also
indicates a stoichiometry (Cu:U = 1.2) of CuUg. The com¬
position of the solutions used in Fig. 8 are given in

Table XII. The composition of the solutions used for Fig.
9 are outlined in the experimental section. A Job's plot
of 6-n-propyl-2-thiouracil and cupric ion in 0.05 M HCIO^
showed no discernible change in the absorbance that could
be assigned to complex formation.
Ultraviolet spectra of dilute acid solutions of
cuprlc-2-thiouracll mixtures. Since the cupric complex
of 2-thiouracil is so strong the equilibrium constant is
best measured in dilute acid so as to maintain significant
concentrations of complexed and free ligand. The ultra¬
violet spectra of acid solutions of the 2-thiouracil-cupri
ion mixtures listed in section III of the experimental were
run. Prepared 0.1 M HCIO^ solutions containing 4.5 x 10“-^
M cupric ion were diluted 1:50 and read on the Cary spec¬
trophotometer. The absorbances at 270 m)i were similar to
that for 2-thiouracil alone. The total change in absorb¬
ance was about 3 percent which is insufficient for good
analysis. The small change in absorbance, due to the
dilute solutions necessary for UV spectra, would cause
large errors in the calculated stability constants. Higher
pH values would cause precipitation of the complex. This
result indicates that the molar absorption coefficient of
the complex may be about the same as the value for the
uncomplexed ligand. Since the polarographic data indicated
a strong complex does exist then the small change in

85
absorbance is probably due to the low concentrations
necessary for UV spectra. The molar absorption coeffi¬
cient of the complex may be about the same as the uncom-
plexed ligand.
Furthermore, the data indicate that no sulfinic acid
(73) is formed and the proposed oxidation of 2-thiouracil
by cupric ion in alkaline solution does not occur in acid
solution (73)» The absorbance of the diluted solutions
-3 -4
containing 4.5 x 10 J M and 2.15 x 10 M cupric ion was
1.305 and 1.250, respectively. The absorbance of 2-thi¬
ouracil alone is calculated to be 1.233* The absorbance
that was expected if oxidation occurred was 0.995 which
was based on formation of the disulfide and subsequent
disproportionation to the sulfinic acid and 2-thiouracil.
The calculated absorbance is based on the molar absorption
coefficients of sulfinic acid and 2-thiouracil given in the
experimental section and Eq. 2.
Spectrophotometric titration of 2-thiouracil. Anal-
ysis (Eq. 3) of the change in absorbance as a function of
increasing pH of aqueous solutions of 2-thiouracil gave a
pK near 7*6. Isosbestic points occur at 302, 258 and 219
SI
nya. The spectra and plots of Eq. 3 are shown in Fig. 10
and 11.

86
Spectrophotometric titration of 2-thiouracll-cuprlc
Ion mixtures. The shapes of the spectra of highly dilute
aqueous mixtures of cupric ion and 2-thiouracil (Fig. 12)
as a function of increasing pH were nearly the same as for
the curves obtained during the spectrophotometric titration
in the absence of cupric ion (Fig. 10) except in the region
of 255 to 305 mp where the absorbance was reduced due to
the complex. A rise in absorbance near 3^5 mp was similar
to the rise in absorbance of spectra analyzed in the Job's
plots. The value of pK (7.5^) was only slightly changed
o.
from the value of 7*3 found in the absence of metal ion.
These results indicate that the cupric complex of 2-thi¬
ouracil absorbs ultraviolet light less than the ligand
does in the pH region 6.0 to 10.0 between 255 and 305 mp.
The complex weakly absorbs ultraviolet light above 305 mp.
-4
Concentrations of cupric ion 0.500 x 10 and larger caused
precipitation which frustrated analysis. The useful data
—4
collected at a cupric concentration below 0.500 x 10 were
analyzed using Eq. 3 at 270, 285, 230, 240, 310 and 320 mp.
The spectra and plots of Eq. 3 are given in Figs. 12 and
13.
Degradation of 2-thiouracil disulfide in the absence
and presence of cupric ion. It is known (73) that hydrol¬
ysis of disulfides proceeds by dispxoportionation according
to the following reaction.

87
2 RSSR + 3 H20 > 3 RSH + RSO^H (Eq. 81)
Cupric ion can oxidize 2-thiouracil to the disulfide in
dilute alkali (14). It was conceivable that the reaction
might also occur in dilute acid but would yield the sul-
finic acid and 2-thiouracil according to Eq. 81, since
it has been shown (14) that the disulfide of 2-thiouracil
is very unstable in dilute acid and produces 2-thiouracil
as one of the products. If cupric ion is present when N
moles of 2-thiouracil disulfide disproportionate the fol¬
lowing series of reactions can be expected.
N RSSR + 2N HgO > 3N/2 RSH + N/2 RSOgH
(Eq. 82)
3N/2 RSH + N Cu+2 > N/2 RSH + N/2 RSSR + N Cu+
(Eq. 83)
N/2 RSSR + N H20 > 3N/4 RSH + N/4 RSOgH
(Eq. 84)
The sum of the final concentrations of RSH in Eq. 83 and
84 is
RSH = 5N/4 RSH
(Eq. 85)

88
The sum of the final concentrations of RSOgH is
^ RS02H = 3N/4 RS02H (Eq. 86)
The final absorbance of the solution will be the sum of the
absorbances of RSH and RSOgH. Ey using Eq. 2, the molar
absorption coefficients of 2-thiouracil (C 13,680) and
its sulfinic acid derivative (£ 2^q 3,510) and the original
concentration of 2-thiouracil disulfide (N = 3*99 x 10“-'M),
the final absorbance can be calculated in the absence and
presence of cupric ion. The expected absorbance in the
absence of cupric ion, as a result of the disproportiona¬
tion expected in Eq. 82 is 0.89. The observed absorbance
when equimolar and twice equimolar concentrations of cupric
ion (N and 2N) as of disulfide were mixed, was 0.83. This
was in reasonable agreement with the expected value of
0.89. The small difference may be due to partial formation
of the sulfonic acid derivative of 2-thiouracil (73).
The expected change in absorbance for the postulated
sequence of Eqs. 82 - 86 when N and 2N relative concentra¬
tions of cupric ion were used would be 12 percent and 23
percent, respectively. However, the final absorbance
values for disulfide in acid solution were the same in
the presence and absence of cupric ion (i.e., 0.83). It
can be concluded that oxidation does not occur in acid

89
solution and further substantiates the spectral studies
which indicated that the molar absorption coefficient of
the cupric complex of 2-thiouracil in acid solution must
be nearly the same as for 2-thiouracil.
Aqueous solubility of thlouraclls as a function of
pH. The intrinsic solubility of 2-thiouracil, 6-n-propyl-
-2-thiouracil, 6-methyl-2-thiouracil, 5-methyl-2-thiouracil,
5,6-dimethyl-2-thiouracil, 5-carboethoxy-2-thiouracil and
6-amino-2-thiouracil varied from 7.97 x 10"*^ to 1.79 x 10“^
Ab-Ab„+
M. Plots of log versus pH were prepared and a
AbH+
typical plot is given in Fig. 2?. The calculated intrinsic
solubilities and pK values are given in Table IV for the
a
various ligands.
Aqueous solubility of 2-thlouracil and 6-n-propyl-2-
-thlouracll as a function of pH in the presence of cadmium
ion. Due to the fact that precipitation of complexes must
be avoided in solubility analyses, the pH range was re¬
stricted to values less than 6.5. This restriction
drastically reduced the concentration of 2-thiouracil anion
and suggested calculations which indicated that only a 1.3
percent increase in total solubility of 2-thiouracil could
be predicted from the stability constants obtained from
the potentiometric titrations data. This increase is less
than experimental error and thus no stability constants
could be obtained by this method.

90
Spectra of aqueous thiouracil-metal Ion mixtures.
The spectra of ligand solutions (2-thiouracil, 6-n-propyl-
-2-thiouracil, 6-methyl-2-thiouracil, 5-methyl-2-thiouracil
and 5.6-dimethyl-2-thiouracil) containing ferrous and
ferric ions showed no evidence of complexation. The spec¬
tra of the solution mixtures were the same as the sum of
the spectra of the ligand and metal ion alone. A typical
plot showing the summation of absorbances for ferric ion
is given in Fig. 4. Ferric ion absorbs in the ultraviolet
with a X at 295 mu but ferrous ion does not absorb.
Solutions containing ferrous ion often contained small
amounts of ferric ion due to the ease of oxidation. The
amount of contaminating ferric ion in the ferrous ion
solutions was easily corrected for by running blank versus
metal ion alone and subtracting the ferric ion absorbance
from the absorbance of the mixture. These ligand solutions
with a seven-fold excess of cadmium and lead showed about
a 10 percent loss of absorbance in the absorption maxima
occurring near 272 mp (Fig. 3)* This change is not enough
to allow accurate calculation of the stability constants
since the changes in absorbance readings as affected by
large changes in metal ion concentration would be small.
Also, since the molar absorption coefficients of the com¬
plexes must be known in order to calculate the stability
constants (74) and the molar absorption coefficient of MU+

91
is not easily measured, spectral measurements using cadmium
and lead were abandoned in favor of potentiometry.
Aqueous ligand solutions with cupric ion but no added
acid showed a large reduction in the absorbance of the 272
mjj. maximum and a rise in absorbance correlated with the
large change in pH found by the titration studies.
Isosbestic points occurred near 306 and 255 in all
cases. Due to the difficulty in measuring molar absorption
coefficients as previously cited and the occurrence of
precipitation of the cupric complex in the cells no attempt
was made to determine stability constants.
Spectra of sterically blocked ligands in the presence
of metal ions. The aqueous spectra of 2-ethylmercapto-4-
-hydroxypyrimidine and N,N -diethyl-6-methyl-2-thiouracil
in the presence of cupric, lead, cadmium, ferrous and
ferric ion showed no change from the spectra obtained in
the absence of metal ions. This indicates that no com-
plexation occurred and is further evidence that the com-
plexation occurs at the sulfur atom. The spectra of
2-ethylmercapto-4-hydroxypyrimidine and N,N*-diethyl-6-
-methyl-2-thiouracil in the presence of cupric ion are
given in Pigs. 5 and 6.

92
III» Polarography
Determination of ratios of dissociation constants of
oxidized and reduced cuprlc-thlouracil complexes. Plots
of (Ei)c - (Ea)s () versus the negative logarithm of
the molar thiouracil concentration (p[HU]) for 2-thiouracil
6-n-propyl-2-thiouracil and 5*6-dimethyl-2-thiouracil gave
fair straight lines with slopes consistent with the theo¬
retical requirements of Eq. 66. Typical plots of Eq. 66
are given in Fig. 28 and the negative logarithm of the
ratios of the dissociation constants for the oxidized and
reduced complexes (-log KQ/Kr) are given in Table XIII.
The theoretical slopes, equal to 2.303 RT/F, as a function
of temperature are: 21°, -O.O5836; 25°. -0.05916; 30°,
-0.06015; 35°. -0.06114; 40°, -0.06213; 45°, -0.06312.
The polarographic curves became irreversible, as determined
by the E^/^ - Evalue (Eq. 6), at a hydrogen ion con¬
centration lower than 0.01 M. Thus, the perchloric acid
concentration was maintained at 0.05 M or higher in all
cases. The values of E^/^ - E.jy^ as determined from the
polarographic curves were within ten mV of the theoretical
value calculated by Eq. 6 for a one electron change. A
typical polarogram showing the background current and
method of determining E^/^. E^y^ and E-jyg is given in Fig.
14. The intercepts of theAE^g plots similar to those of
Fig. 28 were used to calculate the negative logarithm of
*

93
the ratios of the dissociation constants of the oxidized
and reduced complexes (Table XIII) in accordance with the
relationship of Eq. 66.
The value of the theoretical slope which best fits
the data required that the value of (P - q)/n from Eq. 66
be unity. Since the value of n, the electron change, was
unity then the value of P - Q must also be unity. Polaro-
grams of the ligand alone, as shown in typical Fig. 14,
indicated that reduction of the ligand does not occur.
At the higher temperatures (Table XIII) the line of
best fit for the data (-kE^^ versus - log [HU]) sometimes
had a zero slope. This indicated that the value of P - Q
from Eq. 66 under these conditions was equal to zero and
the values of E^g were independent of the ligand concen-
trat ion.
In general, the negative logarithm of the dissociation
constant ratio decreased with increasing temperature (Table
XIII). This means that with increasing temperature there
is an increase in the dissociation of the oxidized complex
greater than the increase in the dissociation of the
reduced complex.
The derivation of Eq. 66 requires that the concentra¬
tion of ligand be large compared with the concentration
of metal ion. The sensitivity of the polarograph required
-4
a cupric ion concentration of 2.00 x 10 M. This put a

94
_3
lower limit of about 2.00 x 10 M on the ligand concentra¬
tion. The upper limit of the ligand concentration is the
intrinsic solubility of the ligand which is of the order
of 5 x 10"3 M (Table IV). This severely restricted the
range of ligand concentrations available and made accurate
analysis of the intercepts and hence, the ratio of the
dissociation constants more difficult.
The solutions of 6-methyl-2-thiouracil and 5-methyl-
-2-thiouracil gave results similar to those found for 2TU,
PTU and 5.6DMTU except that on standing in the polarographic
cell the waves became much steeper and irreversible. This
behavior is unexplained but since it occurs only in the
polarographic cell it may be due to some interaction with
the mercury drop such as adsorption. The /\Ei values for
6MTU and 5MTU were taken on the first curves obtained from
each solution to avoid the irreversible behavior as much¬
as possible.
The polarographic curves for 5-carboethoxy-2-thiouracil
were very irreversible as indicated by values of
in the range of -0.08 to -0.09. The half-wave potentials
were independent of the ligand concentration, indicating
a value of (P - Q)/n of zero.
Some of the values of the negative logarithms of
KQ/Kr, as noted in the footnotes in Table XIII, are not
accurate due to irreversible polarographic waves and are

95
given only for purposes of attempted correlation with the
antithyroid activity. The other values in Table XIII are
valid.
Polarography of solutions containing cuprous ion in
the presence of 2-thiouracil and 2-thiouracil disulfide.
The polarograms of solution A, page 34, containing solid
CuCl, 0.2648 M HCIO^ and 4.375 x 10"^ M 2-thiouracil gave
a reversible wave (E^/^ - E-y^ = -0.063) and the Ey2 value
of -O.I53 mV was the same as found for a solution contain¬
ing cupric ion and 2-thiouracil under the same conditions.
Only part of the white CuCl dissolved when heated slightly
and the rest slowly turned to orange crystals. The anodic
polarogram was also recorded to look for the wave due to
oxidation of free or complexed cuprous to cupric ion but
none was found.
The polarogram of solution B, containing 2-thiouracil
disulfide in acid solution produced no wave and was the
same as for 2-thiouracll under the same conditions (Fig.
14).
The polarogram of solution C, containing solid CuCl,
3.148 mg. of 2-thiouracil disulfide (corresponding to
0.995 x 10"4 M) and 0.2648 M in perchloric acid gave a
wave with an E- Ey^ value of -O.O78 mV and an Ey2
of -O.O9O V which is very close to the E.y2 value for the
reduction of free cupric to cuprous ion. The CuCl did not
dissolve and what remained turned orange.

96
Polarography of presumed cuprous complex of 2-thloura-
cil. Solutions of the presumed solid cuprous complex of
2-thiouracil and dissolved 2-thiouracil, as given in the
experimental section, were polarographed. When the pre¬
sumed cuprous-2-thiouracil complex xvas run in the absence
of added 2-thiouracil the E-yg value (O.O78 V vs. S.C.E.)
was within 9 mV of the Eyg value (-.0S7 V vs. S.C.E.)
calculated for the reduction of cupric to cuprous ions.
This indicates that the dissolved copper was in the cupric
oxidation state.
Polarograms of the cuprous-2-thiouracil complex in
the presence of added 2-thiouracil gave Ey¿^ - Evalues
of -O.O58 V and an Eyg value the same as that found for
cupric ion in the presence of the same 2-thiouracil con¬
centration.
Addition of concentrated ammonia to fresh, dilute
acid solutions of the cuprous complex of 2-thiouracil gave
an instantaneous blue color typical of cupric ion.
IV. Animal Experiments
Determination of antithyroid activity of 5»6-dimethyl-
-2-thlouracil. The increase in thyroid weight (Table VII)
over the controls was significant but there was no signifi¬
cant differences in the thyroid weights at the three dose
levels used (Table VII). Each of the doses gave the maxi¬
mal response within the experimental error. Thus, it is

97
impossible to compare the antithyroid activity of 5.6-di-
methyl-2-thiouracil (56DMTU) with the antithyroid activity
of 2-thiouracil (2TU) below or at one-half the maximal
response. The dose-response curve of 2-thiouracil used
for comparison was constructed from data in the literature
(10).
The procedure previously used (10) for calculation of
the relative antithyroid activity of antithyroid drugs was
to calculate the ratio of that dose of test drug to the
dose of 2-thiouracil necessary to produce a preselected,
standard increase in thyroid weight. However, in the case
of 5.6-dimeth.yl-2-thioura.cil, 6-methyl-2-thiouracil and
5-methyl-2-thiouracil only single dose determinations were
performed (10). Under the restrictions of single dose
determinations it is not possible to know what part of the
sigmoid, dose-response curve the dose corresponds to. If
the dose happens to produce the maximal response then the
calculated relative activity will be too low by an undeter¬
mined amount. Comparison of the dose used in the literature
(10) and the elicited response with the doses in the present
case show that indeed the response found in the literature
for 5»6-dimethyl-2-thiouracil was maximal. The relative
antithyroid activity of 56DMTU, calculated by the litera¬
ture procedure and using the lowest dose of 56DMTU given
(0.025%) was 2.3 that of 2-thiouracil. The true relative
activity will be larger than 2.3 since a lower dose may
also give the maximal response.

96
In the present case a better comparison of relative
antithyroid activity might be to use relative intrinsic
(maximal) biological activities of the various thiouracils.
These values can be obtained by calculation of the ratio
of the maximal increase in thyroid weight over controls
for the test drug to the masimal increase in thyroid weight
over controls produced by the standard drug, 2-thiouracil.
Calculation of this value is 0.37 by
Relative
Maximal Test Drug Activity - Control_ Intrinsic (Maximal)
Maximal 2TU Activity - Control Biological
Activity
(Eq. 87)
The experimental data and results are given in Table VII.
Determination of antithyroid activity of 6-methyl-
t
-N,N -dlethyl-2-thlouracll.-—The relative antithyroid
activity of 6-methyl-N,N,-diethyl-2-thiouracil was found
to be within experimental error when compared with 2-thio¬
uracil. During the two-week experiment two of the five
treated rats died. One apparently died of starvation due
to refusal to eat the treated diet and the other died
after an injury to its foot. The results of the experiment
are given in Table VIII.

99
Determination of antithyroid activity of 6-n-propyl-
-2-thiouracil and 5-ðyl-2-thiouracil. The doses of
6-n-propyl-2-thiouracil (PTU) and 5-®ethyl-2-thiouracil
(5MTU) used produced maximal thyroid response within
experimental error (Table IX). The relative antithyroid
activity of 5MTU, calculated at the lowest dose used
(0.025$) according to the literature method (10) was 2.0.
This is appreciably larger than the value of 0.7 from the
literature (10). The true relative activity will be greater
than 2.0 due to the fact that a maximal response was
elicited at the lowest dose.
The calculated relative activity of 6-n-propyl-2-
-thiouracil (PTU) was 6.5 but since the lowest dose used
produced maximal response then the true relative activity
will be larger. Thus, the literature value (10) of 11.0
for the relative activity is a better estimate.
The relative intrinsic antithyroid activities of 5MTU
and PTU calculated by Eq. 87 are 0.44 and 1.13» respective¬
ly*

DISCUSSION
I. Structure of Precipitated Cupric-Thlouracil Complexes
Potentiometrlc titrations, elemental analysis and
Infrared spectra data. The mixing of cupric nitrate and
thiouracil solutions gave an immediate drop in pH to about
three and the slow formation of a precipitate (Figs. 1,
15)» The large, immediate drop in pH indicated that a
strong complex or complexes were formed. Since homogeneous
equilibrium conditions are necessary for the determination
of stability constants by titration, this method could not
be used for that purpose.
However, the data could be used to obtain evidence
for the stoichiometry of the precipitated complexes. The
fact that more alkali was consumed to inflection 1 of
curves A, B and C of Fig. 1 than the equinormal amount
of ligand taken (curve G, Figs. 1 and 15) required the
precipitated complex to be either a mixed ligand complex
of the form MUOH or that two protons were dissociated from
the ligand when a cupric ion was bound (Scheme IV).
The dissociation of the second proton from the poten¬
tial hydroxyl group at the four position could produce a
zwitterion (XVI). The possible dimerization of the
100

101
H
Scheme IV

102
zwitterion to give a neutral, insoluble complex with stoi¬
chiometry MgUg (XVII, Scheme IV) is plausible. However,
infrared analysis of the synthesized complex denied this
possibility as it indicated that the carbonyl moiety at
the number four position of the pyrimidine ring was present.
The identity of the carbonyl group was established by an
intense absorption band at 1640 cm.” which is in the re¬
gion of the band assigned (75) to the carbonyl absorption
of the ligand. Elemental analysis of the synthesized com¬
plex is not consistent with a stoichiometry of MgUg. There¬
fore, ring formation to give a stoichiometry of M2U2 is
excluded.
The formation of a chelate structure (VII) has been
considered (25) but is ruled out by the presence of the
1640 cm.”^ carbonyl band, the elemental analysis data and
the fact that it is stereochemically impossible. Another
possibility, barely conceivable on solubility grounds, is
the formation of the sodium salt of the anion produced by
the second dissociation (XVI) and precipitated by addition
of another anion (OH”) to the singly complexed cupric ion.
This also would require the loss of the carbonyl function
which loss was not observed. The elemental analysis of
this species would be expected to indicate sodium ion but
did not. Therefore, the formation of MTJOH (XIX) or com¬
pounds with the same empirical formula as the principal
species precipitated during the titration is consistent
with all of the information.

103
The cupric complexes of 2-thiouracil and 6-n-propyl-
-2-thiouracil whose syntheses were outlined in the experi¬
mental section, had elemental analyses, after extensive
drying, which corresponded to empirical formulae of
Cu2(U“)20 (XXI, Scheme IV). The loss of one mole of water
from the cupric complex of 2-thiouracil was observed. The
presence of the 1640 cm.”'*' carbonyl band has already been
mentioned.
If MUOH (XIX) or compounds of the same stoichiometry
(XX), were the only complexes formed, whether precipitated
or not, inflection 1 of curves A, B and C of Fig. 15 and
inflection 2 of curves A, B and C of Fig. 1 would be ex¬
pected to occur at an alkali consumption equal to twice
the amount of ligand taken. This was not the case. The
inflections occurred at an alkali consumption approximately
1?$ less than twice the amount of ligand taken. This
reduction in the expected titer of alkali consumed can
be rationalized on the assumption of partial formation
of cupric-thiouracll complexes with a stoichiometry of
mu2 (XIV).
If the formed complexes were only MUOH and/or its
polymers then the titer of free, uncomplexed ligand,
measured between inflections 1 and 2 of curves D, E and
F of Fig. 15 and inflections 2 and 3 of curves D, E and
F of Fig. 1, as identified by its pK*, would be equal to
Si
o
the difference between the millimoles of ligand (5*00 x 10 )

104
and the millimoles of cupric ion taken. This was not so
(Figs. 1, 15)» In each instance the titer of free, uncom-
plexed ligand was slightly less than would be expected
under the assumption of only MUOH and its polymer formation
and supported the assumption that partial formation of MUg
occurred which lowered the amount of free thiouracil. It
would be expected that, under more acidic conditions where
the availability of hydroxyl ions is further reduced,
formation of the 2:1 (MU2) complex would predominate.
This formation of an MUg complex argued for by the titra¬
tion curves did not occur in the preparative procedures.
This is probably due to the higher temperatures employed
which allowed the equilibrium to shift in favor of the
MUOH (XIX) and its derived complexes as precipitation
proceeded.
The sharpness of inflections 2 and 1 in curves C of
Figs. 1 and 15, respectively, indicated that the precipi¬
tated complex was not readily destroyed even at relatively
high pH values. This may be due to possible inertness of
the complex (i.e., the rate of achievement of a new equi¬
librium is slow) but it is more probable that the rate of
dissolution was relatively slow. The fact that precipitate
was still present favors the latter interpretation.
Solutions of cupric ion have been shown (76) to exist
with a twin hydroxo bridge structure.

105
XXIII H +2
Cu(^ ^Cu
H
This information is consistent with and analogous to the
structure of the isolated complexes of 2-thiouracil(XX),
the bis (2-thiouracil )-ji-dihydroxodicopper (II), which on
drying and loss of a mole of water would give XXI. The
formula XXI would have the elemental analysis and the
spectrally observed carbonyl bonds. The extreme insolu¬
bility of the complexes frustrated efforts to obtain
molecular weights, nuclear magnetic resonance spectra
and mass spectra.
II. Structure and Stability of Cupric-Thiouracil
Complexes in Homogeneous Solutions
Spectral evidence for the presence of MU^ Cupric Com¬
plexes . Cupric ion modified the ultraviolet absorbance
of the thiouracils (Table III, Fig. 2). This indicated
the formation of a complex and supported the evidence of
the potentiometric titration data which showed a large,
immediate drop in pH when cupric ion and ligand were mixed.
The observed isosbestic points, 25^ and mjx, were strong
evidence for a 1:1 transformation. The decrease in absorb¬
ance near 3^0 mji (Fig. 2). The difficulty in determining
the absorption coefficients of the intermediate complexes
and the fact that slow precipitation occurred in the cells
frustrated determination of the stability constants.

106
Ultraviolet spectra of solutions prepared according
to Job's method (72) of continuous variations (Fig. 7,
Table XII) gave absorbance values at 272 and 3^5 mp which
were plotted against the mole fraction of ligand (Fig. 8).
The maxima of curve A (3^5 mp) and curve B (272 mp) occurred
at a mole fraction of 0.66 of 2-thiouracil. This indicated
a complex with a stoichiometry of MUg.
Ultraviolet spectra (Fig. 26) of solutions with con¬
stant concentrations of 2-thiouracil and equimolar or less
concentrations of cupric ion gave absorbance values at the
272 mp maximum which decreased as the concentration of
cupric ion approached a concentración equal to half the
concentration of 2-thiouracil used. At cupric ion con¬
centrations equal or greater than half the concentration
of ligand the absorbance was constant.
The absorbance at 3^5 mp (Fig. 26) increased as the
concentration of cupric ion approached half the concentra¬
tion of ligand and remained constant above the half-equi¬
molar concentration.
Plots of the absorbances from Fig. 26 at 272 and 3^5
mp versus the ratio of cupric ions to 2-thiouracil concen¬
tration deviated from a constant absorbance at a ratio of
0.50 (Fig. 9) which is indicative of a stoichiometry of
MUg for the cupric complex in homogeneous solution.

107
The stoichiometry of MUg from the Job’s and mole ra¬
tio plots agrees with the potentiometric titration evidence
for the presence of the MUg complex in Scheme IV (XXII).
These facts imply that the MUg (XXII) complex is readily
formed in solution but the higher insolubility of the pre¬
cipitated complex (XX) produces forcing conditions that
destroy it.
Identity of cupric complexes present in acid solution
by polarography. Since precipitation of the cupric com¬
plexes of the thiouracils prevented the calculation of
their stability constants from titration and spectral data,
the polarographic method of analysis was used. Solutions
containing ligand and cupric ion and 0.0500 to 0.2648 molar
in perchloric acid did not precipitate and could be ana¬
lyzed according to Eq. 66. Typical dataare given in Table
VI and a typical polarographic curve is given by Fig. 14.
The slopes of the plots of Eq. 66 (typical plots are
given by Fig. 28) could be fitted best by the theoretical
slopes for the cases of 2-thiouracil, 6-n-propyl-2-thioura-
cil and 5.6-dimethyl-2-thiouracil. These slopes require a
value of unity for P-Q/n from Eq. 66. The values of
Ej/^-Edetermined from the polarographic curves were
within ten mV of the theoretical value calculated by Eq. 6
for a one electron change. Since the value of n, the elec¬
tron change, is unity then the value of P-Q, the differ¬
ence in the number of ligands between the oxidized and
reduced complex, must also be unity.

108
The fact that a one electron change occurred in the
reduction at the dropping mercury electrode means that
the cupric ion was reduced only to the cuprous state and
not to the free metal. This can occur only if the thioura-
cil complex with cuprous ion was very stable and thus pre¬
vented the reduction of the cuprous ion. The prevention
of the reduction of cuprous ion by strong complexation is
not new and occurs in the presence of cyanide ion (77)»
Cuprous ion is a very "soft acid" (78) and would be ex¬
pected to strongly complex with sulfur, a very "soft base"
(78), and prevent the reduction of cuprous ion to copper
metal.
The value of unity for P-Q and n required that the
electrode reaction was the reduction of the 2:1 (MUg)
(XXII, Scheme IV) complex of cupric ion to the 1:1 (MU)
complex of cuprous ion. Since values of the dissociation
constants for the cupric or cuprous complexes are not
known, only the ratio of the dissociation constants of
the cupric and cuprous complexes could be obtained. These
ratios are given in Table III.
In some cases the slopes of the plots of Eq. 66 were
zero which indicated a value of zero for P-Q (Table XIII).
This result means that there is no change in the number
of ligands bound to cupric ion when it is reduced to cu¬
prous ion. Since this occurred at the higher temperatures,

109
It can be explained by the postulate that with cupric ion
under these conditions, the reduction of the 1:1 complex
(XIII, Scheme IV) is preferred with retention of the bound
ligand.
Accuracy and sensitivity of polarographic measurements.
The plots of Fig. 28 show what appears to be a relative¬
ly large degree of scatter in the points defining the
straight lines. This is partly due to the restrictions
placed on the available range of ligand concentrations
by their low intrinsic solubility (Table IV). However,
the accuracy and sensitivity of the data are better than
they appear to be. Other studies in the literature (79)
with larger ranges of ligand concentrations and greater
changes in half-wave potential show more scatter in the
data than in the present study. In any case the error
in the points approaches the + 2.5 mV which is normal
for the Sargent Model XV Polarograph (41).
Effect of temperature and acid concentration on the
negative logarithm of the ratio (KQ/Kr) of the cupric to
cuprous dissociation constants. The values of the log
ratios in Table XIII indicate that the overall dissociation
constant of the 2:1 cupric complex is very much smaller
than the dissociation constant of the cuprous complex.
This is not surprising when the larger positive charge
of the cupric ion is considered.

110
An increase in temperature caused a decrease in the
log dissociation constant ratio (Table XIII). An increase
in temperature is expected to increase dissociation reac¬
tions. Thus, the observed decrease can be attributed to
a larger increase in the dissociation constant of the
cuprous complex relative to the increase in the dissocia¬
tion of the cupric complex.
If we ignore the temperature effects and deal only-
in orders of magnitude the differences between the log
ratios for the ligand dependent (P-Q=l) and independent
plots (P-Q=0) is approximately lCp (Table XIII). This
value can be taken as an approximation of the negative
logarithm of the dissociation constant of a single ligand
from the 2:1 (MUg) cupric complex. The validity of this
comes from the fact that KQ is the overall dissociation
constant for the cupric complex (Eq. 46) which is the pro¬
duct of the dissociation constants for the loss of the
first and second ligand. In changing to a ligand-indepen¬
dent condition the value of -log K /K would differ from
or
the ligand-dependent condition only by the value of the
negative logarithm of the first dissociation constant of
the cupric complex.
The change from a ligand-dependent to ligand-indepen¬
dent condition occurs over a relatively small temperature
and acid concentration range. It would be expected that
a slower transition from a finite to zero slope would be

Ill
observed but this was not the case. The data in each
instance was fitted best by either the theoretical finite
or zero slope. A postulated mechanism which might account
for the observed behavior is that the diffusion of the 1:1
cupric complex (MU+) up to the electrode may be much faster
than the diffusion of the neutral, MUg species. If this
were the case, the lowering of the concentration of the
MU+ species just outside the diffusion layer would be ex¬
pected to cause a further increase in the dissociation of
the MUg species to MU+ and could conceivably prevent any
of the MUg complex from ever arriving at the electrode
surface and being reduced. The postulated mechanism is
an "all or none" type and would be expected to yield the
type of data observed.
III. Cadmium and Lead Complexes of Thlouracils
Structure and stability constants of cadmium and lead
complexes of thiouraclls. Solutions of lead or cadmium
nitrate and various thiouracil ligands (2-thiouracil, 6-n-
-propyl-2-thiouracil, 6-methyl-2-thiouracil, 5-methyl-2-
-thiouracil, 5.6-dimethyl-2-thiouracil and 5-carboethoxy-
-2-thiouracil) could be titrated with standard alkali
until precipitation occurred (Fig. 16).
The homogeneous portion of the titration curves up
to the point of precipitation (Fig. 16) was analyzed ac¬
cording to Eq. 1? and gave stability constants (K^) for

112
4
1:1 complexes (Table II) of the order of 10 which were
independent of the metal ion concentration. The validity
of Eqs. 26 and 38 is denied (56) since they relate the
ligand anion concentration to experimental quantities
for mixed ligand and polynuclear complexes, respectively,
and are both functions of the free metal ion concentration.
The species present in homogeneous solution must be the 1:1
(MU+) and 2:1 (MUg) complexes (Scheme V) since typical
plots of Eq. 1? (Fig. 21) show finite intercepts.
Analysis of the homogeneous solutions by Eq. 18 (Table
I), which assumes the presence of the MU+ species only,
gave plots (typical Figs. 17-20) which were linear over
the major part of the data. The slopes of the plots (Table
I) varied between a value of one and two as a function of
metal ion concentration. Those cases where the initial
metal ion concentration ([M Jq) was equal to the initial
ligand concentration (¡_HUJq) usually gave slopes of unity
while metal ion concentrations less than equimolar gave
slopes larger than unity but no greater than two. Slopes
greater than one and less than two indicate the presence
of complexes with a stoichiometry of MU2. Since the log
values are usually greater than the log K2 values,
it is expected that the formation of MU+ complexes
(XXIII, Scheme V) would be preferentially formed when the
ligand concentration was insufficient for complete MU2
formation. Equimolar concentrations of ligand and metal
ion would then result in predominant formation of the MU+

113
Scheme V

114
complex and plots of Eq. 18 would give slopes of about
unity. This was observed to be the case (Table I).
Ultraviolet absorption spectra of ligand solutions
containing lead and cadmium ion showed deviations due to
complex formation as expected (Table III). Typical spectra
are given in Fig. 3» The absorbance at the 272 mji peak
was reduced but the absorbance decrease was not enough
to allow a sensitive measurement of the stability constants.
A disadvantage of the use of spectral data for the measure¬
ment of stability constants is the necessity of determining
the molar absorption coefficients of the complexes (?4).
The absorption coefficient of the intermediate complex
(MU+) is very difficult to evaluate and in view of the
small changes in absorbance with large changes in metal
ion no attempt was made to calculate stability constants
by spectral measurements.
Titration of the ligand-metal solutions past the point
of precipitation gave curves which indicated a consumption
of alkali equal to twice the amount of ligand taken (curve
A, inflection 1, Fig. 16). This indicates that the complex
being precipitated is different from the compounds MU+
(XXIV, Scheme V) and MUg (XXII, Scheme V) formed in solu¬
tion and either has a stoichiometry of MUCH or is pro¬
ducing two titratable protons per ligand. This behavior
required that the complex species (MUg) present in homo¬
geneous solution was not the species being precipitated.

115
Consideration of curve A, Fig. 16 shows that the pre¬
cipitated complex is not disrupted by hydroxide ion even
at relatively high pH values, otheritfise the titer of alkali
consumed to inflection 1 would be greater than twice the
titer consumed by the titration of the ligand alone (curve
G). Thus, there is no apparent tendency for the precipi¬
tated complex to be disrupted by high concentrations of
hydroxyl ions, at least to a pH of about 9.0.
Solutions containing less than equimolar amounts of
metal to ligand showed the presence of uncomplexed, free
ligand. The titer of the uncomplexed ligand, measured
between inflections 2 and 3 of curves D, S and F of Fig.
»
16, and having a pK the same as the titration of the
o.
ligand alone, also indicated the precipitation of a complex
with a stoichiometry of 1:1, metal to ligand, or some
multiple of 1:1. Thus, a given increase in the moles of
metal ion used resulted in an equal decrease in the moles
of ligand left uncomplexed. Further proof of the precipi¬
tation of a complex with a stoichiometry of 1:1 is that
a slope of one was obtained when the moles of metal ion
used was plotted against the moles of free, uncomplexed
1igand.
The precipitates produced by titration of solutions
containing 2-thiouracil in the presence of lead and cadmium
ions were isolated, washed and the infrared spectra re¬
corded. The precipitate produced from the lead solution

116
was isolated at pH 7*5 and that from the cadmium solution
at pH 10.5. In both cases the infrared spectra were identi¬
cal to the spectra of the synthesized complexes with a 1:1
stoichiometry indicating MU or some multiple.
Synthesis of cadmium and lead complexes of thlouracils.
The equilibria in Scheme V are based on the law of mass
action and the relative solubility of the 1:1 and 2:1 com¬
plexes as determined by titration. Deliberate attempts to
test Scheme V were conducted by preparation of MU^ (2:1)
complexes using forcing conditions of a high ligand con¬
centration and addition of metal ion to ligand and by
preparation of complexes with a stoichiometry of 1:1 using
conditions of equal concentrations and addition of ligand
to metal ion.
The infrared spectra, of prepared bis(2-thiouracil )-
cadmium(Il), bis(2-thiouracil )lead(II) and bis(6-n-propyl-
-2-thiouracil )cadmium( II) (MU,,) prepared under special
conditions of concentration and order of addition as out¬
lined in the experimental section, showed the presence of
a strong absorption at about I63O cm.- This is in the
region expected for the absorption of the carbonyl group
at the four position of thiouracils (80). The infrared
, 1
spectrum of 6-methyl-N,N -diethyl-2-thiouracil contained
— 1
no bands between 1440 and 1660 cm.- , as expected, which
is where the C=N ring absorption has been show, to occur
(81). The infrared spectra of these complexes showed the

presence of strong absorption bands between 1440 and 1660
cm. , indicating that C=N bonds were present and that the
nitrogens were not bound to the metal ion. Elemental
analysis, given in the experimental section supported the
assigned structures (XXII, Scheme V). This evidence of
presence of C=N and C=0 bonds and elemental analysis is
conclusive proof that the structure of these complexes,
when prepared under the designated forcing conditions,
is the 2:1 complex represented by XXII, Scheme V.
The infrared spectra of bis(2-thiouracil-cadmium(II))
bis(2-thiouracil-lead(II)) and bis(6-n-propyl-2-thiouracil
-lead(II)) (MgU2) had infrared spectra with no band above
1570 cm.” This means that the carbonyl band near I63O
_ 1
cm. (80) was not present and that the carbonyl group at
the four position of the thiouracil ring did not exist.
The bands assigned (81) to the C=N group did exist however
This information coupled with the elemental analysis given
in the experimental section further supported the assigned
polynuclear, cyclic structure (XVII, Scheme V).
Further evidence for the assigned structure was de¬
duced from the fact that the pK of 2-ethylmercapto-4-
3,
-hydroxypyrimidine is 7.01 compared to 7.46 for 2-thioura.-
cil. The effect of alkyl substitution at the sulfur atom
tended to make the potential hydroxyl group at the four
position a stronger acid. Substitution by a cation on
the sulfur as in MU+, would be expected to cause an even

118
larger increase in the acid dissociation of the 4-hydroxy
group. This postulated dissociation of the 4-hydroxy group
will produce a source of negative charge which can attract
previously complexed cations. Since the pK* of the 4-hy-
9
t
droxy group in the cation is much lower than the pK of
9
the parent thiouracil, the concentration, in the region
of pH 5 to 6, of the species dissociated at the four posi¬
tion will increase much faster than the concentration of
the anion of the parent ligand and very much faster than
the hydroxyl ion content of the solution.
Recrystallization of the polynuclear complex of cad-
mium-2-thiouracil (XVII) from ammonia solution gave a sub¬
stance whose infrared spectra was almost exactly the same
as for the original complex (XVII) but whose crystal struc¬
ture was decidedly polymeric in nature. The crystals
looked as if they had been converted to paper pulp. It
is postulated that the ammonia complexed with the cadmium
and opened the ring structure of XVII. On removal of the
ammonia the material took on the polymer structure (SVIII).
The effect of the relative solubility of the species
actually precipitated from solution compared with the
solubility of the complex in homogeneous solutions was
important. Both the polynuclear comples (MgUg) and the
2:1 complex (MU^) were formed from the same intermediate
(MU+). As long as homogeneous solution conditions prevail,

119
it was expected that only MU+ and MUg complexes exist.
As soon as the dissociation of the 4-hydroxyl group on
MU+ begins to give MU, then *'orms w^lch soon exceeds
its solubility and precipitates. The precipitation upsets
the equilibrium in favor of the formation of more MU and
of the complex that has precipitated. Thus, the complex
remaining in solution (MU,,) can never reach its solubility
limit and therefore, only one complex would precipitate.
This behavior was seen in the case of the cadmium and
lead complexes where MUg complexes were initially formed,
as determined by potentiometric titration and Eq. 17, but
only polynuclear complexes having the ring structure and
a stoichiometry of MgU2 were found in the precipitate.
The one exception was the case of cadmium-6-n-propyl-
-2-thiouracil in which the 2:1 complex precipitated instead
of the polynuclear complex (Fig. 22). This means that the
solubility of the 2:1 complex in this case was lower than
the solubility of the polynuclear complex and precipitated
first.
Special case of titration of 6-n-propyl-2-thiouracll
solutions containing cadmium ion. The potentiometric
titration of solutions of cadmium and 6-n-propyl-2-thioura-
cil at 25*0° and 35*0° did not exhibit the consumption of
alkali necessary to produce mixed ligand (MUOH) or poly¬
nuclear (MgUg) complexes as x^ras observed in the cases of

120
lead and cadmium with the other ligands. Inflection 1,
curve A of Fig. 22 occurred at an alkali titer equal to
the titer consumed by the ligand alone (curve D). This
occurred at all cadmium ion concentrations equal or
greater than txiice the concentration of ligand taken.
When the metal ion concentration was greater than twice
the ligand concentration the formation of metal hydroxides
could be observed by titration beyond the inflection 1 of
curve A. When the concentrations of metal and ligand were
equal (1.00 x 10 J M) there was an inflection (curve B,
Fig. 22) which was larger than that found in curve A and
which also occurred at an alkali titer equal to the titer
of ligand taken. This occurred because all of the metal
ion xtfas complexed and none was available for formation of
metal hydroxides. Only when the amount of metal ion taken
was less than half the amount of ligand was there any free,
uncomplexed ligand observed (curve C). In this case the
precipitated complex vías not the mixed ligand or polynu¬
clear complex but was the simple 2:1 complex (MUg) which
existed in homogeneous solution.
Effect of ligand structure and temperature on stabil¬
ity constants of cadmium and lead complexes. Among the
ligands which complex cadmium and lead strongly (2-thioura-
cil (2TU), 6-n-propyl-2-thiouracil (PTU), 6-methyl-2-thi-
ouracil (6MTU), 5-methyl-2-thiouracil (5MTU), 5»6-dimethyl-
-2-thiouracil (5> 6DMTU), and 5-carboethoxy-2-thiouracil

121
(5CETU)), the stability constant for the formation of MU+,
K^, is largest for 56DMTU and smallest for 5CETU (Tables
I and II). The values of log Kg, the constant for the
formation of MUg, are more variable and are less accurate.
However, in most cases the log Kg values are largest for
the 5.6DMTU complexes also. The stability constants for
the cadmium and lead complexes of the other ligands (PTU,
2TU, 6MTU and 5MTU) are grouped fairly close together and
exhibit no outstanding differences. The effect of alkyl
substitution of thiouracil at the five and six positions
on the stability constant for the formation of MU+ can be
most clearly seen in Fig. 24. The value of F is largest
for Pb-5¿>DMTU among the lead complexes and is largest for
Cd-5^DMTU among the cadmium complexes. The value of -foF
for the Cd-5CETU complex has the smallest value and repre¬
sents the least stable complex.
It appears that simultaneous alkyl substitution at the
five and six position of thiouracil may increase the abil¬
ity to complex metal ions. Single and no substitution at
the five or six positions make little difference in the
ability of the ligand to complex metal ions.
Substitution of the electronegative carboethoxy group
at the five position would be expected to reduce the elec¬
tron availability at the sulfur atom, para to the five
position, and thus reduce the sulfur's electron charge

122
and its ability to attract and bind to cations. This re¬
duced ability to attract and bind metal cations was re¬
flected in the lower stability constant (Table II).
The effect of temperature on the value of log is
given by Fig. 24. Due to the requirement of constant
when plotting the change in free energy versus the absolute
temperature, the range of temperature used (20°) was pur¬
posely restricted (71). The slopes of Fig. 24 (^S) and
the values of AH calculated from Eq. 80 are given in Table
XI. Although the error in the values of As are relatively
large certain trends are evident.
The values for ¿H, the change in enthalpy for forma¬
tion of the MU+ complex, are generally negative especially
in the case of the lead complexes. A negative value of
AH indicates release of energy and formation of a stable
bond while positive values of indicate a higher stabil¬
ity of the complex. The AH values for the cadmium com¬
plexes of 6MTU and 5CETU are slightly positive but the
corresponding values of As are also positive and larger
than the others. The value of As for the cadmium complex
of 6MTU probably has a large error in view of the lack of
strict linearity of the plot in Fig. 24. The even larger
value of As for the cadmium complex of 5CETU appears to
be real however, and could be due to the fact that 5CETU
is the most soluble of the ligands (Table IV) and therefore,
the large entropy factor may represent a rearrangement of

123
the water of hydration. All of the values of AS for com-
plexation are positive except for the value for the lead
complex of PTU. The positive values may reflect the loss
of water from the hydration sphere of the metal ions when
complexation occurs.
The larger stability constants for the lead complexes,
compared to the cadmium complexes (Table II) are seen to
result from the larger, more negative values of the enthal¬
py (Table XI). In the case of the lead complex of PTU a
negative value for AS was obtained. A possible rationali¬
zation may be that it is due to restriction of the longer
n-propyl side chain which is the only feature distinguish¬
ing it from the other ligands.
The acid dissociation constants (Table X) of the
ligands and the associated thermodynamic parameters (Table
XI, Fig. 25) are related to the stability of the cadmium
1
and lead complexes. As the pKa of the ligand rises the
stability of the corresponding lead and cadmium complex
also rises. This relation reflects the increasing negative
charge density of the sulfur which results in a greater
attraction for cations in general. The negative values
of the entropy change, except for PTU which is anomalous,
are assigned to the increased ordering of the hydration
sphere induced by the production of ions.
Alkylation of the sulfur with an ethyl group as in
2-ethylmercapto-4-hydroxypyrimidine destroyed the ability

124
to complex metal Ions. Since the potential hydroxyl group
at the four position still existed but no complexation was
observed, it can be asserted that the binding of metal ions
by the oxygen at the four position in the parent thioura-
cils does not exist.
»
Alkylation of the N-l and N-3 nitrogens, as in N,N -
-diethyl-6-methyl-2-thiouracil also destroyed the ability
of the thiouracils to complex metal ions. Substitutions
on both the ring nitrogens or the sulfur would prohibit
the tautomeric formation of an anionic sulfur. This is
further evidence that complexation occurred at the sulfur
atom and was dependent on the formation of the sulfur
anion.
IV, Complexation of Other Metal Ions with Thiouracils
Potentlometric titration of thiouracil solutions con¬
taining nickel or zinc. Potentiometric titration of li¬
gand solutions containing nickel or zinc gave very similar
curves (Fig. 23)» The homogeneous part of the titration
curves, up to the point of precipitation as indicated on
Fig. 23. was analyzed using Eqs. 17 and 18 and gave values
of log and log Kg which were approximately 2 units
smaller than those observed for the cadmium and lead cases
(Tables I and II).
The titer of alkali consumed to inflection 1 of curve
A was greater than twice the titer of alkali consumed by

125
the ligand alone. This indicated that the complex was not
stable at high pH values and was disrupted by hydroxyl ions
to give the nickel and zinc hydroxides. This is not unex¬
pected since the stability constants of the nickel and zinc
complexes are about a hundred times smaller than those for
lead and cadmium.
Potentiometric titration of thiouracils in the presence
of other metal ions. Potentiometric titration of the
ligands in the presence of ferric, ferrous, cobalt, calcium
and manganese ions gave no indication of any complex forma¬
tion. Further evidence for the lack of complexation by
ferric and ferrous ions was given by the ultraviolet spec¬
tral data.
Spectra of ligands in the presence of ferric ion (Table
III, Fig. 4) showed no evidence of complexation. The spec¬
tra of the mixtures were simply the sums of the absorbances
of the ligand and metal ion. This indicated no interaction
and hence, no complexation. The plotted points of Fig. 4
are the summation of curves 1 and 2 and fall on the spec¬
trum of the mixture (curve 3)» Spectra of ferrous ions
gave the same results as was found for ferric ion (Table
III). These results support the results of the potentio¬
metric titration data.

126
Spectra of sterically hindered ligands In the presence
of metal Ions» Substitution of an alkyl group on the
sulfur of 2-thiouracil would be expected to block the forma¬
tion of metal ion complexes if the anionic sulfur form of
thiouracil is the complexed species. Ultraviolet spectra
of 2-ethylmercapto-4-hydroxypyrimidine in the presence of
cupric, lead, cadmium, ferrous and ferric ions showed no
change from the spectra obtained in the absence of metal
ions (typical Fig. 6, Table III). The lack of change
indicated no interaction of the ligand with metal ions
and was good evidence for the lack of complex formation.
Formation of the anion of thiouracil can also be blocked
by substituting alkyl groups for the dissociable hydrogens
on the nitrogen atoms at positions one and three. When
this was done, as in N,N -diethyl-6-methyl-2-thiouracil,
the ultraviolet spectra of the molecule in the presence
of cupric, lead, cadmium, ferrous and ferric ions (typical
Fig. 5» Table III) showed no change from the spectra of
the ligand in the absence of metal ions. This indicated
no complex formation and was further support for metal
complexation at an anionic sulfur in the sterically free
ligands.
Relative tendency of metal ions to complex with
thiouracils. The order of decreasing stability of metal
complexes of thiouracils is
Cu+2 » Pb+2 > Cd+2 » Ni+2 Zn+2.

12?
Metal ions which do not complex thiouracils are
Fe+3, Fe+2, Co+2, Ca+2 and Mn+2.
The "natural order" (82) for the first row transition
metals, Ni+2, Cu+2, Zn+2, shows the highest complexation
stability at cupric ion and is a direct consequence of
their ionic radii (83) and follows the order of the
heats of hydration (84).
When ligands approach cupric ion, the energy of the
normally degenerate (having equal energies) 3d electron
orbitals is changed, causing some orbitals to become
lower in energy and some higher (85). The nine 3d
electrons of cupric ion will preferentially fill the
lower energy orbitals first and will give a decrease
in the total energy which is called the Crystal Field
Stabilization Energy (CFSE). The CFSE is theoretically
Q
a maximum in a d ion in an octahedral field but in the
case of cupric ion the d 2 and d 2 2 orbitals are
"2-i A " jr
degenerate (83). The degeneracy is removed by
Q
Jahn-Teller stabilization causing the 3d configuration
to become most stable (86). Jahn-Teller stabilization
results from the fact that if the ground level of a system
is composed of degenerate energy states, a distortion of
the system must occur to remove the degeneracy and make
one energy state lower. In the case of cupric ion there
are two degenerate energy states and one is raised and
the other is lowered in energy. This change in energy

128
states results in an axial distortion of the octahedral
cupric ion to a square-planar configuration (86).
The 3d orbitals of sulfur are empty and may accept
electrons by back donation from the partially filled 3d
orbitals of cupric ion to form it bonds thus adding to the
strength of the complex. The complexes of cadmium and
lead may owe part of their relatively strong bonding to
back donation of electrons from their filled 4d and 5d
orbitals, respectively, to the empty 3d orbitals of sulfur
(87).
Pearson (78) has classified metal ions and ligands
in terms of "soft" and "hard" acids and bases. In this
classification those metal ions which complex by ionic
attraction are called "hard" acids while those metal ions
are called "soft" which complex primarily by formation of
covalent bonds.
Sulfur anions bound to conjugated systems as in thi-
ouracil and thiophenol anions rank near the top as soft
bases (85). According to Pearson’s classification a soft
base forms the strongest bonds with soft acids. Cadmium
(78) and silver ions (78, 88) are typical soft acids while
calcium, manganese and ferric ions are typical hard acids
(78). Lewis acids that fall into a borderline class are
ferrous, cobaltous, nickel, cupric, zinc and lead ions.
The stability of ferrous and cobaltous ion complexes would
be expected to be loxver than found for nickel since their

129
ionic radii are larger and heats of hydration are lower.
Lead ion is also a borderline case and probably owed its
relatively high stability constant with thiouracils to
its large size and polarizability.
V. Biological Effects
Effect of thlouracll structure on antithyroid activ¬
ity. Uracil, the oxygen analog of thiouracil, has no
antithyroid activity (89). This establishes that a sulfur
at the two position of thiouracils does promote anti¬
thyroid activity.
Substitution on the sulfur of 2-thiouracil abolishes
its antithyroid activity (10) and preliminary studies
showed that 6-methyl-N,N -diethyl-2-thiouracil has minimal
activity (Table VIII). Substitution of a nitrogen for the
carbon at the six position of 2-thiothymine(5-methyl-2-
-thiouracil) gives 2-thio-6-azathymine which has appreci¬
able antithyroid activity (90). When the N-l and N-3
nitrogens of 2-thio-6-azathymine are substituted with
methyl groups all antithyroid activity is lost (90). This
is consistent with the lack of antithyroid activity of
6-methyl-N,N -diethyl-2-thiouracil (Table VIII).
Thiouracils with sterically blocked sulfur atoms do
not complex with the tested metal ions and do not have
antithyroid activity. The implication is that complexation

130
or the molecular parameters that control complexation are
important to the antithyroid activity of thiouracils.
Thus, the presence of an electronegative carboethoxy group
at the five position of thiouracil, which would reduce the
electron density at the sulfur, reduces the stability of
metal complexes and drastically reduces the antithyroid
activity of 2-thiouracil (10). Other thiouracil deriva¬
tives substituted at the five position with electronega¬
tive groups also have much reduced antithyroid activity.
Such compounds are 5-cyano-2-thiouracil and 5-carbcxy-2-
-thiouracil (10) and would be expected to reduce the
electron density at the sulfur atom.
A possible mechanism of antithyroid activity of
thiouracils is their oxidation to the disulfide by iodine
in the thyroid gland (91)» Since oxidation is the loss
of electrons, it would be expected that substituents
which increase the electron density of sulfur would
also increase the antithyroid activity. However, this
increase in electron density is also what would be
expected if complex formation were the mechanism.
If we assume that either complexation or oxidation
is the antithyroid mechanism of thiouracils, then the
pK values of the ligands (Table X), which are a measure
Si
of the sulfur electron density and which are related to
the log values for the complexes, could be used to

131
predict the antithyroid activity of thiouracil derivatives.
However, accurate data on ED^q (effective dose at 50 per¬
cent of maximal response) values of the thiouracils must
first be obtained.
Relation of stability constants with antithyroid
activity. Since the relative intrinsic activity is the
ratio of the maximal thyroid response of the drug to
2-thiouracil and was independent of dose, it was felt
that this ratio could be calculated with greater
confidence than the Apparent Relative Activity (Tables
VII-IX) which would be dose dependent. (See Results.)
Attempted correlation of log values for the
cadmium and lead complexes with the Relative Intrinsic
Activity (RIA) gave straight lines which were parallel
to the Relative Intrinsic Biological Activity axis
indicating that Relative Intrinsic Biological Activity
is independent of the stability constants. The values
of the negative logarithm of the ratio of the dissocia¬
tion constant of the cupric complex to the dissociation
constant of the reduced complex for the cupric complexes
were also independent of the Relative Intrinsic Biological
Activity. The maximum biological response, as measured by
the Relative Intrinsic Biological Activity, is a result of
the saturation of receptor sites, and is a parameter that
is independent of concentration.

132
It may be argued that If the strength of the drug
binding, in an occupied receptor site, were related to
the intensity of the biological response, the Relative
Intrinsic Biological Activity would correlate with the
stability constant. This argument is denied by the
lack of correlation.
The limitations of the calculated Apparent Relative
Activity (ABA) constants that were used have already been
mentioned in the section on results. However, if proper
data were available, they may be a better parameter for
correlation with the stability constants since they are
concentration dependent and the mechanism of drug action
may involve a transfer or chemical process that is con¬
centration dependent.
Since a possibility existed that there may be a
relationship between the negative logarithm of the ratio
of the dissociation constant of the oxidized complex
to the dissociation constant of the reduced complex
with the Apparent Relative Activity, a plot (Fig. 29)
was constructed using the data from the literature (10)
as well as that found in Tables VII and IX. Fig. 29
suggests a relationship although it is recognized that
the slope of the line depends largely on the position
of the point for 5-carboethoxy-2-thiouracil. This rough
relationship suggests that further studies along these
lines may prove fruitful.

Relation of Intrinsic solubility of ligands to their
relative antithyroid activity. In addition to the pos¬
sibility that the negative logarithm of the ratio of the
dissociation constant of the cupric complex to the dis¬
sociation constant of the cuprous complex may relate to
biological activity of antithyroid compounds, the possi¬
bility that intrinsic solubility plays a role in biologi¬
cal activity should be considered. The thiouracil deriva
tives which have the highest Apparent Relative Activity
also have the highest intrinsic solubility (Table IV).
The water solubility of these drugs may be one important
aspect of their activity since their transport through
the body depends on this established concentration
gradient and is necessary to give a concentration at the
biophase sufficient to elicit the desired biological
response. The lower the solubility of these drugs, the
greater the limitations on the transport of drug to the
receptor site.
For example, the presence of an amino group at the
six position of 2-thiouracil might be expected to cause
appreciable antithyroid activity by virtue of the probabl
increased electron density at the sulfur. However, the
antithyroid activity of ó-amino-2-thiouracil is very low
(10). This compound also has the lowest intrinsic

134
solubility which would be expected to contribute to its
low Apparent Relative Activity by reducing the concen¬
tration available to the thyroid gland.
However, it is realized that the transport of drugs
through the body is not solely a function of their solu¬
bility and their related rates of solubilization prior
to absorption. Partition coefficients, degrees of pro¬
tein binding, crystalline polymorphic forms, all of these
items may have significant additional effects on the
transport of the thiouracil to its site of action.

SUMMARY AND CONCLUSIONS
1. Metal complexes of thiouracils substituted with
alkyl groups at the five and six positions have been shown
to be most stable in the case of cupric ion, less stable
in the case of cadmium and lead and least stable in the
cases of nickel and zinc. The complexation of thiouracils
by ferric, ferrous, cobalt, calcium and manganese does not
occur.
2. The structures of precipitated cupric complexes of
thiouracils, in the pH range three to five, were determined
by potentiometric titration, infrared spectra and elemental
analysis and have the structure
H
135

136
3. Cupric complexes in homogeneous solutions have the
structure
- Cu -
hr
as determined by ultraviolet spectra, potentiometric
titration and polarography, up to a pH of at least three.
At temperatures near 40° there is evidence that the 1:1
complex may be the principal species in solution.
4. The negative logarithm of the ratio of the dis¬
sociation constants of the 2:1 cupric and 1:1 cuprous
thiouracil complexes in dilute acid solution, have been
determined by polarography and have values in the range
ten to eleven.
5. The cadmium and lead complexes of thiouracils in
homogeneous solution have stability constants of about
4 3
10 for , lO-' for K„ and a structure of

137
for the fully complexed metal ion which is in equilibrium
with the 1:1 complex.
6. The precipitated cadmium and lead complexes of
thiouracils have been shown to exist in two forms. High
concentrations of ligand produce complexes of the form MU?,
while high concentrations of metal ions in the presence of
thiouracils at pH values above 6.5 give insoluble complexes
with the structure
7. Determination of the effect of temperature on the
stability constants of the cadmium and lead complexes gave
values of the thermodynamic parameters enthalpy and entropy
which indicated that the stability of the complexes is
directly related to the electron charge density on the
sulfur atom of thiouracils. Substitution at the five
position Tvith an electronegative carboethoxy group signifi¬
cantly lowered the thermodynamic stability constant of the
complexes while 5»6-diraethyl-2-thiouracil had the highest
stability constant.

138
8. The differences in the stability constants of the
thiouracil complexes of the various metal ions have been
accounted for on the basis of the pKa of the ligands and
on the electron structure of the ligand and metal ions
and indicate that the bonds between metal and sulfur are
principally covalent in nature.
9. These studies suggest a relationship between the
loss of ability of thiouracil derivatives to complex metal
ions and the loss of antithyroid activity, for example,
2-ethylmercapto-4-hydroxypyrimidine and 6-methyl-N,N -
-diethyl-2-thiouracil. Further studies should be made
to determine the validity of this suggestion.

APPENDIX
A. Tables

140
TABLE I—COMPOSITION OF 30LUTI0NSa AND ESTIMATED STABILITY
CONSTANTS13 AT VARIOUS TEMPERATURES ON THE POSTULATE OF 1:1
METAL COMPLEXES OF SUBSTITUTED THIOURACILS FROM POTENT10-
METRIC TITRATIONS
Ligand0
Metal
103[m+2]0
°c
Sloped
-4
10 ^
Log K1e
2TU
Ni+2
60.00
25.0
1.13
6.00
25-0
1.11
4.00
25.0
1.01
0.0269
2.43
1.60
25.0
1.08
1.00
25.0
1.20
Cd+2
2.00
25.1
1.00
1.62
4.21
1.80
25.3
1.36
1.60
24.9
1.12
1.40
25-3
1.23
1.20
25.1
1.36
1.00
25.1
1.22
0.80
25.1
1.59
0.60
25.1
1.33
0.40
25.0
1.60
0.20
25.0
1.60
Pb+2
2.00
25.0
1.00
4.83
4.68
1.80
25.2
1.15
1.60
25.2
1.15
1.40
25.1
1.11
1.20
25.0
1.11
1.00
24.9
1.08
0.80
24.9
1.12
0.60
24.9
1.10
0.40
24.9
1.08
0.20
24.9
1.07
PTU
Cd+2
1.60
25.3
1.43
1.00
25.3
1.66
0.40
25.3
1.42
0.20
25.3
2.06
Pb+2
2.00
25.8
1.00
6.60
4.82
1.60
25.8
1.00
5.90
4.77
1.00
25.8
1.12
0.40
25.8
1.24
0.20
25.8
1.13
Ni+2
2.00
25.9
1.91

141
TABLE I (continued)
Ligand Metal
103[M+2]q °C Slope lO-4^
Log K1
6mtu
5,6DKTU
5MTU
Cd+2
2.00
26.0
1.60
25-5
1.00
25-5
0.40
26.0
Ni+2
2.00
26.0
pb+2
2.00
26.0
1.60
25-5
1.00
25.5
0.40
25*0
Pb+2
2.00
25.8
1.60
25.0
1.00
25.3
0.40
25.3
Cd+2
2.00
25.5
1.60
25.5
1.00
25.3
0.40
25.5
Ni+2
1.00
25.5
pb+2
2.00
25.0
1.60
25.3
1.00
24.5
0.40
24.5
Cd+2
2.00
25.0
1.60
25.0
1.00
25.0
0.40
25.3
Ni+2
2.00
26.0
1.60
25.3
f
1.00
25.0
Pb+2
2.00
25.0
Cd+2
2.00
24.8
0.80
25.0
1.25
1.33
1.32
1.96
1.53
1.00
4.57
4.66
1.00
1.22
1.22
6.30
4.80
1.00
11.00
5.04
1.00
11.00
5.04
1.00
1.37
11.00
5.04
1.00
1.14
1.32
1.31
3-38
4.53
1.38
1.00
1.06
1.20
1.20
5.61
4.75
1.07
1.32
1.42
1.93
1.49
1.21
1.31
2.29
4.36
1.25
1.75
5CETU

142
TABLE I (continued)
Ligand
Metal
103[m+2]0 °c
Slope
t-.
0
1
•p-
Pi
f—1
Log K1
Pb+2
2.00
45.0
2.00
35.0
—
1.20
35.0
0.80
45.0
—
5CETU
Cd+2
2.00
35-4
1.54
2.00
44.85
1.73
0.80
34.5
2.28
0.80
44.85
2.03
6mtu
Cd+2
2.00
44.85
1.00
2.14
4.33
2.00
35*2
1.15
0.80
45.00
1.46
0.80
35.25
1.47
Pb+2
2.00
44.90
1.02
2.00
35.1
1.00
5.50
4.74
0.80
44.95
1.19
0.80
35.15
1.10
PTU
Cd+2
2.00
44.90
1.58
2.00
35.1
1.49
1.20
44.90
1.42
0.80
34.9
1.48
Pb+2
2.00
44.85
1.00
2.98
4.47
2.00
34.8
1.08
4.47
4.65
0.80
44.85
1.05
0.80
34.8
1.15
5,6DMTU
Pb+2
2.00
45.35
1.00
7.25
4.86
2.00
34.85
1.00
10.5
5.02
0.80
44.9
1.32
0.80
34.85
1.06
Cd+2
2.00
44.85
1.23
2.00
34.85
1.11
0.80
45.OO
1.67
0.80
34.95
1.33
2TU
Pb+2
2.00
44.85
1.00
3.47
4.54
2.00
34.95
1.10
3.80
4.58
0.80
45.15
1.16
0.80
34.85
1.10

143
TABLE I (continued)
Ligand Metal 103[m+2]q °C Slope 10~4K1 Log Kx
Cd
+2
5MTU Cd+2
Pb
+2
2.00
45.00
1.10
2.00
34.7
1.31
0.80
45.00
1.25
0.80
34.7
1.26
2.00
44.75
1.14
2.00
34.85
1.16
0.80
44.95
1.44
0.80
34.90
1.28
2.00
44.85
1.02
2.00
35*05
1.00
0.80
44.85
1.14
0.80
34.90
1.08
3.51 4.55
4.57 4.66
aFinal concentration of ligands was 0.002 M, ionic strength
was 0.006 M and the initial ¥0101116 was 25*00 ml.
â– u.
dK^ and log values were derived from the intercept
values of plots of log ^2- versus p[lf], where ñ is the
degree of formation of the complex, in accordance with
log —— = pK^ + p[U~] for those cases where the slope
was consistent with the theoretical expectation of unity.
cLigand abbreviations: (2TU) 2-Thiouracil, (PTU) 6-n-
-Propyl-2-Thiouracil, (6MTU) 6-Methyl-2-Thiouracil, (5.6DMTU)
5,6-Dimethyl-2-Thiouracil, (5MTU) 5-Methyl-2-Thiouracil,
(5CETU) 5-Carboethoxy-2-Thiouracil.
dSlope of plot of log ^-2- versus p[U-].
e ñ
Calculation of constants and slopes limited to data from
homogeneous part of titration curves.
immediate precipitation occurred.

144
TABLE II—COMPOSITION OF SOLUTIONS AND ESTIMATED LOGARITH¬
MIC STABILITY CONSTANTS, K1 AND K0, AT VARIOUS TEMPERATURES
ON THE POSTULATE OF MIXED 1:1 AND 2:1 METAL COMPLEXES OF
SUBSTITUTED THIOURACILS FROM POTENTIOMETRIC TITRATIONS
Ligand3 Metal 103[M+2]q °C Log Log K2b
2TU
Pb
+2
Cd
+2
2.00
25.0
4.68
3.07
1.80
25*2
4.69
3.69
1.60
25.2
4.70
3.5^
1.40
25.1
4.72
3.47
1.20
25.0
4.73
3.42
1.00
24.9
4.76
3.37
0.800
24.9
4.76
3.49
0.600
24.9
4.84
3.47
0.400
24.9
4.83
3.41
0.200
24.9
4.75
3.49
4.74
+
0.05
3.44
+
O.15
2.00
34.95
4.52
3.24
0.800
34.85
4.62
3.40
4767
+
0.05
T32
+
0.08
2.00
44.85
4.52
2.62
0.800
45.15
4.49
3.37
433
+
0.01
2.99
+
0.37
2.00
25.1
4.21
3.12
1.80
25.3
3.88
4.00
1.60
24.9
4.10
3-57
1.40
25.3
4.03
3.55
1.20
25.1
3.85
3.94
1.00
25.1
4.02
3.46
0.800
25.1
4.13
3.58
0.600
25.1
3.98
3.31
0.400
25.0
4.07
3.36
0.200
25.0
3.96
3.03
47Ü2
+
0.10
3.49
+
0.29
2.00
34.?
4.07
3.68
0.80
34.7
4.10
3.86
W7ÃœE
+
0.01
3.77
+
0.09
2.00
45.0
4.05
3.20
0.800
45.0
3.93
3.70
3.99
+
0.06
3745
+
0.25

145
TABLE II (continued)
Ligand
Metal
io3[m+2]0
°c
Log Kx
Log K2
Ni+2
60.00
25.0
2.46
3.78
6.00
25.0
2.59
3.57
4.00
25.0
2.38
1.69
1.60
25.0
2.60
2.12
1.00
25.0
2.40
2.84
2779 +
0.09
27ZE
+
0.71
PTU
Pb+2
2.00
25.8
4.79
3.34
1.60
25.8
4.76
1.00
25.8
4.84,
3.48
0.400
25.8
4*67d
3.26
0.200
25.8
4.50d
— — — —
7779 +
0.03
J7J5
+
0.09
2.00
34.80
4.58
3.50
0.800
34.80
4.69
3.32
4761 +
0
•
O
3.41
+
0.09
2.00
44.85
4.43
3.32
0.800
44.58
4.45
3.23
4.44 +
0.01
3.27
+
-3-
0
•
0
Cd+2
2.00
25.0
4.16
4.22
1.60
25.3
4.73
1.00
25.3
4.17d
4.81
0.400
25.3
4.31a
4.20
0.200
25.3
— — — —
— —
4.16 +
0.005
4.49
+
0.28
2.00
35.1
3.95
4.64
0.800
34.9
3.78
4.97
J7EE +
0.08
778T5
+
0.17
2.00
44.9
4.01
4.35
1.20
44.9
3.65
4.19
3.83 +
0.18
7727
+
0.08
Ni+2
2.00
25.9
1.34,
3.49d
1.60
25-9
0.30a
4.46d
1.00
25.9
1.36
3.53
rTH +
0.01
3.^1
4*
0.02
Zn+2
2.00
25.9
2.16
3.62

146
TAELE II (continued)
Ligand Metal 103[m+2]q °C Log ^ Log K2
6mtu
5MTU
Pb+2
2.00
26.0
4.63
3.29
1.60
2 5*5
4.73
3.33
1.00
25.5
4.68
3.45
0.400
25.0
4.73
3.00
4769
+
0.04
3T26
+
0.16
2.00
35.1
4.70
3.51
0.800
35.1
4.80
3.59
5775
+
0.05
3.55
+
0.04
2.00
44.9
4.54
3.28
0.800
44.95
4.52
3.48
4751
+
0.01
3T3I?
+
0.10
Cd+2
2.00
26.0
4.15
3.43
1.60
25.5
4.16
3.59
1.00
25.5
4.°9ri
3.70
q
0.400
26.0
3.67d
4.15
1
4.13
+
0.03
3.57
+
0.11
2.00
35.2
4.31
3.39
0.800
35.2
4.35
3.96
4.33
+
0.02
JT92
+
0
•
0
2.00
44.85
4.26
3.82
0.800
45.0
4.18
3.80
4722
+
0.04
T7HT
+
0.01
Ni+2
2.00
26.0
4.13
3.18
Pb+2
2.00
25.0
4.75
3.14
1.60
25.3
4.74
3.34
1.00
24.5
4.85
3.25
0.400
24.5
4.89
3.57
47BÜ
+
0.06
3.32
+
0.16
2.00
35.05
4.65
3.25
0.800
34.9
4.73
3.12
4769
+
0.04
3.18
+
0.06
2.00
44.85
4.51
3.30
0.800
44.85
4.59
3.13
4.55
+
0.04
3.21
+
0.08

147
TABLE II (continued)
Ligand
Metal
io3[m+2]0
°c
Log K1
Log K2
Cd+2
2.00
25.0
4.28
3.47
1.60
25.0
4.27
3.75
1.00
25.3
4.22,
3.92,
0.400
25.0
3.79a
4.48a
4.25 +
0.03
T77T +
0.18
2.00
34.85
4.14
3.82
0.800
34.9
4.25
3.93
4719 +
0
•
0
1787 +
0.07
2.00
44.75
3.97
3.23
0.800
44.95
4.11
3.85
77Ü7 +
0.07
337 +
0.31
Ni+2
2.00
26.0
2.30d
2.84
1.60
25.3
2.72
2 -74d
1.00
25.0
2.61
1.63d
2787 +
0
•
0
2.79 +
O
•
0
5,6DMTU
Pt>+2
2.00
25.8
5.01
4.04
1.60
25.0
4.98
3.86
1.00
25.3
5.°3d
3.83
0.400
25-3
4.79a
3.92
57ÃœI +
0.02
379T +
0.08
2.00
34.85
5.04
4.37
0.800
34.85
4.99
4.72
57ÃœT +
0.02
4.54 +
0-
T—1
•
O
2.00
45.35
4.81
4.23
0.800
44.9
4.82
4.13
47HT +
0.00
47TH +
O
•
O
'ax
Cd+2
2.00
25-5
4.52c
1.60
25.5
4.36
3.99
1.00
25*3
4.37
4.24
0.400
25.5
4.37
4.11
4717 ±
-3"
O
O
.
O
47TT +
0.10
2.00
34.85
4.50
4.23
0.800
34.95
4.40
4.40

148
TABLE II (continued)
Ligand Metal 103[m+2],
Log K^_
Log K2
2.00
44.85
4.28
4.21
0.800
45.0
4.19
4.45
4723
+
0.04
+
0.12
Ni+2
1.00
25-5
3.03
2.99
5 CETU
2.00
25*5
ppt.
ppt.
0.800
25.5
ppt.
ppt.
2.00
35.0
ppt.
ppt.
0.800
35.0
ppt.
ppt.
2.00
45.0
ppt.
ppt.
0.800
45.0
ppt.
ppt.
Cd+2
2.00
24.8
3.57
3.74
0.800
25.0
3.48
3.56
3.52
+
0.04
3.65
+
0.09
2.00
35.4
3.50
3.72
0.800
34.5
—
—
2.00
44.85
3.79
3.98
0.800
44.85
3.55
3.94
3.67
0.12
3T96
+
0.02
a
Ligand abbreviations: (2TU) 2-Thiouracil, (PTU) 6-n-
-Propyl-2-Thiouracil, (6MTU) 6-Methyl-2-Thiouracil, (5MTU)
5-Methyl-2-Thiouracil, (5.6DMTU) 5.6-Dimethyl-2-Thiouracil,
(5CETU) 5-Carboethoxy-2-Thiouracil.
>>
Log and log Kg values were derived from the slope and
intercept values of plots of ñ/(l-ñ)[U-] versus =py £u-]
in accordance with = K1 + —- [U ]K1K„.
(1-H)[U ] 1 n-1 L 1 2
These are the mean values of the stability constants.
The error is expressed as the standard deviation.
^This value is not included in the calculation of the mean
value because of large difference from other values.
eLog K^Kg values are given for: PTU with 0.200 x 10”^ m
Cd+2 at 25.3°, 8.66; 5CETU with 0.800 x 10"3M Cd+2 at
34.5°, 7.52.

TABLE III—COMPOSITION OF AQUEOUS SOLUTIONS AND ULTRAVIOLET ABSORBANCE OF AQUEOUS
THIOURACIL-METAL ION MIXTURES AT 25°
FbiorBaHce~o?~
Liganda [Ligand] Metal [Metal] Absorption X Metal Ion at
*max of Li^and
2TU
1.00 x 10
-4
Cu
+2
Cd
+2
Pb
+2
0 4
1.328
272
1.00
X
1.120
272
7.00
X
10"4
0.987
272
0
1.614
212
0 -4
1.376
271
1.00
X
10 4
1.325
271
7.00
X
10 4
1.290
271
0
1.669
212
o ,,
1.370
273
1.00
X
)- -3-
i i
o o
rH rH
1.320
273
7.00
X
1.250
273
0
1.673
212
0 -4
1.310
272
1.00
X
10-4
1.530
272
7.00
X
10 4
>2
272
0 4
1.641
212
1.00
X
10 4
>2
212
7.00
X
10-4
>2
212
0.200
VO

TABLE III (continued)
Absorbance of
Ligand [Ligand] Metal [Metal] Absorption X Metal Ion at
W of Li«and
PTU
Pe
+ 2
1.00 x 10
-4
Cu
+2
Cd
+2
Pb
+2
0 4
1.210
274
1.00
X
io:j
1.225b
274
7.00
X
10 4
1.295"
274
0 -4
1.561
213
1.00
X
10 4
1.553
213
7.00
X
10 4
1.550
213
0 4
1.520
273
1.00
X
10 4
1.060
273
7.00
X
10
1.056
271
0
1.585
213
0 -4
1.540
273
1.00
X
10 4
I.528
273
7.00
X
10 4
1.390
273
0
1.600
213
0 4
1.490
273
1.00
X
10 4
1.462
273
7.00
X
10"4
1.423
273
0
1.568
213
0 4
1.548
274
1.00
X
10“J
1.760
274
7.00
X
10'
>2
274
0 Is
1.602
214
1.00
X
10’4
>2
214
7.00
X
10~4
>2
214
0.200
o

TABLE III (continued)
Absorbance of
Ligand [Ligand] Metal [Metal] Absorption X Metal Ion at
W of Li£and
ÓMTU
Fe
+2
1.00 x 10-4 Cu+2
Gd
+2
Pb
+2
0 4
1.395
274
1.00
X
i°i
1.408
274
7.00
X
io-4
1.440
274
0 4
1.520
213
1.00
X
10 4
1.535
213
7.00
X
10“4
1.565
213
0 4
1.512
275
1.00
X
io t
1.000
275
7.00
X
10"4
0.945
275
0
1.585
213
0 4
1.471
275
1.00
X
10 4
1.460
275
7.00
X
10"4
1.386
275
0
1.565
213
0 4
1.505
274
1.00
X
10 4
1.480
274
7.00
X
10-4
1.440
274
0
1.582
213
0 -4
1.550
275
1.00
X
10-4
1.765
275
7.00
X
10 4
>2
275
0 h
1.621
214
1.00
X
io"£
>2
214
7.00
X
10~4
>2
214
0.221

TABLE III (continued)
Absorbance of
Ligand [Ligand] Metal [Metal] Absorption \ Metal Ion at
of Ligand
max
5MTU
1.00 x 10
-4
Pe
+2
Cu
+2
Cd
+2
Pb
+2
0 4
1.566
275
1.00 x
10 4
1.611
275
0.026
7.00 x
10~4
1.571
275
0.182
0 4
1.615
214
1.00 x
10 I
1.672
214
0.050
7.00 x
10 4
1.633
214
0.350
0 4
1.535
275
1.00 x
10 4
1.534
275
7.00 x
10 4
1.528
2 75
0
1.575
213
0 4
1.534
275
1.00 x
10 4
I.O56
275
7.00 x
10 4
1.071
272
0
1.610
213
0 -4
1.545
275
1.00 x
1° t
1.506
275
7.00 x
10 4
1.422
275
0
1.610
213
0
1.523
275
1.00 x
1° 4
1.475
275
7.00 x
10 4
1.438
275
0
1.600
213
ro

TABLE III (continued)
Ligand [Ligand]
5,6DMTU
1.00 x 10~4
Absorbance of
Metal [Metal] Absorption X Metal Ion at
"max of Li^d
Fe+3
0 A
1.525
275
1.00
X
10_4
1.750
275
7.00
X
10 4
>2
275
0 k
1.600
214
1.00
X
10"a
>2
214
7.00
X
10'4
>2
214
Fe+2
0 -4
1.530
275
1.00
X
10 4
1.522
275
7.00
X
10 4
1.538
275
0 4
1.620
214
1.00
X
10l4
10 ^
1.620
214
7.00
X
1.650
214
Gu+2
0 -4
10 4
1.763
276
1.00
X
1.470
276
7.00
X
10 4
1.325
276
0
1.439
216
Cd+2
0 -4
1.752
276
1.00
X
10 4
1.722
276
7.00
X
10 4
1.717
276
0
1.422
216
Pb+2
0 4
1.763
2 77
1.00
X
10 4
1.736
277
7.00
X
10 4
1.704
277
0
1.433
216
VjJ

TABLE III (continued)
Absorbance of
Ligand j_Ligand] Metal [Metal] Absorption \ Metal Ion at
Xmax 0f Li^d
2EM4HP
Fe
+2
1.00 x 10
-4
Cu
+2
Cd
+2
0 4
1.665
276.
1.00
X
10 4
1.873
276
0.210
7.00
X
10“4
>2
276
0 4
1.395
216
1.00
X
1CU
10 4
>2
216
7.00
X
>2
216
0 4
1.675
276
1.00
X
10 4
1.698
276
0.020
7.00
X
10~4
1.715
276
0.140
0 4
1.395
216
0
1.00
X
10 4
1.405
216
0.035
7.00
X
10“4
1.412
216
0.280
0 4
0.550
284
1.00
X
10 4
1.561
284
7.00
X
10~4
1.570
284
0 4
0.555
283
1.00
X
10 4
O.56O
283
7.00
X
10 4
0.573
283
0 -4
1.180
230
1.00
X
10-4
1.350
230
0.158
7.00
X
10 4
>2
230
0
0.860
203
4^

TABLE III (continued)
ÁbsoFbañce-of-
Ligand [Ligand] Metal [Metal] Absorption \ Metal Ion at
of Ligand
max
Pb
+2
Fe
+2
0 4
0.550
283
1.00
X
10 4
0.547
283
7.00
X
10~4
0.544
283
0 -4
1.175
230
1.00
X
10 4
1.400
230
7.00
X
10~4
>2
230
0 4
0.860
204
1.00
X
10"k
>2
204
7.00
X
10“4
>2
204
0 -4
0.549
283
1.00
X
10 4
0.755
283
7.00
X
10 4
1.830
283
0 -4
1.185
230
1.00
X
10 4
>2
230
7.00
X
10"4
>2
230
0
0.855
203
0 -4
0.555
283
1.00
X
10 4
0.555
283
7.00
X
10 4
O.565
283
0 -4
1.200
231
1.00
X
10 4
1.210
231
7.00
X
10-4
1.250
231
0 4
0.872
204
1.00
X
10 4
0.872
204
7.00
X
10"4
0.915
204
0.222
0.230
1.225
v^n

TABLE III (continued)
Absorbance of
Ligand [Ligand] Metal [Metal] Absorption X Metal Ion at
'max 0f Li®and
NNDE6MTU
1.00 x 10
-4
Gu
+ 2
Cd
+ 2
Pb'
+2
0 4
1.320
278
1.00
X
10 4
1.321
278
7.00
X
10"4
1.332
278
0
1.525
213
0 4
1.315
278
1.00
X
10 4
1.316
278
7.00
X
10"4
1.336
278
0
1.540
222
0 4
1.312
277
1.00
X
10 4
1.286
277
7.00
X
10"4
1.279
277
0
1.540
222
aLigand abbreviations? (2TU) 2-Thiouracil, (PTU) 6-n-Propyl-2-Thiouracil, (6MTU)
6-Methyl-2-Thiouracil, (5MTU) 5-Methyl-2-Thiouracil, (5.6DMTU) 5.6-Dimethyl-2-
-Thiouracil, (2EM4HP) 2-Ethylmercapto-4-Hydroxypyrimidine, (NNDE6MTU)
N#N -Diethyl-6-Methyl-2-Thiouracil.
^Sample sat for 15 minutes exposed to air. Oxidation of ferrous to ferric ion caused
a high absorbance value.
C «.¿i
Calculated from absorbance of 1 x 10 M metal ion.
^Solvent was 0.001 M HC10^.

TABLE IV—INTRINSIC SOLUBILITIES3 OF LIGANDS AND pK^ VALUES DETERMINED FROM SOLUBILITY
DATA AT 25°
Ligand Solubility (M/L) pKQ
2-Thiouracil
5-53
X
1
o
*—1
7-5 2
6-n-Propyl-2-Thiouracil
7.07
X
cn
1
o
1—1
7.80
6-Methyl-2-Thiouracil
3.75
X
1-»
o
1
7.9^
5-Methyl-2-Thiouracil
3-58
X
10"3
7.80
5,6-Dimethyl-2-Thiouracil
8.79
X
i
o
*—1
—
5-Carboethoxy-2-Thiouracil
7-97
X
10"3
—
6-Amino-2-Thiouracil
1.79
X
i
o
T—Í
—
determined from the absorbance of 0.1 M HCIC^ saturated ligand solutions.
b Ab-Ab +
Calculated from the intercept values of plots of log :— versus pH according to
Ab-Ab.T+ , Ab„+
log iE_ = pH - pK .
AV
*v3

TABLE V—EFFECTS OF CONCENTRATION OF CUPRIC ION AND PER¬
CHLORIC ACID ON DIFFUSION CURRENT, HALF-WAVE POTENTIAL,
E1/2 AND E^/4 - E1/¿+ VALUE FOR POLAROGRAPHIC REDUCTION
OF CUPRIC NITRATE9-
[Cu+23 [HC104] IjjHa” E°/2 E3/4-e1/4 °c
0.001
0.1
8.85
9.15
0.01
9.05
9.09
0.001
9.12
0.0001
9.12
8.89
8.77
8.86
0.0008
0.001
7.32
7-34
0.0006
0.01
5.49
5.44
5.^3
0.0005
0.0001
4.47
4.44
4.41
0.0004
0.001
3.57
3.56
3.51
0.0003
0.01
2.74
2.76
2.74
2.76
2.73
2.74
2.68
2.73
0.0413
-O.O333
21.0
0.0383
-O.O322
21.3
0.0413
-O.O302
21.8
0.0413
-O.O312
22.0
0.0413
-O.O323
22.0
0.0393
-O.O327
22.1
0.0410
-0.0273
21.0
0.0430
-0.0283
21.0
0.0420
-O.O262
21.0
0.0402
-O.O341
21.3
0.0402
-O.O33I
21.3
0.0402
-O.O33O
21.0
0.0412
-O.O29I
21.1
0.0428
-O.O27O
21.2
-O.O32I
20.8
—
-O.O323
20.8
—
-O.O332
20.8
__
-O.O3OI
20.2
—
-O.O3II
20.3
--
-O.O32I
20.3
0.0412
-O.O3II
21.5
0.0422
-O.O32O
21.5
0.0442
-O.O32O
21.5
0.0422
-O.O32O
21.5
0.0442
-O.O32O
21.5
0.0432
-O.O3IO
21.5
0.0422
-O.O3II
21.5
0.0422
-O.O330
21.5

159
TABLE V (continued)
[Cu+2] [HC104] iD)tób E¡/2 e3A-e1/4 °c
0.0002 0.1
1.82
0.0412
-0.0360
21.3
1.85
0.0413
-O.35I
21.3
1.87
0.0393
-0.311
21.3
1.82
0.0382
-O.O32I
21.3
1.82
0.0382
-O.O35I
21.4
0.0001 0.2648
0.943
0.0200
-0.0780
20.7
0.976
0.0220
-O.O76O
20.8
O.972
0.0292
-0.0594
21.0
0.948
0.0292
-0.0454
21.0
0.976
0.0322
-O.O374
21.1
0.976
0.0342
-0.0394
21.0
9,
All solutions were
0
.2 M in
NaClO^ and
contained
0.001%
Triton.
bThe ip values were
measured
at El/2.
Q
Average values of
E./2, not
including
data for 0.
0001 M
cupric ion is 0.0411
vs. 5.
C.E.
^These values were
obtained
from a single solution
left
in the cell and ran six times. The shape of the curve
changed and became more reversible.

160
TABLE VI—EFFECTS OF 6-n-PROPYL-2-THIOURACIL AND PERCHLORIC
ACID CONCENTRATIONS ON E3/4 - E1/4 AND (E¿)c ~b(Ei)J VALUES
OF 2.00 x 10"4 MOLAR CUPRIC ION AT 21°
103[PTU]C E3/4-Ei/4,Volts [HC104] (Ei)C-(E¿)S + a, Volts
4.375
-0.0582
0.0500
-0.1324 + 0.0008
3.875
-0.0573
-O.I35O + 0.0018
3.375
-0.0593
-0.1324 + 0.0041
2.875
-0.0572
-0.1266 + 0.0031
2.375
-0.0612
-0.1255 + 0.0030
1.875
-0.0592
-O.II99 + 0.0026
4.375
-0.0582
0.0800
-0.1266 + 0.0013
3.875
-O.0582
-O.I319 + 0.0009
3.375
-0.0533
-O.I305 + 0.0005
2.875
-0.0572
-O.I258 + 0.0019
2.375
-0.0572
-0.1193 + 0.0022
1.875
-0.0553
-O.II5I + 0.0004
1.375
-O.057I
-0.1119 + 0.0025
4.375
-0.0552
0.100
-0.1248 + 0.0012
3.875
-0.0612
-0.1288 + 0.0027
3.375
-0.0543
-0.1223 + 0.0020
2.8?5
-0.0612
-O.IO67 + 0.0137
2.375
-0.0572
-0.1135 + 0.0085
1.875
-0.0601
-0.0789 + 0.0001
1.375
-0.0602
-0.1045 + 0.0024
4.375
-O.O572
0.200
-0.1206 + 0.0035
3.875
-O.O562
-0.1266 + 0.0028
3.375
-O.O582
-O.I2O9 + 0.0040
2.875
-O.O572
-0.1185 + 0.0043
2.375
-O.O613
-O.II98 + 0.0030
1.875
-O.O578
-0.1112 + 0.0036
1.375
-O.O563
-0.1066 + 0.0014
aThe difference in half-wave potentials of the complexed
and simple cupric ions.
^lonic strength is constant at 0.20 M using sodium per¬
chlorate. All solutions contained 0.001^ Triton as
maximum suppressor. The diffusion current was constant
at 1.48pA.
C6-n-Propyl-2-Thiouracil.

TABLE VII—EFFECT OF 5.6-DIMETHYL-2-THIOURACIL ON RAT THYROID WEIGHT AFTER TWO WEEKS
OF DRUG DIET
Apparent
% 5.6DMTU Estimated Dose3 Thyroid Weight Relative,
W Diet mg./100gm. Body mg./lOOgm. Body Weight Activity0
Weight/Day (2TU=1.0)
Relative
Intrinsic0
Activity
(2TU=1.0)
0.00
—
8.11
—
—
—
6.00
—
—
—
6.91
—
—
—
6.05
—
— — — —
— — — —
5*71 d
6ÍJ5 + 0.87d
— — —
0.025
0.965
16.03
2.78
0.965
15.30
2.57
0.965
12.71
1.80
0.965
14.51
2.33
0.965
14.29
14.57 + 1.11
2.27
2.35 + 0.33
161

TABLE VII (continued)
Apparent Relative
% 5,6DMTU Estimated Dose Thyroid Weight Relative Intrinsic
W Diet mg./lOOgm. Body mg./lOOgm. Body Weight Activity Activity
Weight/Day (2TU=1.0) (2TU=1.0)
0.050
0.100
1.930
13.32
0.?4
1.930
13.55
O.76
1.930
13.15
0.72
1.930
11.70
O.56
1.930
14.47
0.86
13.24 + 0.89
0.73
3.86
12.98
0.45
3.86
14.64
0.57
3.86
10.86
0.30
3.86
10.16
0.25
3.86
10.47
0.27
ii.82 +1.72
0.37
0.09
0.12
0.37
aSince diet consumption was not measured the dose was estimated from the diet con¬
sumption found for the experiment using 5-^ethyl-2-thiouracil. The estimated diet
consumption is 3*86 gm./lOOgm. body weight per day.

TABLE VII (continued)
â– j_
°The apparent relative activity was calculated by
Thyroid Weight of a Drug Treated Rat Mean of Thyroid Weights of Controls
Thyroid Weight of a 2TU Treated Rat Thyroid Weight of 2TU Controls
where the denominator, calculated at the same dose as the numerator, was obtained
from a standard log dose-response curve for 2-thiouracil from the literature
(Endocrinology, 3?t 456, (19^5))* The literature values for the 2-thiouracil
doses (mg./100gm. Body Weight/da.y) and associated thyroid weights (mg./lOOgm.
Body Weight) used for the standard curve are: 0.04, 7.6; 0.13, 7*6; 0.4, 8.6;
1.2, 11.8; 3*6, 21.4; 10.0, 23»4, respectively.
cThe relative intrinsic activity was calculated by
Maximal Thyroid Weight of a Drug Treated Rat — Mean of Thyroid Weights of Controls
Maximal Thyroid Weight of a 2TU Treated Rat — Thyroid Weight of 2TU Controls
where the denominator has a value of 16.0.
^The error is expressed as + a.

TABLE VIII—EFFECT OF 6-METHYL-N.N'-DIETHYL-2-THIOURACIL ON THE RAT THYROID WEIGHT AFTER
TWO WEEKS OF DRUG DIET
% 6MNN DETU Diet Consumption Dose
in Diet gm./lOOgm. Body mg./100gm.
Weight/Day Body Weight
Thyroid
Weight
mg./100gm.
Body Weight
Apparenta
Relative
Activity
(2TU=1. 0)
Relative ,
Intrinsic0
Activity
(2TU=1.0)
7.44
—
4.44
—
8.04
—
5-83
—
5.28
—
4.44
—
7.58
—
4.41
—
7.38
— — — —
3.22
5T47 + 0.8
6.59
6.59
6.27
0.1
0.11
9.79
8.79
5.26
0.05
0.05
6.54
6.54
9.19
F79Ü + 1.66
0
vH
•
0
+ 1
O
oa th
• •
o|o
0.29
O.15 + 0.10
aThe apparent relative activity was calculated by
Thyroid Weight of a Drug Treated Rat Mean of Thyroid Weights of Controls
Thyroid Weight of a 2TU Treated Rat Thyroid Weight of 2TU Controls
where the denominator, calculated at the same dose as the numerator, was obtained
from a standard log dose-response curve for 2-thiouracil from the literature

TABLE VIII (continued)
(Endocrinology, 37» 456, (19^+5 )) • The literature values for the 2-thiouracil doses
(mg./100gm. Body Weight/day) and associated thyroid weights (mg./lOOgm. Body Weight)
used for the standard curve are: 0.04, 7.6; 0.13, 7*6; 0.40, 8.6; 1.20, 11.8; 3*60,
21.4; 10.0, 23.4, respectively.
The relative intrinsic activity was calculated by
Maximal Thyroid Weight of a Drug Treated Rat — Mean of Thyroid Weights of Controls
Maximal Thyroid Weight of a 2TU Treated Rat — Thyroid Weight of 2TU Controls
where the denominator has a value of 16.0.

TABLE IX—EFFECT OF 6-n-PROPYL-2-THIOURACIL AND 5-METHYL-2-THIOURACIL ON THYROID WEIGHT
AFTER TWO WEEKS OF DRUG DIET
% PTU Thyroid Weighta
in Diet (jng./100gm.
Body Weight)
Id
Dose Apparent
(mg./lOOgm. Relative
Body Weight/Day) Antithyroid
Activity
(2TU=1.0)
Diet Consumption Relative0
(gm./lOOgm. Body Intrinsic
Weight/Day) Activity
(2TU=1.0)
0.00
0.020
0.040
7.45
6.95
5.26
5.35
3.61
3772 +
1.36
28.55
1.31
25.68
1.36
23.08
1.31
22.71
1.18
25.00 +
2.34
1.29
29.21
2.24
24.21
2.72
33.15
2.63
31.09
2.66
22.11
2.35
27.95 + 4.16 2732
8.79
9.09
8.32
9.06
9.12
8787
+
0,
.30
4.56
6.58
3.76
6.84
3.47
6.59
3.86
5.90
3.91 +
0.40
£747
+
0,
.34
2.17
5.62
1.50
6.81
2.28
6.59
2.09
6.65
1.33
5.89
1.87 +
0.38
6.31
+
0,
.46
1.20
1.39
166

TAELE IX (continued)
% PTU Thyroid Weight
in Diet (mg./100gm.
Body Weight)
Dose Apparent Diet Consumption Relative
(mg./lOOgm. Relative (gm./100gm. Intrinsic
Body ’Weight/Day) Antithyroid Body Weight/Day) Activity
Activity (2TU=1.0)
(2TU=1. 0)
0.080 21.76
25.00
29.47
25.86
23752 +2.74
5.32
1.06
4.28
1.32
4.06
1.66
4.77
1.36
4.61
1773 + 0.21
6.66
5.35
5.08
5.97
3775 + 0.61 1.24
% 5MTU
in Diet
0.00
0.0250
5.31
5.51
3.29
5.34
4.04
5.00
4.92
1.64
4.68
5.34
05 + O.71
4.57 +
10.31
_d
__d
_d
9.27
O.54
2.67
2.16
11.24
1.43
1.21
5.73
13.76
0.76
3.58
3.05
11.00
1.08
0.81
4.32
TT7TT + 1.5
7793
0)5 + 1.11
3752 +
1.5
1.3
0.41
O'
-o

TABLE IX (continued)
% 5MTU
Thyroid Weight
Dose
Apparent
Diet Consumption
Relative
in Diet
mg./100gm.
(mg./100gm.
Relative
(gm./lOOgm.
Intrinsic
Body Weight)
Body Weight/Day)
Antithyroid
Body Weight/Day)
Activity
Activity
(2TU=1.0)
(2TU=1.0)
0.0500
17.65
2.71
1.09-
5-42
14.63
e
6
e
11.48
2.06
0.58
4.13
11.57
1.17
1.65
2.34
13. «3 + 2.5
1798"
1.10 + 0.44
3797 + 1.3
0.59
0.100
13.07
2.58
0.71
2.58
15.37
5.14
0.72
5.14
10.86
4.09
0.43
4.09
11.87
3.46
0.54
3.46
11.57
3.83
0.50
3.83
12.55 + 1.6
3782
T57I8 + 0.11
3782 + 0.83
0.50
aThe error is expressed as the standard deviation of the mean.
The apparent relative activity was calculated by
Thyroid Weight of a Drug Treated Rat Kean of Thyroid Weights of Controls
Thyroid Weight of a 2TU Treated Rat Thyroid Weight of 2TU Controls
where the denominator, calculated at the same dose as the numerator was obtained
from a standard log dose-response curve for 2-thiouracil from the literature
(Endocrinology, 37, 456, (1945)). The literature values for the 2-thiouracil

TABLE IX (continued)
doses (mg./100gm. Body Weight/Day) and associated thyroid weights (mg./100gm.
Body Weight) used for the standard curve are: 0.04, 7*6; 0.13, 7*6; 0.4, 8.6;
1.2, 11.8; 3*6, 21.4; 10.0, 23.4, respectively.
cThe relative intrinsic activity was calculated by
Maximal Thyroid Weight of a Drug Treated Rat — Mean of Thyroid Weights of Controls
Maximal Thyroid Weight of a 2TU Treated Rat — Thyroid Weight of 2TU Controls
where the denominator has a value of 16.0,
^This rat died on the fourteenth day of experiment.
eThis rat died on the fifteenth day of experiment.

170
TABLE X—NEGATIVE LOGARITHMS OF DISSOCIATION CONSTANTS
(pK* )a OF LIGANDS AS A FUNCTION OF TEMPERATURE13
SL
pKa + o
Ligand
2TU
25
7.46
+
O
•
o
35
7.22
45
7.09
PTU
25
7.76
+
0.05
35
7.48
45
7.17
6mtu
25
7.73
+
0.08
35
7.6 5
45
7.41
5MTU
25
7.71
+
o
•
o
35
7.57
45
7.35
5,6DMTU
25
8.08
+
0.06
35
8.06
45
7-76
5CETU
25
6.43
+
0.02
35
6.40
45
6.27
2EM4HP
25
7.01
+
0.05
Q t
pK is equal to the pH at half-neutralization according
Q.
to pK& = pH - log .
b
c
LHlJ] 2
The ionic strength, was 0.006 M.
Ligand abbreviations: (2TU) 2-Thiouracil, (PTU) 6-n-
-Propyl-2-Thiouracil, (¿MTU) 6-Methyl-2-Thiouracil,
(5.6DMTU) 5» 6-Dimethyl-2-Thiouracil, (5MTU) 5-Methyl-2-
-Thiouracil, (5CETU) 5-Carboethoxy-2-Thiouracil,
(2EM4HP) 2-Ethylmercapto-4-Hydroxypyrimidine.

TABLE XI—ENTROPY AND ENTHALPYa VALUES FOR THE IONIZATION OF THIOURACILS (HU)
AND FOR THE FORMATION OF MU+ BY COMPLEXATION WITH CADMIUM AND LEAD
Cadmium Lead Ionization
Ligand AS e.u.c AH Kcal.b e.u.c Kcal.b e.u.c Kcal.b
2TU
14.0
-1.4
6.0
-4.6
-4
9.0
PTU
—
—
-4.5
-7.9
8.6
13.2
6MTU
20.0
0.2
10.0
-3.5
-13.4
6.6
5MTU
4.2
-4.5
6.4
-4.6
-10.5
7.4
5,6DMTU
7.5
-3.8
8.0
-4.5
-13.5
7.1
5CETU
27.5
2.3
— — — —
-17.6
3.5
aEnthalpy values are for 37.5°«
bError in enthalpy values is + 0.1 Kcal.
£
Error in entropy values is approximately + 10 e.u.

172
TABLE XII—COMPOSITIONS AND ABSORBANCES OF 2-THIOURACIL
-CUPRIC ION SOLUTIONS FOR JOB'S PLOTS
lo\cu+^] 104 [2TU] Absorbance, Absorbance,
272 mp 345 MM
1.00
0.000
1.389
0.000
0.900
0.100
1.179
0.016
0.800
0.200
1.010
0.025
0.700
0.300
0.839
0.036
0.600
0.400
0.714
0.030
0.500
0.500
0.589
O.O36
0.400
0.600
0.450
0.031
0.300
0.700
0.312
0.025
0.200
0.800
0.201
0.022
0.100
0.900
0.085
0.015
0.000
1.000
0.000
0.000

TABLE XIII—NEGATIVE LOGARITHM OF THE RATIO OF THE DIS¬
SOCIATION CONSTANT51 OF THE OXIDIZED COMPLEX TO THE DIS¬
SOCIATION CONSTANT OF THE REDUCED COMPLEX OF CUPRIC-THI-
OURACILS (Kq/K ) AS A FUNCTION OF TEMPERATURE AND ACID
CONCENTRATION
2-Thiouracil
-Log K0/Kr
[hcio4]
21°
0
0
35°
4¿
0.2648
10.38
—
—
0.200
10.66
9.99
9.67
6.75b
0.100
10.12
9.49
9.30
—
0.0800
10.11
9.56
9.27
—
0.0500
10.22
10.29 +
9.38
0.20 J75Ó + 0.
9.29
23 9T3S + 0.16
[HC104]
6-n-Propyl-
-Log
2-Thiouracil
Ko/Kr
21°
35°
40°
0.200
11.30
10.74
10.51
0.100
11.61
10.37
10.13
0.0800
11.63
10.35
7.56°
0.0500
11.41
TT74Ü +
10.36
0.14 10.45 + C
7.49°
1.I6
5,6-Dimethyl
-Log
-2-Thiouracil
Ko/Kr
LHC104]
21°
0.200
11.61
0.100
11.27
0.0800
11.33
0.0500
11.12
11.33 + 0.18
35
40'
7.98c
10.84

174
TABLE XIII (continued)
6-Methyl-2-Thiouracil
-Log Ko/Kr
1—1
K
O
• M
O
1 l
21°
35°
0.200
10.82
—
0.100
10.83
—
0.0800
11.02
—
0.0500
10.72
7.87d
1Ó.85 + 0.11
5-Methyl-2-Thiouracil
-Log Ko/Kr
lhcio4j
21°
35°
0.200
10.90
—
0.100
10.81
—
0.0800
10.70
—
0.0500
10.70
7.64
10.77 +
0.08
5-Carhoethoxy-2-Thiouracil
-Log K0/Kr
[HC1C4]
21°
35° 40°
0.200
6.12d
—
—
0.100
—
—
—
0.0800
—
—
—
0.0500
7.02e
5 - 61e
5.^5e
aThe ratios are calculated according to
(Vc " (Vs - nF
In
K
/ _ . \P-Q
|_H+] X±\ n
K
- —r1 Y- ln LHU].
a
The error is expressed as +

175
TABLE XIII (continued.)
bThe polarographic curves may have contained two waves very
close together but due to the small difference was treated
as a single curve. The value of log KQ/Kr in the table is
for correlation purposes only and is not accurate. The
plot of ^Ei versus log[2TU] had a zero slope.
Q
The plots had a zero slope. The polarographic curves
2
had E^/^-E^y^ values of -0.045 to -0.068 volts. The tabu¬
lated values of log K /K in the table are valid.
or
^The polarographic curve is not constant and becomes
steeper and the E| value more negative. The E^/^-E^y^
value is about -0.035 volts. The E¿ value is indeoendent
2
of the ligand concentration. The value of K /K in the
or
table is not accurate and is included only for correlation
purposes.
eThe polarographic curve is very irreversible and gave
E^/^-E-jy^ values of about -0.09 volts. The curves were
stable and gave zero slopes in the /^E^ plots. The log
KQ/Kr values in the table are very approximate and are
given for correlation purposes.

B. Figures

Fig. 1—Potentiometric titration curves of filtered, day
old cupric nitrate-2-thiouracil mixtures with )i = 0.006
at 25*0°. Twenty ml. aliquots were 2.00 x 10“3 M in
2-thiouracil and (A), 2.00 x 10~3 M; (B), 1.60 x 10”3 M;
(C), 1.40 x 10~3 Ms (D), 1.00 x 10"3 M; (E), 6.00 x 10“4M
(F), 2.00 x 10“4 M; (G), zero M in cupric nitrate. Curve
—4
H is the titration of 25 »0 mis. of 4.00 x 10 M cupric
nitrate. The titer of alkali between inflections 2 and
3 are: (D), 1.22 x 10"2; (E), 2.50 x 10~2; (F),
3.52 x 10“2 meq.

HO°N ’baw
178
pH

Fig. 2—Ultraviolet spectra of mixtures of cupric nitrate
and 6-methyl-2-thiouracil in water at 25*0°. Key: (1),
1.00 x 10’4 M Cu+2 (no absorbance); (2), 1.00 x 10~4 M
ó-methyl-2-thiouracil (6MTU); (3). 1.00 x 10"4 M Cu(N0^)2
and 1.00 x 10”4 M 6MTU; (4), 7.00 x 10“4 M Cu(N0^)o and
_k ~ — J c
1.00 x 10 M 6MTU.

Absorbonce
180

Fig. 3—Ultraviolet spectra of mixtures of lead nitrate
and 2-thiouracil (2TU) in water at 25«O0» Key: (1),
1.00 x 10~4 M Pb+2; (2), 1.00 x 10“4 M 2TU; (3),
1.00 x 10~4 M 2TU and 1.00 x 10“4 M Pb+2; (4), 1.00 xl0“4
M 2TU and 7.00 x 10~4 M Pb+2.

182
Wavelength, m/¿

Fig. 4—Ultraviolet spectra of mixtures of ferric nitrate
and. 6-methyl-2-thiouracil (6MTU) in water at 25»0°« Key:
(1), 1.00 x 10"4 M Fe+3; (2), 1.00 x 10~4 M 6MTU; (3).
1.00 x 10“4 M 6MTU and 1.00 x 10“^ M Fe+^. The points
are the summation of the absorbances of (1) and (2).

Absorbance
184

Fig. 5—Ultraviolet spectra of mixtures of cupric nitrate
and N,N -diethyl-6-methyl-2-thiouracil (NNDEbMTU) in water
at 25.0°. Key: (1), 1.00 x 10“4 M Cu+2; (2), 1.00 x 10~4
M NNDE6MTU; (3), 1.00 x 10“4 M Cu+2 and 1.00 x 10"4 M
NNDEÓMTU; (4), 7.00 x 10”4 M Cu+2 and 1.00 x 10"4 M
NNDEÓMTU.

Absorbance
186
Wavelength, m/¿

Fig. 6—Ultraviolet spectra of mixtures of cupric nitrate
and 2-ethylmercapto-4-hydroxypyrimidine (2EM4HP) in water
at 25.0°. Key: (1), 1.00 x 10"4 M Cu(N03)2; (2),
1.00 x 10~4 M 2EM4HP; (3). 1.00 x 10“4 M Cu(N03)2 and
1.00 x 10“4 M 2EM4HP; (4), 7.00 x 10“4 M Cu(N03)2 and
1.00 x 10"4 M 2EM4HP.

Absorbance
188

Fig. 7—Plot of absorbance at 272 m)i versus the mole fraction of
cupric nitrate of aqueous solutions of cupric nitrate and
2-thiouracil at 25*0°. Maximum concentrations of metal ion and
_k
ligand are 1.00 x 10 M. The upper line assumes no interaction.

190
© Q
eouoqjosqv
Mole Fraction of Cu

Fig. 8—Job's continuous variations plots of absorbance of aqueous
mixtures of cupric nitrate and 2-thiouracil at; (A), 3^5 nya; (B),
272 mja.

Absorbance
Mole Fraction of Cu*2
\D
ro
Absorbance

Fig. 9—Plots of ultraviolet absorbance at 272 and 3^5 mji
[Cu+^]/[2TU] where [2TUQ is constant at 5«00 x 10”^ M.
versus

Absorbance
Ratio of Cone, of Gu*2 to 2TU
VO

Fig. 10—Ultraviolet spectra of aqueous 2-thiouracil
solutions as a function of pH at 25*0°. C2TU] =
1.00 x 10"4 M.

Absorbance
Wavelength, mjn

Fig. 11—Plots of absorbance of aqueous 2-thiouracil
solutions versus pH at 230 mji (01), 240 m)i (â–¡), 270
mp (O)* 310 mp (Ql), 320 mp (X) and 285 mp (#).
The pK values (pH at half-neutralization) are
d.
indicated.

Absorbance
vo
CD

Fig. 12—Ultraviolet spectra of an aqueous solution con¬
taining cupric nitrate ([Cu+^j = 0.200 x 10“^ M) and
2-thiouracil ([2TUJ = 1.00 x 10"^ M) as a function of
pH at 25.0°. Key: (1), pH 5.59; (2), pH 6.72; (3).
pH 7.40; (4), pH 7.72; (5), pH «.10; (6), pH 9.60.

• Absorbance
200
Wavelength, m/¿

Fig. 13—Plots of absorbance of an aqueous solution containing
cupric nitrate (£Cu+^j = 0.200 x 10~^ M) and 2-thiouracil
([2TUj = 1.00 x lO"4 M) versus pH at 230 mp (3), 240 mp (â–¡),
270 mp (O). 310 mp O), 320 mp (X), and 285 mp (#). The
apparent pK values (pH at half-neutralization) are indicated.

Absorbance
202

Fig. 14—Polarogram of 2.00 x 10~^ M cupric nitrate in
the presence of 1.875 x 10”^ M 2-thiouracil and 0.0800
M HCIO^ at 21.0°. Segments A and E are portions of the
dashed difference curve obtained by subtraction of the
+2
ligand curve from the ligand + Cu curve. Lines B, C
and D are at 3/4, 1/2 and 1/4 of the vertical distance
between A and E and intersect the dashed difference
curve at ei/2 and El/4*

Micro Amps
204
Volts vs S.C.E.

Pig. 15—Potentiometric titration curves of aqueous
mixtures of cupric nitrate and 2-thiouracil containing
precipitated complex with )x = 0.006 at 25*0°, Twenty
-five ml. solutions were 2.00 x 10“3 M in 2TU and (A),
2.00 x 10”3 M; (B), 1.80 x 10“ 3 M; (C), 1.60 x 10”3 M;
(D), 1.00 x 10"3 M? (E), 6.00 x 10”4 M; (F), 2.00xl0“4
M; (G), zero M in cupric nitrate. Curve H is the titra-
“ ”* ,
tion of 25«0 mis. of 4.00 x 10 M cupric nitrate. Titer
of alkali between inflection 1 and 2 are: (D), 1.57x10“
(E), 2.90 x 10“2; (F), 4.30 x 10"2 meq.

Meq. NaOH x 10'
206-
(O
pH

Fig. 16—Potentiometric titration curves of aqueous
solutions of lead nitrate and 2-thiouracil with p = 0.006
at 25.0°. Twenty-five ml. solutions were 2.00 x 10” 3 M
in 2-thiouracil and (A), 2.00 x 10” 3 M; (B), 1.60 x 10” 3
Ms (C), 1.40 x 10” 3 Ms (D), 1.20 x 10”3 Ms (E), l.OOxlO”3
Ms (F), 6.00 x 10”^ Ms (G), 6.00 x 10”4 Ms (H), zero M
in cupric nitrate. The titer of alkali between inflec¬
tions 1 and 2 are; (B), 6.5O x 10”3; (C), 1.14 x 10”2;
(D), 1.76 x 10’2; (E), 2.10 x 10’2; (F). 2.30 x 10"2;
(G), 3«03 x 10”2 meq.

106.6
95.94
85.28
74.62
63.96
53.30
42.64
31.98
21.32
I 0.66
0

Fig. 1?—Plot of log (l-n)/n against the negative logarithm
of 6-n-propyl-2-thiouracil anion concentration (p[PTU"3)
obtained from lead nitrate (2.00 x 10" 3 M) - PTU
(2.00 x 10"3 M) mixture in water with p = 0.006 at 25»8°.
Slope of plot is 1.00 and log is 4.82.

60
40
30
20
,c«0
8
6
4
3
2
I
210
5.0
6.2
6.6

Fig. 18—Plot of log (1-ñyñ against the negative logarithm
of 6-n-propyl-2-thiouracil anion concentration (p[PTU**])
obtained, from lead nitrate (1.00 x 10“ 3 M) - PTU
(2.00 x 10“3 M) mixture in water with p = 0.006 at
25*8°. Slope of plot is 1.12.

212
p [PTU'j

Fig. 19—Plot of log (l-ñ)/ñ against the negative loga
rithm of 6-n-propyl-2-thiouracil anion concentration
(p£PTU“j) obtained from lead nitrate (2.00 x 10"^ M)
- PTU (2.00 x 10"3 M) mixture in water with p. = 0.006
at 25»8°. Slope of plot is 1.13.

214
p [PTlf]

Fig. 20—Plot of log (l-ñ)/ñ against the negative loga
rithm of 6-n-propyl-2-thiouracil anion concentration
(p[PTU~j) obtained from lead nitrate (4.00 x 10"^ M)
- PTU (2.00 x 10"3 M) mixture in water with p = 0.006
at 25.8°. Slope of plot is 1.23»

216
p [PTU‘]

Fig. 21—Plots of ñ/( l-n)[2TU] against
( _
n-2
n-1.
[211;“] from aqueous
mixtures of 2-thiouracil ([2TU] = 2.00 x 10“ 3 M) and lead nitrate
with p = 0.006 at 25*0°. Key: (1), 2.00 x 10“3 M Pb+2; (2),
1.80 x 10“3 M Pb+2; (4), 1.40 x 10"3 M Pb+2; (7), 8.00 x 10”4 M
+2
(10), 2.00 x 10"4 M Pb+2,
Pb
slope is K^K2
Intercept on ordinate is and

«o
/
O
Q>

Fig. 22—Potentiometric titration curves of aqueous mix¬
tures of cadmium nitrate and 6-n-propyl-2-thiouracil
(PTU) with jx = 0.006 at 25.0°. Twenty-five ml. solutions
were 2.00 x 10"3 M in PTU and (A), 2.00 x 10“ 3 M; (B),
1.00 x 10“3 M; (C), 4.00 x 10“4 Mj (D), zero M in Cd+2.
Curve E is the titration of 25 mis. of 8.00 x 10“4 M
cadmium nitrate.

Meq. NaOH
220
10 9 8 7 6 5 4
pH

Fig. 23—Potentiometric titration curves of aqueous
mixtures of nickel nitrate and 2-thiouracil (2TU) with
H = 0.006 at 25,0°. Twenty-five ml. solutions were
2.00 x 10-3 M in 2TU and (A), 2.00 x 10~3 M; (B),
1.20 x 10"3 M; (C), tí.00 x 10-3 M; (D), 4.00 x 10~3
M; (E), zero M in nickel nitrate. Curve F is the
¿L
titration of 25 mis. of 4.00 x 10 M nickel nitrate.

Meq. NaOH
222
Q I I I I
II 10 9 8 7 6 5
pH

Fig. 24—Plot of the change in free energy (&F) versus the absolute
temperature calculated from the stability constants of the 1:1
complexes (MU+) for the cadmium and lead nitrate complexes of
2-thiouracil (2TU), 6-n-propyl-2-thiouracil (PTU),
5,6-dimethyl-2-thiouracil (56DMTU), 6-methyl-2-thiouracil (6MTU),
5-methyl-2-thiouracil (5MTU) and 5-carboethoxy-2-thiouracil (5CETU).
Ionic strength is 0.006.

-AF Calories

Fig. 25—Plot of the change in free energy (A.F) versus the absolute
temperature calculated from the acid dissociation constant of
2-thiouracil (2TU), 6-n-propyl-2-thiouracil (PTU),
5,6-dimethyl-2-thiouracil (56DMTU), 6-methyl-2-thiouracil (6MTU),
5-methyl-2-thiouracil (5MTU) and 5-carboethoxy-2-thiouracil.
Ionic strength is 0.006.

A F Calories
Temperature, °K
226

Fig. 26—Ultraviolet spectra of aqueous solutions at
25»0° containing 2-thiouracil (5*00 x 10“^ M) and
cupric nitrate. The molarities of cupric ion were:
(0), zero M; (0.5), 0.500 x 10*’-’ M; (1), l.OOxlO--’
M; (2), 2.00 x 10"5 Ms (3), 3.00 x 10" 5 M and (5),
5.00 x 10"5 M.

Absorbance
228
Wavelength, m//

Fig. 27—Plot of log (Ab - AbH+)/AbH+ against pH for
aqueous, saturated 2-thiouracil solutions at 25*0°
according to log (Ab - Ab^+ )/Ab^+ = pH - pK& where Ab
is the absorbance of a saturated solution in the
buffer region and AbfI+ is the absorbance due to the
intrinsic solubility as measured in 0.100 M HCIO^.
The slope is 1.00 and pK is
9.
7.52.

230
pH

Fig. 28—Plots of -Í (Ei)„ - (Ei)_J against the negative logarithm
2 c 2S
of the 6-n-propyl-2-thiouracil concentration ([PTUJ) obtained from
— ¿1.
aqueous solutions 2.00 x 10 M in cupric nitrate, 0.200 M in
NaClO^ and 5.00 x 10~2 M (â–¡), 8.00 x 10"2 M (O). 1.00 x 10'1 M
(A) and 2.00 x 10“^ M (#) in HCIO^ at 21.0°. Each point is the
mean of three determinations and all solutions contained 0.001$
Triton X-100 as maximum suppressor.

-log [PTU]
232

Fig. 29—Plot of the logarithm of the Apparent Relative
Antithyroid Activity versus the negative logarithm of
the dissociation constant ratio of the 2:1 cupric
-ligand and 1:1 cuprous-ligand complexes with thioura-
cils. Antithyroid activity is relative to 1.00 for
2-thiouracil. Key: PTU, A; 5.6DMTU (exper.). A;
5.6DMTU (lit. ),O; 6MTU, V; 5MTU (lit.), X; 5MTU
(exper.), ; 2TU,#; 5CETU,CZ1. The error in -log
KQ/Kr is the standard deviation and is indicated by
the bars. The value of -log KQ/Kr for 5CETU was
estimated by adding the average difference between
the values of -log KQ/Kr for the plots of versus
-log HU with finite and zero slopes to the value of
-log KQ/Kr found from plots with zero slopes. This
was necessary in order that the values of -log K0/Kr
would correspond to the same species (MU^ and MU).

log k,
Relative Antithyroid Activity
o _ _ o
, Q 2 2 o o
N>
-P”

REFERENCES
(1) Rendiría, G., Sarcione, E. J., Lee, C. J.,
Barrett, H. W., Proc. Soc. Exptl. Biol. Med., 9¿. 350 (1957).
(2) M¿rch, P., Acta Pharmacol. Toxicol., 1, 106
(1945).
(3) Paschkis, K. E., Cantarow, A., Rakoff, A. E.,
Tillson, E. K., J. Pharmacol. Exper. Therap., S3, 270 (1945).
(4) Williams, R. H., Kay, G. A., Jandorf, B. J.,
J. Clin. Investigation, 23. 613 (1944).
(5) Williams, R. H., Kay, G. A., J. Clin. Endocri¬
nol. , 4, 385 (1944).
(6) Astwood, E. B., Harvey Lectures, 40, 195 (1944).
(7) Pitt-Rivers, R., Physiol. Rev., 30, 194 (1950).
(8) Brown, J. H. U., Barker, S. B., "Basic Endo¬
crinology" 1st ed., F. A. Davis Co., Philadelphia, Penn.,
1962, pp. 132-152.
(9) Anderson, G. W., Halverstadt, I. F., Miller,
W. H., Roblin, R. 0., J. Am. Chem. Soc., 67, 2197 (1945).
(10) Astwood, E. B., Bissell, A., Hughes, A. M.,
Endocrinology, 37. 456 (1945).
(11) Williams, R. H., Kay, G. A., J. Medical Science,
213. 198 (1947).
(12) "New and Nonofficial Drugs,” 1st. ed., J. B.
Lippincott Co., Philadelphia, Penn., Í963• P* 759.
(13) Goodman, L. S., Gilman, A., "The Pharmacological
Basis of Therapeutics," 3rd ed., Macmillan Co., New York,
N. Y., 1965, p. 1489.
(14) Miller, W. H., Roblin, R. 0., Astwood, E. B.,
J. Am. Chem. Soc., 67, 2201 (1945).
235

236
(15) Wayne, E. J., Koutras, D. A., Alexander, W. D.,
"Clinical Aspects of Iodine Metabolism," 1st ed., Blackwell
Scientific Publications, Philadelphia, Penn., 1964, p. 39»
(16) Garrett, E. R., personal communication.
(17) Seven, M. J., Johnson, L. A., "Metal-Binding in
Medicine," 1st ed., J. B. Lippincott Co., Philadelphia,
Penn., I960, pp. 321-324.
(18) Koch, H. J., Jr., Smith, E. R., J. Clin. Endo¬
crinol . , 16, 123 (1956).
(19) Kasanen, A., Viitanen, I., Acta Med. Scandinav.,
153. 467 (1956).
(20) Weiss, B., J. Biol. Chem., 201, 31 (1953).
(21) Fawcett, D. M., Kirkwood, S., J. Biol. Chem.,
205, 795 (1953).
(22) Taurog, A., Potter, G. D., Chaikoff, I. L.,
J. Biol. Chem., 213, 119 (1955).
(23) Serif, G. S., Kirkwood, S., J. Biol. Chem.,
233, 109 (1958).
(24) Libermann, D., Nature, 164, 142 (1949).
(25) Libermann, D., Bull. Ste. Chim. Biol., 31,
1325 (1949).
(26) Weiss, R., Venner, H., Z. Physiol. Chem., 317.
82 (1959).
(27) Latimer, W. M., "Oxidation Potentials," 2nd ed.,
Prentice-Hall Inc., Englewood Cliffs, N. J., 1952, p. I85.
(28) Koppel, H. C., Springer, R. H., Robins, R. K.,
Cheng, C. C., J. Org. Chem., 26, 792 (1961).
(29) Brown, D. J., "The Pyrimidines," 1st ed.,
Interscience Publishers, New York, N. Y., 1962, p. 614.
(30) Short, L. N., Thompson, H. W., J. Chem. Soc.,
168 (1952).
(31) "The Merck Index," 7th ed., Merck and Co.,
Rahway, New Jersey, I960, p. 865.

237
(32) Brown, D. J., "The Pyrimidines," 1st ed., Inter¬
science Publishers, Inc., New York, N. Y., 1962, p. 615»
(33) Brown, D. J., "The Pyrimidines," 1st ed., Inter¬
science Publishers, Inc., New York, N. Y., 1962, p. 619.
(3*0 Brown, D. J., "The Pyrimidines," 1st ed., Inter¬
science Publishers, Inc., New York, N. Y., 1962, p. 586.
(35)Lacey, R. N., J. Chem. Soc., 839 (195*0*
(36) Brown, D. J., "The Pyrimidines," 1st ed.. Inter-
science Publishers, Inc., New York, N. Y., 1962, p. 613•
(37) Reilley, C. N., Schmid, R. W., Anal. Chem.,
30, 947 (1958).
(38) Reilley, C. N., Schmid, R. W., Lamson, D. W.,
Anal. Chem., 30, 953 (1958).
(39) Pribil, R., Kondela, Z., Matzska, B., Collection
Czechoslov. Chem. Communs., 16, 80 (1951)*
(40) Gillam, A. E., Stern, E. S., "Electron Absorp¬
tion Spectroscopy," 2nd ed., Edward Arnold Publishers,
Ltd., London, England, I960, p. 289.
(41) "Instruction Manual, Sargent Recording Polaro-
graph, Model XV," Section 1.1, 1963.
(42) Kolthoff, I. M., Lingane, J. J., "Polarography,,,
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 362.
(43) Rossotti, F. J. C., Rossotti, H., "The Deter¬
mination of Stability Constants," 1st ed., McGras - Hill
Co., New York, N. Y., 1961, p. 176.
(44) Colichman, E. L., J. Am. Chem. Soc., 72, 4036
(1950). “ ~ “
(45) Schmid, R. W., Reilley, C. N., J. Am. Chem.
Soc., 80, 2087 (1958). ~
(46) Kolthoff, I. M., Lingane, J. J., "Polarography,"
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 228.
(47) Tomes, J., Collection Czechoslov. Chem.
Communs., 9» 81 (1937)*

(48) Tomes, J., Collection Czechoslov.
Communs., 9. 81 (1937).
(49) Tomes, J., Collection Czechoslov.
Communs., 9, 150 (1937*1^
(50)Hume, D. N., DeFord, D. D., Cave,
J. Am. Chem. Soc., 73, 5323 (1951)*
Chem.
Chem.
G. C. B.,
(51)Cotton, F. A., Wilkinson, G., "Advanced
Inorganic Chemistry," 1st ed., Interscience Publishers,
Inc., New York, N. Y., 1962, p. 750.
(52) Furman, N. H., "Standard Methods of Chemical
Analysis," 6th ed., D. VanNostrand Co., Inc., Princeton,
New Jersey, 1962, p. 251.
(53) Furman, N. H., "Standard Methods of Chemical
Analysis," 6th ed., D. VanNostrand Co., Inc., Princeton,
New Jersey, 1962, p. 560.
(54) Bjerrum, S., "Metal Ammine Formation in
Aqueous Solution," P. Haase and Sons, Copenhagen,
Denmark, 1941.
(55) Hedberg, D. D., "Sargent Equilibrium Constants
of Inorganic Compounds," E. H. Sargent and Co., Chicago,
Ill., 1963.
(56) Hossotti, F. J. C., Rossotti, H., "The Deter¬
mination of Stability Constants," 1st ed., McGraw - Hill
Co., New York, N. Y., 1961, pp. 355-358.
(57) Kolthoff, I. M., Lingane, J. J., "Polarography
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, pp. 211-235.
(58) Prutton, C. F., Marón, S. H., "Fundamental
Principles of Physical Chemistry," 2nd ed., Macmillan
Co., New York, N. Y., 1956, p. 530.
(59) Kolthoff, I. M., Lingane, J. J., "Polarography
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 214.
(60) Kolthoff, I. M., Lingane, J. J., "Polarography
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 43.
(61) Ibid., pp. 18-46.
(62)Ibid., p. 191

239
(63) Kolthoff, I. M., Lingane, J. J., "Polarography,"
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 52.
(64) Heyrovsky, J., "Principles of Polarography,"
1st ed., Academic Press, New York, N. Y., 1966, p. 105•
(65) Kolthoff, I. M., Lingane, J. J., "Polarography,"
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 53.
(66) Latimer, W. M., "Oxidation Potentials," 2nd
ed., Prentice-Hall Inc., Englewood Cliffs, New Jersey,
1952, pp. 12, 30, 184.
(67) Potter, E. D., "Electrochemistry," 1st ed.,
Cleaver-Hume Press, Ltd., London, England, 1956, p. 112.
(68) Lange, N. N., "Handbook of Chemistry," 9th ed.,
Handbook Press, Sandusky, Ohio, 1956, p. 950.
(69) Kolthoff, I. M., Lingane, J. J., "Polarography,"
2nd ed., vol. 1, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 224.
(70) Onstott, E. I., Laitinen, H. N., J. Am. Chem.
Soo., 72, 4724 (1950).
(71) Rossotti, F. J. C. in "Modern Coordination
Chemistry," Lewis, J., and Wilkins, R. G., ed., Inter¬
science Publishers, Inc., New York, N. Y., I960, p. 18.
(72) Job, P., Ann. Chim. Paris, 9, 113 (1928).
(73) Reid, E. E., "Organic Chemistry of Bivalent
Sulfur," vol. 3. 1st ed., Chemical Publishing Co., New
York, N. Y., I960, p. 376.
(74) Rossotti, F. J. C., Rossotti, H., "The Deter¬
mination of Stability Constants," 1st ed., McGraw - Hill
Co., New York, N. Y., 1961, p. 277.
(75) Bellamy, L. J., "The Infrared Spectra of Complex
Molecules," 2nd ed., John Wiley and Sons, Inc., New York,
N. Y., 1962, p. 283.
(76) Berecki-Biedermann, C., Arklv Kemi, 9, 175 (1956).
(77) Kolthoff, I. M., Lingane, J. J., "Polarography,"
2nd ed., vol. 2, Interscience Publishers, Inc., New York,
N. Y., 1952, p. 494.

240
(78) Pearson, E. G., Science, 151» 172 (1966).
(79) Kolthoff, I. M., Lingane, J. J., "Polarography, "
2nd ed., vol. 1, Interscience Publishers, Inc,, New York,
N. Y., 1952, p. 221.
(80) Schneider, W. C., Halverstady, I, F., J. Am.
Chem. Soc., 70, 2626 (1948).
(81) Bellamy, L. J., "The Infrared Spectra of Complex
Molecules," 2nd ed., John Wiley and Sons, Inc., New York,
N. Y., 1962, p. 283.
(82) Livingstone, S. E., Quart. Rev., 19. 386 (1965).
(83) Irving, H., Williams, R. J. P., J. Chem. Soc.,
1953, 3192.
(84)Day, M. C., Selbin, J., "Theoretical Inorganic
Chemistry," 1st ed., Reinhold Publishing Co., New York,
N. Y., 1962, p. 317.
(85)Day, M. C., Selbin, J., "Theoretical Inorganic
Chemistry," 1st ed., Reinhold Publishing Co., New York,
N. Y., 1962, p. 292.
(86)Day, M. C., Selbin, J., "Theoretical Inorganic
Chemistry," 1st ed., Reinhold Publishing Co., New York,
N. Y., 1962, p. 319.
(87) Irving, R. J., Fernellus, W. C., J. Phys. Chem.,
60, 1427 (1956).
(88) Vieback, T., Pharmacopeia Scientia Pharm., 21,
312 (1953).
(89) Astwood, E. B., J. Pharmacol, and Exper. Therajk,
78, 79 (1943). ~
(90) Horak, F., Samel, M., Collection Czechoslov.
Chem. Commun., 30, 1229 (1965).
(91) Stuckey, R. E., J. Pharm. Pharmacol., 1, 382
(1949).

BIOGRAPHICAL SKETCH
Dennis Joseph Weber was born March 30» 1934, at
Kalamazoo, Michigan. In June, 1952, he was graduated
from Santa Ana High School, Santa Ana, California. In
June, 1958, he received the degree of Bachelor of Science
from Western Michigan University. From 1958 to 1962 he
was employed by the Upjohn Company during which time he
attended the Graduate School of Western Michigan University.
He received the degree of Master of Arts with major in
chemistry in August, 1962. From September, 1962, until
the present time he has pursued his work toward the degree
of Doctor of Philosophy.
Dennis Joseph Weber is married to the former Shirley
Ann Standish and is the father of six children. He is a
member of the Society of Sigma Xi.
241

This dissertation was prepared under the direction
of the chairman of the candidate's supervisory committee
and has been approved by all members of that committee.
It was submitted to the Dean of the College of Pharmacy
and to the Graduate Council, and was approved as partial
fulfillment of the requirements for the degree of Doctor
of Philosophy.
December 19. 1967
Dean, Graduate School
Supervisory Committees