This dissertation has beenmicrofUmed exactly as received 69-16,662
O'HALLORAN, Harry John, 1939-TRANSITION METAL COMPLEXES OF0. - AMINOTHIOACIDS.
University of Hawaii, Ph.D., 1969Chemistry, inorganic
University Microfilms, Inc., Ann Arbor, Michigan
TRANSITION METAL COMPLEXES OF
a-AMINOTHIOACIDS
A DISSERTATION SUBNITTED TO THE GRADUATE DIVISION OF TIlE
UNIVERSITY OF HAWAII IN PARTIAL FULFILIMENT
OF 'IHE REQUIREMENTS FOR THE DEGREE OF
OOCTOR OF PHIlOSOPHY
IN CHEMISTRY
JANUARY 1969
By
Harry John O'Halloran
Dissertation Committeel
Jay W. Wrathall t ChairmanJohn W. Gi1jeRobert A. DueeJohn J. NaughtonChristopher Gregory
ABSTRACT
TRANSITION METAL COMPLEXES OF a-AMINOTHIOACms
Metal complexes of several a-aminothioacids have been prepared and
characterized. The ligands employed in this study were the thioacid
analogs of a-aminoacids and included thioglycine (tg). DL-thiovaline
(tv). DL-f}-phenylthioalanine (pta). DL-thiomethionine (bn). and thio-
isoleucine (ti) (a mixture of isomers). The ligands' general structure
is that of the zwitter ion RCH(N1i3
)COS-. Red. sparingly soluble nic-
kel(II) complexes were prepared from each of the ligands. and in all
cases only one product was obtained regardless of the stoichiometric
amounts of reactants involved. The elemental analysis indicate that
the complexes are the one metal to two ligand species. bis (thiogly-
cinato)nickel(II) [Ni(tg)2J• bis(thiovalinato)nickel(II) [Ni(tv)2J,bis(~-phenylthioalaninato)nickel(II) [Ni(pta)2J, bis(thiomethioninato)-nickel(II) [Ni(bn)2J• and bis(thioisoleucinato)nickel(II) [Ni(ti)2J•The low conductivity of these compounds indicate a nonelectrolYte,
molecular structure. The magnetic moments and the electronic spectra
corresponds to the supposition that the donor atoms are arranged in a
square planar configuration surrounding the nickel ions. The complexes
can exist in either the cis (I) or trans (II) form. Unfortunately. the
very slight solubility of these 'complexes rendered the methods usually
employed to distinguish between m and trans configurations experi-mentally infeasible. However, an x-ray analysis is presently in pro-
gress in these laboratories to determine the absolute structure of the
nickel-thioglycine complex.
I II
iv
Attempts to S-alkylate the ligands as either the free aoids or
while ooordinated to the niokel atom failed. Also, efforts to form
thiobridged complexes were unsuooessful.
The reaotion of cobalt(TI) aoetate with thioglyoine produoed an
extremely insoluble, dark maroom oomplex of the oomposition Co(tg)2~
The magnetio measurements suggest that the oompound may be antifer-
romagnetio, indioating some sort of metal"1lletal orbital interaotion,
perhaps by a Co-S--Co delooalization or a direot Co-Co bond. Ef-
forts to obtain oobalt complexes with the remaining ligands were un-
suooessful.
The metal ions Pt(n), Pd(II) and Cu(n) oaused deoomposition of
the ligands. '!he reaotion of the ligands with copper(II) bromide
generated oolorless solutions whioh indioated the formation of Cu(I)
as an intermediate in the deoomposition of the ligands.
A potentiometrio study over a range of temperatures was under-
taken to determine the dissociation constants of thioglyoine and the
stability constants of the nickel(TI)-thioglyaine oomplex, as well as
the thermodynamic values of these reactions. The aqueous equilibria
v
involved include the following I
+ K + - H+H3NCH2COSHa., H
3NCH2COS +,
+ - Kb - H+H3
NCH2COS ...-! H2NCH2COS +
Ni+2 L-KI NiL++ .. .
NiL+ L-K2
+ .. ~ Ni~
L H2NCH2COS-=
This sytem was then compared with the nickel(II)-glycirie system. The
K for thioglycine is much greater than the K for glycine. Also, K.a a-o
is larger for thioglycine than for glycine. Regardless of these facts,
the nickel-thioglycine system forms a more stable complex than the
corresponding nickel-glycine system. The K2 for the nickel-thioglycine
system is actually larger than the KI • The themodynamic data indicate
that this phenomenon is an entropy effect probably caused by the change
in configuration from an octahedral environment to a square planar one
for the nickel ion.
vi
TABLE OF CONTENTS
Abstract •••••••••••••••••••••••••••••••••••••••••••••••••••••••
List of Tables ••••••••••••••• ~ •••••••••••••••••••••••••••••••••
List of Figures •••••••••••••••••••••••••••••••••••••••••••••••
I. Introduction ••••••••••••••••••••••••••••••••••••••••••••
iii
viii
ix
1
A. statement of Problem •••••••••••••••••••••••••••••••• 1
B. Alfred Werner •••••••••••••••••••••••••••••••••••••••
C. Sulfur as a Donor Atom ••••••••••••••••••••••••••••••
D. Electronic Configuration of Square PlanarComplexes •••••••••••••••••••••••••••••••••••••••••••
2
5
7
E. Metal Complexes of Ligands Containing SulfurAtom.s ••••••••••••••••••••••••••••••••••••••••••••••• 10
F. S-Alkylation and S-Dealkylation Reactions of MetalComplexes Containing Su1f'ur Donor Atoms ••••••••••••• 34
G. Ct-Aminothioacids • ••••••••••••••••••••••••••••••••••• 44
II. Experimental •••••••••••••••••••••••••••••••••••••••••••• 46
A. Synthesis of Ligands •••••••••••••••••••••••••••••••• 46
B. Synthesis of Metal Complexes •••••••••••••••••••••••• 52
C. Attempts to React the Nickel Complexes withAdditional Nickel Ions •• 0........................... 59
D. Alkylation Reactions •••••••••••••••••••••••••••••••• 60
E. Visible and Ultraviolet Spectra ••••••••••••••••••••• 61
F. Infrared Absorption Spectra •••••••••••••• ~.......... 61
G. Conductivity Measurements ••••••••••••••••••••••••••• 62
H. Molecular Weight Determinations ••••••••••••••••••••• 62
I. Mass Spectra •••••••••••••••••••••••••••••••••••••••• 62
J. Magnetic Measurements ••••••••••••••••••••••••••••••• 62
vii
K. Stability Constant Measurements ••••••••••••••••••••• 63
L. Elemental Analysis •••••••••••••••••••••••••••••••••• 66
B. Equilibrium Constants •••••••••••••••••••••••••••••••
A. Magnetic Susceptibility •••••••••••••••••••••••••••••
nIt Calculations •••••••••••••••••••••••••••••••••••••••••••• 67
67
68
Results and Discussion •••••••••••••••••••••••••••••••••• 79
A. Ligands ••••••••••••••••••••••••••••••••••••••••••••• 79
B. Nicke1(II) Complexes •••••••••••••••••••••••••••••••• 82
C. Cobalt Complexes •••••••••••••••••••••••••••••••••••• 92
D. Reactions Involving Other Transition Metals ••••••••• 97
E. A1k,y1ation Reactions •••••••••••••••••••••••••••••••• 97
F. Equilibrium Constants and Thermodynamic Data forNicke1-Thiog1ycine System ••••••••••••••••••••••••••• 100
G. Conclusions ••••••••••••••••••••••••••••••••••••••••• 110
V. Appendix •••••••••••••••••••••••••••••••••••••••••••••••• 113
A.
B.
Data from the Titration Reactions
Programs for the IB1: 360 Computer
• ••••••••••••••••••
•••••••••••••••••••
113
122
VI. Bibliography •••••••••••••••••••••••••••••••••••••••••••• 129
viii
LIST OF TABIES
Table I. Factors Influencing Solution Stabilities ofComplexes ••••••••••••••••••••••••••••••••••••••••• 20
Table II • Melting Points and Elemental Analyses ••••••••••••• 84
Table III. Molar Conductances of Nickel Complexes •••••••••••• 86
Table Dl. Molar Susceptibilities and Magnetic Moments ofCompounds at 2980 K •••••••••••••••••••••••••••••• 89
Table V. U.V. - Visible Spectra ••••••••••••••••••••••••••• 90
Table VI. Magnetic Data of Co(II) Complexes •••••••••••••••• 93
Table VII. Dissociation Constants of Thioglycine •••••••••••• 101
Table VIII. Thermodynamic Data for the Proton Dissociationof Thioglycine and Glycine ••••••••••••••••••••••• 106
Table IX. Stability Constant Data for the Nickel(II)-Thioglycine and Nickel(II) -Glycine System •••••••• 108
Table X. Thermodynamic Data for the Stability of theNickel(II)-Thioglycine System •••••••••••••••••••• 109
Figure 1.
Figure 2.
Figure 3.
Figure 4.
Figure 5.
LIST OF FIGURES
Simple Energy Level Scheme for Square PlanarComplexes •••••••••••••••••••••••••••••••••••••••••••
8Ligand Field Splitting - d •••••••••••••••••••••••••
Plot of log K1 ~. liT for Thiog1ycine ••••••••••••••
Plot of log Kl 2,. liT for the Nickel(n)-Thioglycine System ••••••••••••••••••••••••••••••••••
Plot of log K2 ~. liT for the Nickel(II)-Thi~glycine System ••••••••••••••••••••••••••••••••••
i:x:
8
87
103
107
108
I. INTRODUCTION
A. statement of Problem
The most thoroughly studied of the organic ligands which form
coordination compounds with transition metals are those which contain
oxygen and/or nitrogen. as donor atoms. These are both members of the
second period of the periodic table and there are often striking dif-
ferences in the nature of the complexes formed by these donor atoms as
compared to donor atoms from later periods. Sul:f'ur is one of these
heavier donor atoms. Recently considerable interest has been shown in
sul:fur-containing ligands. However, much work needs to be done con-
cerning sulfur-containing ligands in order to more fully understand
their bonding and complexing characteristics. Therefore, this study
was initiated in the hope of providing further information concerning
metal-sulfur coordination.
The ligands employed in this work, a.-aminothioacids were chosen
for a. number of reasons. Firstly, they are ideally suited structurally
for the formation of metal complexes. Secondly, although stability
constant measurements of a. considerable number of complexes containing
sul.fur ligands have been reported, unfortunately, in most cases, infor-
mation concerning complexes with the slime metal and the analogous oxy-
gen containing ligand is not available. Such is not the case here
since complexes of a.-amino acids and transition metal are well known
and quantitative comparison can be made with the metal complexes formed
bya.-aminothioacids. Also, comparisons of the reactions of a.-amino-
thioacids with those of a.-amino acids are of interest due to the
2
biological significance of a.-amino acids. At the same time many bio-
chemical reaotions involve metal-sulfur bonding. Finally there is re-
latively little information ooncerning complexes formed by sulfur atoms
in an acyl carboxylate position due to the low ohemical stability of
suoh complexes.
The above facts suggested that a detailed study of metal complexes
of a.-aminothioacids would prove to be both interesting and soientifi-
cally rewarding.
B. Alfred. Werner
The formulation of the structural theory of organic chemistry,
which led to the rapid and brilliant development in this area, was a
consequence of the concept of valence and the principle of constancy of
valence. When these same concepts were applied to inorganic ohemistry,
the results were less than rewarding and quite baffling. In his papers
of 1891, 1892, and 1893, Werner (1) discarded the concept of valence as
a fixed directed. foroe with a fixed number of units of definite spatial
distribution which attributes a definite integral valence number to
eaoh atom. Werner postulated two types of valency or bonding for ooor-
dination compounds, consisting of a primary and a seoondary valenoe.
The primary, or ionizable valence, can only be satisfied by negative
ions as in simple salts, whereas, the secondary, or ooordinating va-
lence, oan be satisfied. by neutral moleoules as well as negative ions.
He further stated that the same variety of anion can satisfy both va-
lencies, for example, [co(NH3)fIJclz. Also, there is a fixed number
of seoondary valenoies oalled the coordination number for each central
3
ion.. The secondary valencies have definite spatial arrangements. In
one form or another this concept has remained the basis of the chemis-
try of complex compounds.
For about two decades the mechanism of Werner I s secondary coordi-
nation remained unexplained until Sidgwick and Lowry (2), working in-
dependently in 1923, showed that this secondary coordination was the
result of dative bonds formed by pairs of electrons totally provided by
the coordinating molecules or ions. This approach to bonding in coor-
dination compOlmds was extended by Pauling (3) and developed into the
valence bond theory of metal-ligand bonding, which enjoyed great popu-
larity during the 1930 I S and 1940 IS. The valence bond theory had many
major defects, however, For example, it could not predict or interpret
either spectra or detailed magnetic properties. In fact, it was unable
to account for or to predict even the relative energies of different
structures.
Another approach, the electrostatic crystal field theory, based on
ionic bonding in complexes, was suggested by Langmuir in 1919 and the
quantum mechanical theory was developed by Bethe (4) a decade later.
Basically, this theory considers complexes as composed of a central
metal ion which is surrounded by ligands held only by electrostatic
forces. For negative ions the binding force is simply the attraction
of opposite charges and for neutral molecules the binding force is due
to a dipole-ion interaction. The ligands are then considered as point
charges surrounding the metal ion. The first applications of this new
theory were made by Penney and Schlapp (5), and VanVleck. (6,7)
Although Van Vleck demonstrated the superiority of the crystal field
4
theory, it was essentially ignored for twenty years until Orgel (8)
pointed up the significance of this approach. Crystal field theory
ignores metal-ligand covalent bonds and replaces them with simple
potentials. Experimental evidence has shown that it is not always pos-
sible to disregard such bonding. Therefore, a hybridi~ation of the
pure crystal field theory with the molecular orbital theory of Mullikan
(9), making allowances for orbital overlap between the ligand and the
central metal ion, was de'veloped and is known as ligand field theory.
It is essentially a crystal field approach modified in the direction of
molecular orbital theory because of the difficulty in obtaining simple
answers from molecular orbital theory. Ligand field theory incorpo-
rates the best features of crystal field theory and molecular orbital
theory. However, as the delocali~ation of ligand electrons and orbit-
als becomes more and more important, ligand field theory becomes less
accurate.
The most general theory and sometimes the only one giving a really
satisfactory explanation for bonding in some complexes is the molecular
orbital theory, first applied to complexes by Van Vleck. (6) The mo-
lecular orbital theory postulates that overlap of orbitals will occur,
to some degree, whenever symmetry permits. Therefore, it includes the
electrostatic model with no overlap as one extreme, maximum overlap-
ping as the other extreme, and all intermediates degrees of overlap in
its scope. In fact, both valence bond and crystal field are merely
special cases of molecular orbital theory. Although this is the most
general theory, it is difficult to obtain an exact treatment for com-
plexes which contain many atoms.
5
The past ~enty five years have produced a renaissanoe in the
field of inorganic chemistry. This rebirth of the field was due to a
number of faotors, for example, improved instrumentation, the develop-
ment of theoretical aspects (i.e., crystal field theory, ligand field
theory, and molecular orbital theory), the advent of nuclear chemistry,
and the demands of industry whioh have given inorganic chemistry an ac-
celeration unequalled in the field of chemistry.
C. Su.l.:rur as a Donor Atom
There are only about a dozen or so elements in the periodic table
which exhibit a tendency to act as donor atoms in coordination ohem-
istry. These comprise most of the so-called nonmetallic elements found
at the extreme right of the periodic table. The most thoroughly stud-
ied of these coordinating species have been the ligands of the halide
ions and organic moleoules oontaining oxygen and nitrogen as the donor
atoms. Reoently, however, increased interest has been stimulated oon-
cerning organio ligands containing su.1.:fUr atoms as the coordinating
species. Even with the considerable interest now being afforded sulfur
containing ligands, a great deal of study is needed to understand fully
their bonding and complexing charaoteristics. Sulfur is one of the
heavier donor atoms. It is generally believed that sulfur atoms do not
donate as strongly as do nitrogen and oxYgen donors, and with a great
number of metal ions this is true, the stability of formation following
the order.
FLOl.Nt>S
However, for oomplexes formed from Hg(II), Cu(I), Ag(I), Au(I), Au(III),
6
Pt(II), Pd(ll), Co(ll), Fe(ll), and Cr(ll), the order of stability ap-
pears to be.
5 »N) 0 ) F« Cl (. Br 0 :> N > Cl ) Br > C ) 5e > 5 > I ') As > P ~ TeIt can be seen from this series that sulfur is considerably less eleo-
tronegative than oxygen, and even less than oarbon, whioh has very lit-
tle tendenoy to form transition metal oomplexes (with the exoeption of
olefins and other unsaturated molecules) • It must be remembered though
that the electronegativity of an atom is not a fixed and immutable con-
stant. The variation in the eleotronegativity of an atom will depend
upon the nature of the valenoe state and the c·ther atoms surrounding it.
Another faotor uponwhioh the ooordinating ability of a donor atom
will depend is the total dipole moment of that ligand.
(1)
where~o is the pennanent dipole moment,~i is the induced dipole mo-
ment, a is the polarizability, and E is the induoing eleotrostatio
field. The permanent dipole of H2
0 is greater than that of H25,
7
however, H2S is more polarizable than H20. '!herefore , if metal ions of
high field strength are used, H2S is found to coordinate better than
H20. Although the coordinating ability and the pemanent dipole moment
decrease in the series I H20) ROH )R
20, the opposite order is observed
for the sulfur containing series I H2Sl RSH" R2S. (12) Also the polari-
-2 -zability of sulfur ligands decreases in the order S >RS >R2S, which
must be kept in mind when discussing these ligands. At the same time,
the number of lone pairs of electrons and the electronic charge de-
creases in the same order.
Livingstone (13) has shown that, for both the electrostatic and
covalent models of bond formation between metal ion and a ligand con-
taining either su.l.:fUr or oxygen, the oxygen containing species should
form a stronger bond. As he points out, these models ignore one very
important factor. Sulfur, unlike oxygen, contains low-lying empty d
orbitals which can overlap with and accept back donation of electrons
from the filled d orbitals of the metal ion. This provides for cJ,.r-cl.rr
bond formation which will enhance the overall stability of the metal-
ligand bonding. '!his TT-bonding occurs with latter transition metals
in their normal oxidation states and with early transition metals in
unusually low oxidation states.
D. Electronic Configuration of Sguare Planar Complexes
For square planar complexes, the ligands form a-bonds with the
central metal ion involving the nd ? ..2.' (n+l)p , (n+l)p plus a com-r-y- x y
bination of the (n+l)s and nd 2 orbitals of the metal ion. However,z
there is usually relatively little overlap between the ligand orbitals
M orbitals ML4 orbitals L orbitals
~I ISTRONGLY 0ANTIBONDINGI
I *Iil * 1+ --- TTI TT(n+l)p I
II 2 2 (0*)I x -y
(n+l)s I talxy (TT* )
~ "t~ 2 (0*)J.t ZQ)~ I ~b.:3 (TT*)I':il I xZ,yz
Ind I TT
,-BONDING, NON -BONDING ,-
""~TT
0
'""
I BONDING I~/0FIGURE 1. SJMPLE ENERGY LEVEL SCHEME FOR SQUARE PLANAR COMPLEXES#:
#: H. B. Gray, Prog. Transition Metal ~., !, 239 (1965).
8
9
and the nd 2 orbital of the metal ion; therefore, the orbital structurez
of the metal ion may be considered as dsp2. (14) The ligands may also
form n-bonds with the central metal ion involving the nd and np orbitals
of the donor atom. According to symmetry properties the metal orbitals
which may be involved with n-bonding are nd , nd , nd , (n+1)p ,xy xz yz x
(n+1)p , and (nt-1)p. The n-bonding involves two types of bonding,y z
either electron donation from the ligand to the metal, L.... M, or e19c-
tron donation from the metal to the ligand, L+-M.
A simplified energy level diagram showing the approximate posi-
tions of the most important electronic energy levels in square planar
complexes is given in figure 1. The lowest and most stable molecular
orbitals are the a-bonding orbitals. Next is found the n-bonding and
non-bonding mole~ orbitals. At higher energies is found the so-
called antibonding molecular orbitals derived from the metal d orbitals.
The actual ordering of these four d-antibonding orbitals is still not
well defined and there remains a certain amount of controversary re-
garding them. (14,15,16,17,18,19,20,21,22,23,24) They indeed may
change depending on the nature of the n orbitals provided. by the li-
gand. The important characteristic is the fact that there are four
relatively stable orbitals and one very unstable one.
At still higher energies is found contributions from any n or-
bital systems of' the ligands. At the highest energies we find the
strongly a-antibonding orbitals.
For a dB system the above leads to the anticipation of spectral
absorption bands from three types of electronic transition. "d-d",
ligand-to-metal (L ....M) and metal-to-ligand (M-tL) charge transfer.
10
TherE; are three spin-allowed d-d transition. dxy'" dx?--r' dz2
-+d:r?-_';-'
and d ,d -+ d 2 2. Also, there are intramoleoular L'" M ohargexz yz x-y
transfer bands from all the allowed transitions desorib9d. generally as
a..:r?--I and TT..:rl--I. For oomplexes oontaining TT-acceptor ligandssuoh as those oontaining sulfur atoms, there are M.."L transitions
with electron donation from the three filled d orbitals of the metal
to the ligand orbitals.
E. Metal Complexes of Ligands Containing Sulfur Atoms
1. Su1fur-eontaining Monodentate Ligands
Organic ligands, containing sulfur atoms as the donor atoms, com-
prise a family of ligands which are extremely interesting. Most of the
sulfur-containing ligands are the thio analogs of oxygen-eontaining li-
gands • Although this is not meant to be a review of sulfur-eontaining-
ligands, a brief discussion in this area is desirable for a fuller
understanding of sulfur coordination.
Very early the affinity of thiols for Hg(II) was discovered,
whereby the name ''mercaptan'' was coined. (13) Ethyl mercaptan forms a.
polymeric complex (I) with Ni(II) having the formula Ni(SEt)2. (25)
Palladium compounds of the general f'ormula Pd(SR)2 probably have simi-
lar structures. (26) This exemplifies the bridging ability of the sul-
:fur atom plus the tendency of' sulfur to eXPand its coordination, an
area which will be discussed later. A point in fact here is the mer-
capto-bridged dimeric complexes of' Pd(II) and Pt(II) (i.e., [Pd(SR)~J2
and [Pt(SR)LzJ2
) are not readily split by p-toluidine and other uni-
dentate ligands, whereas, the halogen-bridged analogs are dissociated
11
to monomeric fragments. (27,28)
Et Et Et EtS S S s
"Ni/ "-Ni/ "-Ni/ "--Ni/ "
/ "-S/ "--S/ "-S/ "--S/Et Et Et Et
I
Of the metals which form stable complexes with mercaptans, ahost
all are of the later transition series and are, therefore, capable of
back donation of electrons forming strong M-..L rr-bonds. One glaring
discrePanCy is that of titanium IV which has no available electrons for
back donation yet forms a very stable complex, bis (cyclopentadienyl)-
titanium diethylmercaptide. (29) Attempts to form the analogous di-
CpTiClz + EtSH + base --+ C~Ti(SEt)2+ HOI + base (2)
ethoxide complex failed. The explanation given for the stabillty of
the ethylmercaptide complex suggests that, rhrough rr-bonding, electrons
are donated from the aromatic cyclopentadiene sytem to the titanium
ion, which in turn, donates these electrons via rr-bonding to the sulfur
atom of the mercaptan.
Thioethers, as a general rule, do not coordinate as well as mer-
captans and in fact do not coordinate very strongly to metals apart
from Pt(II), Pd(II), Rh(III), Ir(III), and Hg(II). Dimethyl sulfide
reacts witb. PtC~ to form three isomeric compounds of the general for-
mula Pt(O~(Me2S)2. (30) The V-isomer is the ionic salt [Pt(Me2S)4]
[PtCl~. Jensen (31), using dipole moment measurements, investigated
12
a number or a.- and f)-isomers Pt(~S)2Clz. which were among the first
metal complexes to be studied by this means. He found the a.-isomers
have a dipole moment of approximately 2.4 D, while the f)-isomers were
about 9.0-9.5 D. This clearly demonstrated that the a.-isomers have a
trans oonfiguration (II) and the f)-isomers have a ~ configuration
(III). Only the a.-form of the analogous Pd(II) complexes have been
n
isolated. (27,32) Ipatiev and Friedman (33) have prepared Pd(II) COO\-
plexes with alkyl phenyl sulfides of the composition PhSR·2PdClz. Their
probable structure (IV) are the only known Pd(II) complexes of this
type.
Ph
IR-S Cl Cl Cl Cl,,/,/,/,/
Pd Pd Pd Pd/,/,/'-./"'-
Cl Cl Cl Cl S-RIPh
IV
Unlike dia1kylsulfides. alkyl phenyl sulfides do not coordinate
reaily with mercury.
Cyclic thioethers also form complexes readily. Hendra and Powell
()4) reported thioxan (V) complexes of M(C4HSOS)2 (M::pt,Cu,Cd,Hg) which
13
were monomeric with the ligand bond through sulfur atom only.
v VI
Molybdenum complexes having the formula MoOC13
·L and NoOC132L' have
been prepared (where D=V and L'=VI). (35) The thioxan complex is pos-
sibly chloro-bridged. Davis (36) has reported a compound having the
composition Pt(SCH2CH2NHCH2CH2)C14·HCl but did not report a structure.
Thiourea (Vil) and N-N'-substituted thioureas reduce Cu(II) to
Cu(I), Au(III) to Au(I) , Pt(IV) to Pt(II), and Te(IV) to Te(II). (25)
Although urea coordinates through either the oxygen or the nitrogen
atom (37,38), thiourea coordinates through the sulfur atom only. (38,
39,40,41,42,43,44) X-ray analysis has shown that [Ni(tu)4C12] (tu=
thiourea) is octahedral with the chloride ions trans to one another.
(42,45) The compound Ni(tu)2 (NCS)2 is octahedral and polymeric, with
each sulfur atom of the thiourea bound to two different nickel atoms.
(42) Complexes of N,N~substituted thioureas have been extensively
studied. The Ni(II) complexes attain a variety of stereochemical envi-
ronments (46,47,48,49), for example, the complexes of [Ni(ntu)2X2] (X=
Cl,Br; ntu=1-(1-naphthyl)-2-thiourea) are tetrahedral, whereas the
ethylenethiourea (etu;IX) complexes show entirely different geometries.
The paramagnetic compounds [Ni(etu)4X2] (X=Cl,Br) are octahedral and
have been isolated in .2!! and trans forms, the first isolated
14
geometrical isomers of octahedral Ni(II). The iodo-complex
SIIC
H N/ "NH2 2
VIII x
[Ni(etu)412] is six-coordinate and is a rare example of a tetragonal
diamagnetic Ni(II) complex, while [Ni(etu)4] (CI04\ is diamagnetic and
square planar. As pointed out by Holt and Carlin (48), it is unusual
for one donor atom to create such diversification with Ni(II).
Whereas thiourea acts as a monodentate ligand, thiosemicarbazide
(X) behaves both as a unidentate and a bidentate ligand. (39,42,43)
Thioacetamide, CH3
CSNH2 (tam), gives the following S-bonded tetra-
hedral complexes [M(tam)4]x (M=Cu,Ag; X=C104,NO;) (50); [M(tam)2Clz](M=Co,Fe) and S-bonded, octahedral complexes [M(tam)4C12] (M=Ni,Cd).
(51) Thiobenzamide behaves similarly to thioacetamide. (52) Also
thioaotamide decomposes to give metal sulfides, a fact that is commonly
used for qualitative and quantitative analysis of metal ions.
It has long been known that thioacids (RCOSH) form transition me-
tal complexes. Ulrich (53,.54,55), in 1859, reported the formation of
salts of copper, platinum, and gold with thioacetic acid and thio-
butyric acid. However, attempts to produce the iron salt resulted in
decomposition. Sakurada (55) attempted to prepare copper and iron
complexes of thiobenzoic acid (C6H
5COSH) but could not isolate the com-
plexes. Khaletskii and Yanovitskaya (57,.58) have studied the reaction
15
of p-nitrothiobenzoic acid with ferric chloride. They found that the
acid was oxidized to the disulfide by the iron.
Nickel reacts with thiobenzoic acid to give a red brown solution,
but in the presence of pyridine an intense green solution develops from
which Krebs, et al., (59) obtained a solid complex in which the pyridine
nitrogen was also bonded to the metal. Similar zinc and cadmium com-
plexes were also isolated. .~ the complexes had the general formula
M(C6H5COS)2(Csa~)2 (Where M=Ni,Cd,Zn).
Dazinger (60) reacted cobalt salts with ammonium thioacetate in
aqueous solutions and the cobalt complex was extracted into nonaqueous
solvents, giving blue colored solutions. Although no complex actually
was claimed to have been isolated, the residue from the nonaqueous la-
yer was analyzed for cobalt, nitrogen, and sulfur and the ratio was
11214, respectively. Iron salts were treated similarly and gave red
solutions. It was suggested that thioacetic acid could be employed as
a qualitative reagent for cobalt and iron. Brdicka (61) found that
cobalt catalyzed the decomposition and oxidation of thioacids.
Heavy metal thioacetates are unstable (62) and are hydrolyzed
readily into acetic acid and the sulfides of the metal. Or, what
amounts to the same thing, the hydrogen sulfide from the decomposi tion
of the thioacetic acid precipitates the metals as sulfides. (63) Thio-
acetic acid has been strongly recommended as a substitute for hydrogen
sulfide in qualitative and quantitative analysis. (63,64,65) This has
been shown to be particularly advantageous in dealing with small quan-
tities of metals. (66)
Thioacids have also proved useful in stripping nickel plate
16
selectively from copper. The addition of thioacids to the stripping
solutions reduces the amount of copper stripped away without retarding
the stripping of the nickel plate. (67)
Even though thioacid complexes have been known for over one hun-
dred years, there have been relatively few investigations concerning
these ligands. This situation is probably due to the low chemical
stabilities of both the free ligands and their complexes.
2. Chelates and Polydentate Ligands
When a metal ion combines with a ligand containing two or more
donor atoms so that one or more rings are formed, the resulting complex
is called a chelate compound or metal chelate. The ligand is referred
to as a chelating agent or chelate. The application of Werner's coor-
dination theory made it possible to identify chelate rings and to indi-
cate their significance with respect to the stereochemistry of coor-
dination compounds. It was Werner (1) himself who first recognized the
cyclic nature of coordinate bonding in bis(ethylenediamine)platinum(II)
chloride (XI) and suggested that the 4-coordinate bonds between Ft(II)
and the amino groups of the chelating ligand were co-planar, but it
was Ley (68), while investigating the properties of copper glycinate
17
(m), who first recognized the special significance of the cyclic
structure of complex compounds. The same chelation was first used by
MOrgan and Drew (69) in 1920.
Although the high stability of metal chelate compounds became
qualitatively known and many chelate compounds were subsequently syn-
thesized, theories and concepts of chelate ring formation did not de-
velop beyond the level attained by Werner for many decades. Only
during the -past 15 years have quantitative equilibrium measurements
been made and only during the past five years has a significant amount
of data on heats and entropies of metal chelate formation become avail-
able. On the basis of this recent work, it is now possible to under-
stand more thoroughly the nature of metal chelate rings and the con-
stitutional factors which dete:rmine their special properties.
In his review, Diehl (70) pointed out that a chelating ligand is
more firmly bound than the corresponding monodentate ligand because a
breaking of one bond does not remove the ligand from the area of the
metal ion and, therefore, dissociation does not occur as readily as
with monodentate ligands. Martell and Calving (71), in reviewing the
high stabilities of the alkaline earth-EDTA chelates, have suggested
18
that the heats of coordinate bond formation in solution (i.e., relative
to the aquo metal ion) must be negligible, and that the stabilities of
the aquo chelates must be due to a favorable entropy change associated
with formation or the metal chelate compound. The stabilizing effect
of the chelate rin.g was, therefore, concluded to be an entropy effect.
The term chelate effect was first used in 1952 by Schwarzenbach. (72)
To demonstrate the nature of the chelate effect, he used as models a
bidentate ligand and two unidentate ligands which form coordinate bonds
of equivalent strength with a metal ion. He predicted increased sta-
bility for the metal chelate on the basis of standard statistical con-
siderations. The model required a zero heat of reaction in replace-
ment of two unidentate ligands by one bidentate chelating ligand, so
that the stabilizing chelate effect must be an entrpoy effect.
Schwarzenbach uso compared the stabilities of analogous complexes
and chelates that differ in the number of metal chelate rings but re-
semble each other quite closely in all other respects. For example,
one pail' of ligands studied involved the replacement of two iminodi-
acetate (HN(CHzCOO)2-2) ligands by one ethylenediaminetetraacetate
(EJJl'Ar [(OOCCHz)2NCH2CH2N(CHzCOO)2J-4). The only difference between
the two types of metal complexes formed by these ligands is one addi-tional chelate ring formed by the EDTA complex. Again he demonstrated
the increase stability for increased chelation was due to translational
entropy considerations. Adamson (73), who recalculated the formation
constants of a number of mono- and poly-amine complexes, showed that
the chelate effect due to translational entropy applies only to re-
actions in solution, and that for pure substances (i.e., for the solid
19
state) chelates are not favored over simple complexes.
Recently a large amount of experimental data on stabillty constant
(74) of metal chelates and associated heats and entropies of formation
has revealed that many factors enhance the stabilities of metal che-
lates in solution. Martell (75) has tabulated the factors influencing
solution stabilities of complexes, table 1.
There are a great number of bidentate and polydentate ligands
which contain one or more sulfur atans as donor atoms, and an exhaus-
tive inspection of the complexes formed by these ligands is far beyond
the scope of this dissertation. Therefore, only the more recent and/or
more relevant ones will be discussed here.
Ligands with two thioether groups give similar complexes to those
formed by unidentate thioethers. 1,2-Dia1ky1- and diary1-d1thio-
ethanes, RSCH2CHzSR, form stable complexes with most of the later tran-
sition metals. (30,76,77,78,79,80,81) On the other hand, 4~ethy1
thioveratro1e (XIII) exhibits little tendency to do so. (76,82) The
dithioethers form octahedral complexes with Ni(IT).
Msr'QImle~~Ie
nIT
Chelate ligands having a thioether sulfur and a ~trogen donor
atom, for example, tJ-aminothioethers. RSCHzCHzNHz (83), often coordi-
nate more strongly than dithioethers • The five~embered ring system
TABLE I
FACTORS INFIDENCING SOWTION STABILITIES OF COMPIEXES*
20
Enthalphy Effects Entropy Effects
1. Variation of bord strength 1. Number of chelate rings.#with electronegativities of
Size of chelate rings.#metal ions and ligand donor 2.atoms.
3. Changes of solv#tion on com-2. Ligand field effects. plex formation.
3. Steric and electrostatic re- 4. ArrangJments of chelatepulsions between ligan~ donor rings.groups in the complex.
5. Entropy variations in unco-4. Enthalpy effects related to ordinated ligands.
the conformation of ~he un-6. Effects resulting from dif-coordination ligand.
ferences in configuraional5. Other coulombic forces in- entropies of the ligand in
volved ;n chelate ring for- complex compounds.mation.
*A. E. Martell, Advan. Chem. Ser., 62, 272 (1967).-- -#Effects are either unique for metal chelate ring formation, as con-
trasted to coordination by unidentate ligands, or are factors which
are considerably different when chelation is involved.
21
(XIV) has proved to be a very stable oonfiguration. Mann resolved the
XIV
optical isomers of the pt(IV) complex (XV). thereby illustrating the
asymmetry of the trivalent sulfur atom. (84)
Cl
Cl
XV
It has been established through stability constant measurements
that 2-(2-methylthiomethyl)pyridine (XVI) also acts as a chelate. (85)
8-Metbylthioquinoline (XVII; N-S) gives only 111 complexes of the type
M(N-S)Xz with Pd(II). pt(II). and Hg(II) but the 211 complex
[Cu(N--5)2](C104)2haS been isolated.
22
XVI
SMe
XVil
This same five-membered ring structure (XIV) can be accomplished
by mercaptoamine type compounds, the simplest of which is ~-mercapto
ethylamine. Jensen (86) was the first to study this type of ligand.
He observed two nickel(II) complexes. The first was a green 112 com-
plex bis(~-mercaptoethylamine)nickel(II)to which he assigned the ~
structure based only on its color and insolubility in nonpolar 801-
vents. The second nickel complex was not assigned a definite stoi-
.chiometry but Jensen assumed it contained the coordinated zwitter ion;
however, the sole evidence in support of his conclusion was the pre-
sence of chloride ion in the sample. Several years later EMens and
Gibson (87) prepared a novel gold compound, (~-mercaptoethylamine)
diethylgold (XVIil). Jicha and Busch (88) reinvestigated the ligand
used by Jensen. During the course of their investigation, they
23
prepa~ed the bis(~-mercaptoethylamine)-complexesof Ni(Il) and Pd(II).
Although the unsymmetrical nature of ~-mercaptoethylaminecould con-
ceivably lead to geometric isomers among its metal complexes, there was
no evidence of such isomerization. The methods usually employed in the
separation of two geometric isomers were precluded due to the very
slight solubilities of these uncharged species. However, upon addition
of metal ions to solutions containing these complexes, the trinuclear
compounds (XIX) were formed. These complexes were also produced by re-
action of the proper stoichiometric amounts of the ligand and the metal
ion directly. This structure requires the ~ orientation of the mono-
meric species. A substantial array of trinuclear complexes containing
two different metal ions (M 1 MI) have been isolated in the studies of
Jicha and Busch. (89) A crystal structure determination of the tri-
nuclear nickel complex showed the presence of three square planar nic-
kel(II) atoms held together by sulfur bridges. In addition to the para-
magnetic trinuclear Co(II) complex [Co(CO(NH2
CH2CH2S)2)2JXz, Busch andJicha (90) also prepared the diamagnetic trinuclear Co(IlI) compound
[Co(Co (NH2CH2CH2S)3)2JX2' The proposed structure (XX) is accomplished
24
xx
by the sharing of the sulfur atoms of two tris(~llercaptoethylam1ne)
cobalt(III) ion with a third Co(III) ion so that an octahedral arrange-
ment surrounds each Co(III) atom. An analogous compound was produced
by reacting CO(NHzCHzCH2S)3 with Ni(II) according to equation 3. The
complex was found to be spin free, an indication that the nickel has an
octahedral environment much the same as the central cobalt atom in the
(3)
trinuclear Co(III) complex. Although Busch and Jicha could find no in-
dication of geometric isomers, Brubaker and Douglas (91) resolved the
optical isomers of the trinuclear Co(III) oomplex and studied the e18c-
tronic dichroism spectra. Their findings and conclusions supported
structure (XX).
25
The reactions of n-diketones with ~-mercaptoethylamineand Ni(II)
ions have resulted in the synthesis of 2,2 1-d1methyl(ethanediylidenedi-
nitrilo)diethanethiolonickel(n), 2-methyl-2 1-ethyl(ethanediylidenedi-
nitrilo)diethanethiolonickel(n), and 2-methyl-2 I -pantyl(ethanediylio-
denedinitrilo)diethanethiolonickel(II) according to equation 4. (92)
CH2C~
RI I \5"'C~" /--ot-'" I Ni + 2H30+ (4)R""'C~N/ "5
\C~C'XXI
(5)
Although compound (XXI) is diamagnetic 8.ftd square planar, the similarly
prepared (equation 5) compound (XXII) is paramagnetio and octahedral
(93), which illustrates the stabiliZing effect mercaptide ions have on
the square planar configuration of Ni(II).
A substantial number of N- and N,N-d.1substituted ~-mercapto
ethylamine nickel(n) complexes have been prepared and characterized.
26
(94) The compounds studied were the monomeric complexes NiL2
where
L =H2NCH2CH2S-, (CH3)2NCH2CH2S-, (C2H;)2NCH2CH2S-, n-C3H7NHCH2CH2S-,i-C
3H
7NHCH
2CH
2S-, t-C4H9
NHCH2
CH2S-, n-C6H13NHCH2CH2S-'
n-CaH17NHCH2CH2S-' n-CloH2lNHCH2CH2S-, C6H;CH2NHCH2CH2S-, or
C6H;CH2CH2NHCH2CH2S-, The Ni(II) complexes formed from these various
N~substituted ~-mercaptoethylamineswere found to fall into three
groups depending upon their color and solubility properties. Magnetic
moments, molecular weights, and infrared and visible spectra all indi-
cate that these complexes are square planar nickel(II) complexes. Only
two complexes are good crystalline complexes, the bis(N-iroproly-~
mercaptoethylamine) and the bis(N,N-dim.ethyl-~-mercaptoethylamine)com-
plexes. Although the dipole moment measurements of these two complexes
were limited by solubility consideration, Root and Busch were able to
place the dipole moment at a maximum of 3 D, which would indicate a
trans configuration. Girling and Amma (9;) have reported a crystal
structure for bis(N,N-dim.ethyl~~ercaptoethylamine)nickel(II)showing
the complex to have a square planar trans structure (XXIII).
Nickel(n) complexes of two addition mercaptoamines, 2,2'-
dimercaptodiethylamine, HN(CH2CHzSH)2' and
27
methyl-2,2 1-dimercaptodiethylamine, CH3N(CHZCHzSH)Z' have been reported.(96) In view of the great tendency for nickel(II) to form four-coor-
dinate complexes, a dimeric bridged structure (XXIV) was proposed for
these compounds. Cryoscopic molecular weight studies provided support
for this structure.
studies by Wrathall and Busch (97) have lead to the characteriza-
tion of three kinds of nickel(II) complexes with Z-(2-mercaptoethyl)-
O CH2 2N 's"Ni/
N/ "-S
O CHCH/2 Z-XXVI
28
XXVII
The different types of complexes were found to be readily conver-
tible one to another under the proper conditions according to the fol-
lowing (equation 6).
NiClz
(6)
Co(ll), Ni(ll), Ag(l), Hg(II) and Pb(II) complexes of 0-
aminobenzenethiol (XXVIII) and 6"111ercaptopurine (XXIX) are more stable
than the corresponding ~gen-ehelates. (98,99) Bis(2-aminobenzene-
thiolo)nickel(II) (XXX) was first reported by Hieber and Bruck (100,
101), who suggested that its insolubility was due to a polymeric
29
XXVIII XXIX xxx
2
structure involving sulfur bridges. SUbsequentl:y, Livingstone (102)
showed that this compound is diamagnetic and assigned to 'it a square
monomeric structure.
Dithiooxamide (rubeanic acid) forms insoluble c,omplexes which are
polymeric. (10.3,104) Rose and oo-worker (105) have shown that struc-
ture (XXXI) exists with x equal to at least 20.
XXXI
Kluiber (106,107) has reported Cu(II) and N1(n) oomplexes of N-
alkyltbiopioolinamides having structure XXXII. He obsl3rved that N-n-
30
R NR
RN R
butyl-6-methylthiopicolinamide reduced Cu(II) to Cu(I) and was later
able to isolate a 1.1 copper(I) complex. (108) ~
There are a considerable number of metal complexes formed from
ligands which have two sulfur atoms (dithiols) acting as donor atoms.
For example, four membered chelate rings are formed by dithioacids
(109), alkyl xanthates (110), dialkyl-dithiocarbamates (111,112) and
-dithiophosphates (lll,113,114) having the general structure XXXIII.
M
XXXIII a
XXXIII b
XXXIII c
XXXIII d
XXXITI
Y = R--C; dithioacids
Y =R-o--C; alkyl xanthatesY =R
2N; dialkyl-dithiocarbamates
Y = (00) P; dialkyl-dithiocarbamates2
Transition metal complexes possessing a square planar
31
configuration were severely limited before 1960 to the dB metal ions
Ni(II), Pd(II), pt(TI), Au(III), Rh(I), and Ir(I), with the exception
of a handful of Co(TI) and Cu(II) complexes. The best way to stabilize
the square planar geometry, it was reasoned, was to involve the metal
atom in an extensive 'IT orbital network such that the Ti stabilization
would overcome the extra stability gained when the p orbital is usedz
to form extra C1 bonds with axial groups. To achieve this end, numerous
complexes were synthesized containing the a-dithiol groups according to
structures (XXXIV) and (XXXV). (115 through 132)
R =CH3
Co, Rh, Ni, Pd, pt, Cu, Ag, Au, V, Cr, Mo, W, Re, Ru and Os
XXXIV a
XXXIV b
XXXIV c
XXXIV d
XXXIV e
M=Fe,n =2, 3
XXXIV
R = CN (mnt)R = Ph (sdt)R = CF
3R=H
x
R
XXXV
XXXV a R = N (bdt)
XXXV b R = CH3
(tdt)
x
x = 0, -1, -2, -3
32
The complexes involving the metal families Fe through Cu exhibit
two very unusual properties. First, the square planar geometry is
stabilized over a large range of metals. Second, for any given metal
and ligand, the complex is capable of stable existence in more than one
oxidation state. The various oxidation states are simply related by
one-electron transfer reactions. For example, [Ni(sdt)2] -n has been
isolated as n = 0, 1, and 3. (11.5,116) The oxidation of [Ni(sdt)2] -n
to [Ni(mnt)2] - has been described as an oxidation of Ni(II) to Ni(III).
(127,129) However, an alternative formulation or these complexes has
been described involving Ni(II) and the radical anion (XXXVI). (11.5,
116,121,133)
-3
R S - -- - --5 R
"c/: :"c/II I : II
/c,,: l/C"R 5-------5 R
XXXVI
Experimental evidence (119) has indicated that the cobalt com-
plexes [(co(mnt)2)2] -2 and [(Co(tdt)2)2] -2, which are spin paired,
quite probably involve cobalt in a dB configuration. At any rate, the
behavior of these complexes was completely different from any usual d6
Co(III) complex.
Complexes or (sdt) with V, Cr, Mo, W, Re, Ru and Os were origi-
nally reported as bis square planar complexes. (1)4) Further investi-
gation, however, established these to be six coordinate in nature.
33
(119,120,124,128,130) Attempts by Lani'ord and Arche~ (135) to resolve
the tris complexes [co(mnt)3J-3 and [MO(SZCZ(CF)2)3J, respectively,failed raising some doubts conoerning the conventional ootahedral for-
mulation of these systems and adding strength to the previously pro-
posed supposition (119) that these complexes might possess a trigonal-
prismatic structure. Finally, with the x-ray investigation of the
Re(sdt)3 complex (136,137), it was found that the rhenium ion was sur-
rounded by six sul:t'ur atoms in a nearly perfect, trigonal-prismatic
configuration. The other tris complexes showed simi] ar configurations.
(1)8,139,140)
An interesting and significant observation has been pointed out by
Gray (135) concerning the x-ray investigations of both the square planar
and trigonal-prismatic complexes. Independent of the coordination geo-
metry or the central metal ion, the S-S distance in these complexes
always takes a value close to 3.05 A. This relatively short S-S dis-
tance has been taken to indioate that there are interligand bonding
forces present in these complexes whioh are considerably stronger than
in classical octahedral, tetrahedral, or square planar complexes.
Smith, et al., (141) have prepared a novel complex comparable to
the ones formed by the a.-dithiols. They reacted Co(Il) and Ni(Il) with
1,Z-bis(mercapto)-o-earborane, obtaining square planar complexes of the
general structure (XXXVIII).
-2
XXXVII
F. S-Alkylation and S-Dealkylation Reactions of Metal Complexes Con-
taining SuJ.:rur Donor Atoms
Early in the history of complex inorganic compounds, it was re-
cognized that the properties (i.e., color, solubility, magnetic moment,
reactivity, etc.) of a metal ion could be altered depending on the type
ligand to which it was bound. Coordination chemistry has been studied
mainly by inorganic chemists whose principal interest concerned the
metal ions, and therefore, the behavior of the ligand did not attract
much attention. However, recent d9Velopnents have instigated investi-
gations concerning the reactions of coordinated ligands. F~r example,
research in biochemistry has shown the importance of coordinated metal
ions in biochemical synthesis, and in energy storage and transfer. The
search for polymers that can stand high temperatures has produced a
variety of interesting ligand reactions. Also, the role of coordina-
ting compounds as reaction intermediates has suggested the possibility
of controlling the course of organic syntheses and the nature of the
products by coordinating the reactants with metal ions. These develop-
ments have demonstrated the significance of the reactions of coordinated
3S
ligands. One area which has received attention recently has been the
alkylation of coordinated sulfur atoms. It is probably quite apparent
by now that sulfur has a tendency to form more than two bonds when in-
volved with metal ions. This is obvious from the reflection on the
many bridged complexes formed by sul.:rur containing ligands. Also, the
fact that thio-esters readily coordinate with metal ions illustrated
the ability of sul.:rur to expand its coordination.
It has long been !mown that S-alkylation occurs when certain metal
complexes of thiols are treated with alkyl halides. Some of the ear-
liest examples of this type reaotion include the alkylation of
Pt(SCZHS)Z by reacting it with ethyl or methyl iodide to form
Pt(SRZ)zIz (14Z,143), the reaction of powdered HgS with ethyl iodide at
1000 C to give Hg(S(CZHS)Z)zIZ (144), and the reaction between
Hg(SCZHS)Z and ethyl iodide which gave two products, HgIZ .S(CzHs)z and
Hg(S(CZH
S)Z)ZI
4• (14S) These early investigators, however, were not so
much interested in the metal complexes as they were in methods of pre-
paring thioethers.
Ewens and Gibson (87) demonstrated the reactivity of the coordi-
nated mercaptide group by carrying out the alkylation of diethyl-~
mercaptoethylaminegold(III) (xvnI) with alkyl halides according to
equation 7. This addition reaction occurred without destruction of the
complex but was complicated by the formation of oils, and pure com-
pounds could be isolated as solids only in the form of picrate salts.
Adams and coworkers (146,147) have shown that copper(I) derivatives of
36
alkyl halides to form thioethers.
The previously mentioned metal complexes of mercaptoamines synthe-
sized by Busch and coworkers has led the way for detailed studies of
reactions of the coordinated mercaptide group. Their experiments (148,
149) indicated that the coordinated mercaptide ions can be transformed
into thioether chelates without breaking the metal-sulfur bond. Kine-
tic measurements were made on several of these alkylation. However,
these data were complicated by consecutive reactions, dissociative
equilibria, ligand interchange, polynuclear complex formation through
bridging sulfur atoms, and solvent competition. Nevertheless, some
significant conclusions were drawn from interpretation of these rate
data.
Alkylation of the simple mononuclear nickel(II) chelate of ~
meroaptoethylamine produced the octahedral complex (xxxvnI). However,
x
~ ~
N N"H2C/" / CH2I Ni I~C /' CH2
"5 "5/I IR R
X
XXXVIIT
this reaction proved to be deceptively complex. The trinuclear complex
of nickel (XIX) was found to be &n intermediate in this reaction• When
this complex (XIX) was alkylated, the product was also the octahedral
37
complex (XXXVIII). Busch, et al., have proposed the following pos-
sible steP1'ise scheme (equations 8-11) for the alkylation of
Ni(NH2CHzCHZS)Z·
[Ni(NHZCH2CHzS-R)2Iz] + 2[Ni(NH2CHZCHzS)z] ~
[Ni(Ni(NHzCHzCHz)z)z]Iz + ZNH2CH2CHzS--R (9)
[Ni(Ni(NHZCH2CHZS)Z)2]X2 + 4RX~ Z[Ni(NH2CH2CH2S~)Z]
+ Ni+2 + ZX- (10)
The first step involves the alkylation of a small amount of the un-
charged bis complex which is in solution. Because NH2CHzCHZS- is a
more strongly coordinating ligand than NHZCHZCHZS-R, equation 9 occurs
with the formation of the trinuclear complex, which could be isolated
by interrupting the reaotion. '!he exoess alkylating agent then attaoks
the more soluble trimeric species.
The most definite kinetic measurements were oarried out on the
alkylation of bis(methyl-Z,ZI-dimeroaptodiethylamine)diniokel(II)
(XXIV) with alkyl halides. Alkylation oocurred only at the terminal
sulfur atoms, the bridged sulfur atoms being unaffeoted. The small
aotivation energies found in this study and the one mentioned just
38
above suggested the possibility of a pre-equilibrium, as shown in equa-
tion 12, in which the metal acts as an e1ectrophi1e polarizing the
k1+ RX )0·k
-1
(12)
halogen atom. The carbonium ion thus formed then attacks the sulf"ur
atom. A more recent study (150), involving the alkylation of nicke1(II)
complexes of a member of N,N-disubstituted f:3-mercaptoethylamine, has
substantiated this mechanism. These complexes, as mentioned earlier,
have a trans configuration (XXIII) and, therefore, cannot form the tri-
nuclear intermediate (XIX) formed by the bis(f:3-mercaptoethy1amine)nic-
ke1(II) complex, thus simplifying the kinetic investigation.
Rose, Root, and Busch (151) have extended the classes of reactions
of coordinated mercaptides to include those in which the a1ky1ating
agent contains a function group which can complex with the nickel ion.
They reacted the chloroacetate anion with mercaptoamine complexes of
nickel, generating octahedral complexes, [Ni(R2NCH2CH2SCH2COO)2J, con-
taining two tridentate ligands each of which possess three dissimilar
donor atoms, nitrogen, sultur, and oxygen.
39
Sulfur atoms linked to the metal atom in position 9.!! to each
other may react with bifunctional as well as monofunctional alkylating
agents, thus forming new chelate rings. Thompson and Busch (152) using
this principle developed an elegant synthetic application or these li-
gand reactions by constructing macrocyclic chelates. By reacting 2,2'-
dialkyl(ethanediylidanedinitrilo)diethanethiol nickel(II) complexes
having structure (XXXIX). This synthesis of a polycyclic chelate is an
excellent illustration of a metal ion acting as a template, holding
reactive groups in jurlaposition so that complicated multistep reaction
may occur in a sterically highly selective manner. A kinetic study
+ )
pHz~/ ,R' N Br S-~----.C-?" /I Ni
C / "-R""'" ~ N S-C'Rt ~
\ Br / --G
CH2"CH2
XXXIX
40
recently completed (lS3) indicated a pre-equilibrium mechanism for this
reaction similar to the one shown in equation 12.
Lindoy and Livingstone (154) carried out alkylation of bis(2-
aminobenzenethiolo)n1ckel(n) (xxx:) by reacting it with methyl iodide,
although attempts to employ 2-chloromethylpyridine or benzyl chloride
as alkylating agents failed.
An interesting study on the alkylation of the previously mentioned
a-dithiols was conducted by Schrauzer and Rabinowitz. (lSS) They
readily alkylated [Ni(sdt)2J-2 (XXXIV b) and [Ni( (S2C2 (CH3)2)2J-2(XXXIVe), with a number of alkyl halides forming complexes of the
general structure (XL).
XL
R =C6HSI CH3R' =CH
31 C2HS; C"'7; n-C4H9; t-C4H9; sec-csi1.71 CH2CH2CHI CH2C6HS
Only the dialkylated speoies oould be obtained and all attempts to fur-
ther alkylate the oomplex failed. It is also significant that both of
the su1:fur atoms a1kylated belonged to the same ligand. Attempts at
maorooyclization analogous to the reaotion of Busoh (equation 13),
described above, also failed. The nucleophillcity of the su1:fur atom,
41
as expected, was found to depend on the inductive effect of the
ethy1enedithio1 carbon substituent, since the nickel bisma1eonitrile
dianion ([Ni (mnt)2]-2 , (XXXIV a» could not be converted into the di-
a1ky1ated derivative. The complex with R = C6H5 and R' = CH2C6H5 wasfound to be light sensitive and debenzylated on irradiation with visi-
ble light. Pre11minary investigation of the trigonal prismatic di-
anions [M(Sdt)3J-2 (M=V,Mo,W) indicated that instead of forming thedesired methylated compounds they underwent 1,4-S-dea1ky1ation fol-
lowed by decomposition.
S-Dea1ky1ation has been known for a long time. In 1883, B1om-
strand (156) reported the S-demethylation of dimethyl sulfide in the
presence of pt,(II) according to equation 14.
(14)
Attempts to prepare a gold(Ill) complex of 8-methy1thioquinoline
(XVII) (N-SCH3
) led to the S-demethy1ation of the ligand and the isola-
tion of a gold(III) complex of 8-quinolinethi01 [AU(N-5)C12J. (157,
158)
Mono and bis chelate complexes of dimethy1-o-methy1thiopheny1-
arsine (XLI) (As-SCH3) of the types Pd(AS-SCH3)~' Pd(AS-SCH3)2Xz'
Pt(As-SCH3
)2Xz, Pt(AS-SCH3)I2 , and [pt,(As-SCH3
)2] [PU4] have been
reported. (159,160)
42
XLI
S-Demethy1ation of the ligand occurred when these compounds were re-
fluxed in dimethylformamide according to equations 15, 16, and 17.
(161)
~As S
M(AS-SCH3)2X2 > ""'M/ + 2CH3X (15)
AS/ "s~
,.......---....As S X
M(As-SCH3
)X2
)0 "M/ "M/ + 2CHf (16)x/ ""'s/ "As~
These S-demethy1ation reactions were said to be comparable to the
Zeise1 cleavage of an ether by hydrogen halides. The initial reaction
43
in the latter case involves protonation of the ether to form an oxonium
ion and cleavage then occurs by nucleophilic attack by the halide ion
on the protonated ether as shown in eCi,uation 18. A similar type of
nucleophilic attack mechanism was postulated for the S-demethylation
reaction (equation 19).
~
RSCHJ
+ M""Z + x- --+ ~CHJ + x- -+ RSM+ + CHf (19)
The complexes M(As-S)Z which had been prePared by demethylation
of the thioether complexes M(As-SCH3)zXz were easily S-alkylated,
whereas, the thiolo-bridged complexes Mz(As-S)ZXZ
could not be alky-
lated. The S-debenzylation of Pd(AS-SCHZ
C6H5
)CIZ
gav~ Pdz(AS-S)ZCIZ'
indicating that the phenomenon is not restricted to S-demethylation.
Livingstone, et a1., have also observed S-demethylation of 0-
methylthioaniline (XLII) (16Z) and diphenyl-o-methylthiophenylphos-
phine (XLIII) (163) when these ligands react with some metal ions.
01 NHZ, SCH3
XUI
The Schiff base formed by o-methylthiobenzaldehYde and N,N-
d.1ethylethylenediamine (~), which has three potential donor atoms
44
(N-N-5CH3), was found to form five coordinate complexes with co(n)
and Ni(n) having the general formula M(N-N-SCH3
)X2
• (164) In
solutions of inert solvents, the nickel halides set up a temperature-
dependent equilibria between a tetrahedral, where the sulfur atom is
not bound to the metal, and a five coordinated species. Nickel(n)
iodide catalyzes demethylation of the sulfur atom, according to
equation 20, with the formation of a square planar complex (XLV). This
complex would not be alkylat~d once formed.
CH2~---CHI I 2
CH= N N-(C2H.5)2
+ NiI2)----S
"CH3
XLIV
I~ 12CH-N N--(C2H.5)2- 'Ni / + °11:31 (20)·1---S/ "I
XLV
G. a.-Aminothioacids
In 19.53 Wieland and Sieber (164) first reported the preparation of
+a group of oompounds classified as a.-aminothioacids, NH
3CH(R)COS-.
45
These are the sulfur analogs of cr.-amino acids. A few additional papers
(166,167,168,169,170,171,172) have appeared but they dealt mainly with
the preparatory procedures. They also elucidated the melting and de-
composition points, the ultraviolet spectra, and the chromagraphic be-
havior. A few chemical characteristics have been mentioned but only in
a qualitative manner. These compounds have biological significance due
to their s:imilarity to cr.-amino aoids, the building blocks of proteins.
cr.-Aminothioacids are ideally suited structurally for the formation
of metal oomplexes. They would be expected to form five membered che-
late rings with the nitrogen and either the sulfur or the oxygen atoms
functioning as the donor atoms when coordinated to metal ions.
As illustrated above, investigations of metal complexes containing
ligands in which the sul£ur atom is found in a variety of structural
forms have been carried out. a-Aminothioacids contain a sulfur atom
which is in an acyl carboxylate position, similar to tbioacids. In
light of the aforementioned low chemical stability of thioacid com-
plexes, similar instability with cr.-aminothioacid complexes were anti-
cipated. However, the presence of the nitrogen atom, with its ability
to act as a donor atom. thus forming a chelate, was considered a stabi-
lizing factor which might permit complex formation without the decom-
position of the ligand. The stability of the -N-C-C-S- chelate
structure has already been discussed.
It was deemed important, to the further elucidation of metal-
sulfur bonding, to prepare and study the complex formation of transi-
tion metal complexes with a-aminothioacids.
46
II. EXPERIMENTAL
A.. Synthesis of Ligands
The reaotion employed for the synthesis of the ligands was oar-
ried out under striotly anhydrous oonditions. All apparatus was oven-
dried at 1100 for at least one hour. Solvents were dried by oonven-
tional methods and redistilled just prior to use.
1. Preparation of Thioglyoine, NH2CH2COSH
The first step in the preparation of thioglyoine was the synthesis
of oxazolid-21-dione. Three separate approaohes, whioh are slight
variations of the methods by Farthing (173), Coleman (174), and Bailey
(175, were employed for this purpose and are desoribed below.
(a.) Carbonyl ohloride was passed into redistilled benzyl aloo-
hoI (282 g.), with stirring, and oooling with ethanol-solid oarbon
dioxide, at such a rate that the internal temperature remained at -200
to -300 • After the internal temperature began to fall, the flow of
oarbonyl ohloride was disoontinued and the reaotion mixture was allowed
to attain room temperature. Dry, oompressed air was passed through the
solution overnight. The flask was evaouated at the asperator for five
minutes and the solution was filtered.
This solution (benzyl ohloroformate) was then added to a solution
of glyoine (196 g.) dissolved in 1.5 liters of 4N-sodium hydroxide.
After oooling, the reaotion mixture was extraoted with ether, and the
ethereal layer rejeoted. The aqueous layer was stirred with aotivated
carbon and then filtered. The filtrates and washings were cooled, with
stirring, in ice. Conoentrated hydrochlorio aoid was added dro}Hdse
47
until the mixture was acid to Congo-red. A white precipitate was ob-
tained at this point. Recrystallization from chloroform-light petro-
leum gave white needles, m. p. 1210 (N-carbobenzyloxyglycine) • Yield-
130 g. (6;%) C10HllN04 (M.W. 209.20).
Elemental Analysis I
Theoretical. C 57.41
Found I C 57.25
H 5.30
H 5.35
N 6.70
N 6.72
The N-carbobenzyloxyglycine (15 g.), acetic anhydride (15 rol.),
and thionyl chloride (10 ml.) were warmed to boiling and thionyl chlo-
ride was added until the solid dissolved. The solution was cooled and
treated with petroleum ether. Oxazolid-215-dione which separated was
collected, washed with ether and dried. Yield. - 7.0 g. (9316). This
product was immediately used for the preparation of thioglycine due to
its extreme sensitivity to moisture.
(b.) Finely ground glycine (25 g.) was stirred with dioxane (500
ml.) in a 2-1. round bottom flask fitted with gas leads and the flask
was immersed in a water bath at 400 • Carbonyl chloride was passed
through the solution for four hours. Next dry air was passed through
the solution overnight. The solvent was removed by using a rotc eva-
porator while the water bath was maintained at 400 • A white crystal-
line material was obtained. Yield. - 26.9 g. (80%). This was immedi-
ately used to prepare thioglycine.
(c.) Anhydrous sodium carbonate (106 g.) and glycine (75 g.) were
dissolved in water (450 ml.) and the solution was stirred and filtered.
Methanol (2000 ml.) was slowly and continuously added to the solution
48
dur:tng one hour, with stirring. A wite precipitate appeared after
several hundred milliliters had been added. The white, finely divided
salt (disodium methylamine-laN-dicarboxylate), was filtered off, washed
with methanol and then with ether, and dried at 1000 • Yield - 123 g.
(75%).
Disodium methylamine-l.N-dicarboxylate (50 g.) and ethyl acetate
(750 mle) were stirred at room temperature and thionyl chloride (25
ml.) was added. When the reaction had finished and the orange mixture
had cooled to room temperature, ethyl acetate (500 mle) was added and
the mixture refluxed for ten minutes and filtered hot. The filtrates
were concentrated by distillation to about 500 rol. Petroleum ether
(500 rol.) was added, the solution cooled, and oxazolid-2.5-dione fil-
tered off. Yield - 7.1 g. (23%). This was immediately used to prepare
thioglycine.
The same method (169) was employed for the preparation of thio-
g~cine regardless of the method used to obtain the oxazolid-2.5-dione
and is described below.
(d.) The oxazolid-2'5-dione (7 g.) was dissolved in N,N-dimethyl-
formamide (70 rol.). This solution was then added dropwise to a solu-
tion containing N, N-d1methylformamide (140 mle) and triethylamine (2.5
ml.), which had been saturated with hydrogen -sulfide while in an ice
bath. H2S was passed through this solution for two hours. The solu-
tion was transferred to an erlenmyer flask, stoppered and stored over-
night in a refrigerator. Part of the reaction product crystallized out
overnight. 'l.'he precipitation was completed by the addition of absolute
ether (2 1.). The white crystall1 ne product was filtered, washed with
49
absolute ether several times and dried in the air. Yield - 4.7 g.
(75%). The raw material (4.7 g.) was covered with methanol (100 ml.)
in a beaker. The methanol was brought to a boil in a water bath and
water was added droptdse until all the solid dissolved. The solution
was then cooled to obtain the purified thioglycine. The solution was
then cooled to obtain the purified thioglycine. Yield - 3.3 g. (54%)
C2H
5NOS (M.W. 91.13 g.).
Elemental Analysisl
Theoretical I C 26.36
Found I C 26.46
H 5.53
H 5.59
N 15.37
N 15.44
S 35.18
S 35.00
S 24.07
S 23.44
N 10.52
N 10.61
H 8.32
H 8.50
2. Preparation of DL-Thiovaline, (CH3)2CHCH(NH2)COSH
Finely ground DL-valine (25 g.) in dioxane (500 ml.) was treated
with carbonyl chloride at 400 for four hours. The solution was treated
as in (b) above to yield a light yellow oil. The resulting oil could
not be crystallized. Therefore, it was dissolved in N,N-dimethyl-
formamide and treated as in (d) above. Upon the addition of the ether,
a white precipitate was obtained. Yield - 26.1 g. (92%). The preci-
pitate could not be recrYStallized. C'?llNOS (M.N. 133.21 g.).
Elemental Analysisl
Theoretical I C 45.08
Found I C 44.68
3. Preparation of DL~-Phenylthioalanine,C6Hf!i2(NH2)COSH
Finely ground DL~-phenylalanine (25 g.) in dioxane (500 ml.) was
treated analogously as in (b) and (d) above. The white precipitate
50
obtained could not be recrystallized. Yield - 26.3 g. (96%) C9HII
NOS
(M.W. 181.26 g.).
Elemental Analysis.
Theoretical. C 59.66
Found. C 58 .33
H 6.58
H 6.42
N 7.73
N 7.63
S 17.10
S 17.59
4. Preparation of DL-Thioisoleucine, CH3CH2CH(CH3)CH(NH2)COSH
Finely ground DL-isoleucine (25 g.) in dioxane (500 ml.) was
treated analogously as in (b) and (d) above. The white precipitate ob-
tained could not be recrystallized. Yield - 26.8 g. (9S%) C6~3NOS
(M.W. 147.24 g.).
Elemental Analysis.
Theoretical. C 48.79
Found. C 48 .42
H 8.90
H 9.17
S 21.79
S 21.44
5. Preparation of DL-Thiomethionine, CH3SCH2CH2CH(NH2)COSH
Finely ground DL-methionine (25 g.) in dioxane (500 ml.) was
treated analogously as in (b) above to yield a dark brown oil which
could not be crystallized. The oil was dissolved in N,N-dimethyl-
formamide and treated as in (d) above. Upon the addition of ether, a
yellow precipitate was obtained. The solution was filtered and the
precipitate was washed many times with ether with the result that the
yellow color was evidently washed out and a white precipitate remained.
Yield - 24.3 g. (88%). The product could not be recrystallized.
csallNOS2 (M.W. 165.28 g.).
Sl
Nitrogen Analysisl
Theoretical, thiomethionine 8.48
methionine 9.39
Found, 9.01
6. Attempted Preparation of DL-Thioa1anine, CH3CH(NH2)COSH
(a.) DL-Alanine, finely ground (2S g.), in dioxane (SaO ml.) was
treated with carbonyl chloride at 400 for five hours. The solution was
treated as in (b) above to yield a pale brown oil described by Farthing
(173) as the unstable oxazolid-2IS-dione, 4-methy1-DL-Oxazolid-2,S-
dione. This was dissolved in N,N-dimethylformamide and treated as in
(d) above. Upon the addition of the ether, an oil ensued. In attempt-
ing to work up the oil, it turned to a brown paste which was insoluble
in water.
(b.) Benzyl chloroformate was prepared as in (1. a) above. This
solution was then added to a solution of DL-alanine (80 g.) dissolved
in Sao m1. of 4N-sodium hydroxide and treated as in (1. a) above. Again
the unstable oil of the oxazolid-2IS-dione, 4-methy1-DL-oxazolid-2,S-
dione was obtained. This was treated as in (1. d) above and an oil
ensued which, like in the previous attempt, became a paste.
7. Attempted Preparation of N-Phenylthioglycine, C&iSNHCH2COSH
Finely ground N-pheny1glycine (2S g.) in dioxane (SaO ml.) was
treated analogously as in (1. b) and (1. d) above. A brown precipitate
was obtained but gave the nitrogen analysis for N-pheny1glycine.
52
Nitrogen Analysis.
Theoretical. N-phenylthioglyoine 8.38
N-phenylglycine
Found.
B. Synthesis of Metal Complexes
1. Preparation of Bis(thioglyoinato)nickel(II). [Ni(NH2CH2COShJ.
One gram (1.10 x 10-2 moles) of thioglycine was dissolved in 100
ml. of water. To this was added, with stirring, 1.37 grams (5.5 x 10-3
moles) of Ni(C2H302)2·4H20 whioh had been dissolved in 50 ml. of water.
A bright red precipitate began to form immediately. The solution was
filtered and the precipitate was washed with absolute ethanol and then
ether. The resulting red precipitate was dried B!. vacuo over H2S0
4•
Yield - 1.2 g. (91~) (m.p. 180 d) [Ni(NH2CH2COS)2J (M.W. 240.97 grams).
Elemental Analysis.
Theoretical. C 20.ll
Found. C 20.39
H 3.33
H 3.37
N 1l.72
N 1l.74
S 26.89
S 27.10
2. Preparation of Bis(thiovalinato)nickel(II) [Ni( (CH3hCHCH(NH2)COS2J,
Ni(tv)2
(a.) One gram (7.5 x 10-3 moles) of thiovaline was suspended in
100 ml. of water. The thiovaline was not wetted by the water. To this
suspension was added, with stirring, 0.94 grams (3.7 x 10-3 moles) of
Ni(C2H302)2 ·4H20 which had been dissolved in 50 ml. of water. With con-
tinual stirring most of the thiovaline dissolved producing a dark red
53
solution. The solution was filtered to remove any of the solid aoid
whioh had not dissolved. A red-orange preoipitate rapidly formed. The
solution was filtered, the preoipitate washed with ethanol and ether,
and dried i:!l vaouo over H2S0
4• Yield - 0.5 g. (50%) (m.p. 252 d)
[Ni«CH3
)CHCH(NH2)COS)2] (M.W. 325.13 grams).
Elemental Analysis.
Theoretioal. C 37.17
Found. C 37.78
H 6.24
H 6.39
N 8.67
N 8.68
S 19.85
S 20.02
(b.) One gram (7.5 x 10-3 moles) of thiovaline was suspended in
100 ml. of aoetonitrile. forming a oolloid but not dissolving. To this
suspension was added, with stirring. 0.94 grams (3.7 x 10-3 moles) of
Ni(C2H302)2.4~0 whioh had been dissolved in 50 ml. of aoetonitrile.
The solution immediately beoame olear and dark red in oolor. A red-
orange preoipitate rapdily began to form. The solution was filtered
and the preoipitate was washed with absolute ethanol and then ether.
The produot was dried B1 vaouo over H2S04 • Yield - 0.9 g. (90%) (m.p.
252 d) [Ni«CH3)2CHCH(NH2)COS)2] (325.13 grams).
Elemental Analysis.
Theoretioal. C 37.17
Found. C 37.36
H 6.24
H 6.33
N 8.67
N 8.78
S 19.85
S 19.61
3. Preparation of Bis(thioisoleuoinato)niokel(n),
[Ni(CH3CH2CH(CH3)CH(NH2)COS) ], Ni(ti)2
(a.) One gram (6.8· x 10-:-3 moles) of thioisoleuoine was suspended
in but not wetted by 100 ml. of water. To this solution was added,
-3with stirring, 0.85 grams (3.4 x 10 moles) of Ni(C2H302)2'4H20 which
had been dissolved in 50 ml. of water. The solution was stirred and
filtered to remove any of the solid which had not dissolved. Rapidly a
red-orange precipitate formed. The solution was filtered, the precipi-
tate washed with absolute ethanol and then ether, and dried in vacuo
over H2S04 , Yield - 0.7 g. (59%) (m.p. 248 d)
[Ni(CH3
CH2
CH(CH3
)CH(NH2
)COS)2] (353.15 grams).
Elemental Analysis.
Theoretical. C 41.04
Found. C 41.23
H 6.89
H 6.78
S 18.26
S 18.18
(b.) One gram (6.8 x 10-3 moles) of thioisoleucine was suspended
in 100 ml. of acetonitrile without dissolving. To this suspension was
added, with stirring, 0.85 grams (3.4 x 10-3 moles) of Ni(C2H302)2'4H20
which had been dissolved in 50 ml. of acetonitrile. The solution be-
came clear and dark red in color. A precipitate rapidly fonned. The
solution was filtered, the precipitate washed with absolute ethanol and
then ether and dried !!l vacuo over H2
SO4
, Yield - 0.9 g. (75%) (m.p.
248 d) [Ni(CH3
CH2CH(CH3)CH(NH
2)COS)2] (M.W. 353.15 grams).
Elemental Analysis.
Theoretical. C 41.04
Found. C 41.14
H 6.89
H 6.92
N 7.98
N 8.10
S 18.26
S 18.25
4. Preparation of Bis(s-pheny1thioalaninato)nickel
55
(:J-phenylthioalanine were suspended in 150 mle of H20. To this was
added 50 ml. of water in which 1.03 grams (4.13 x 10-3 moles) of
Ni(C2H302)2·4H20 had been dissolved. The solution was stirred and fil-
tered to remove the undissolved acid. A red-orange precipitate rapidly
ensued. The solution was filtered and the precipitate washed with ab-
solute ethanol and then ether. The precipitate was dried 1u vacuo over
H2S04• Yield - 1.1 g. (54%) (m.p. 217 d) [Ni(C
6H
5CH
2CH(NH
2)COS)2]
(M.W. 419.23 grams).
Elemental Analysis.
Theoretical. C 51.57
Found. C 52.56
H 4.81
H 4.76
N 6.68
N 6.45
S 15.30
S 15.08
(b.) One and five tenths grams (8.25 x 10-3 moles) of f:J-phenyl-
thioalanine were suspended in 100 rol. of acetonitrile forming a col-
loid. To this suspension was added, with stirring, 1.03 grams (4.13 x
10-3 moles) of Ni(C2H30Z)2·4H20 which had been dissolved in 50 rol. of
acetonitrile. The solution became clear and dark red in oolor and a
red-orange precipitate rapidly formed. The solution was filtered, and
the precipitate washed with absolute ethanol and then ether, and dried
~ vacuo over H2SO4• Yield - 1.4 g. (81%) (m.p. 217 d)
[Ni(C6H5CH
2CH(NH
2)COS)2] (M.W. 419.23 grams).
Elemental Analysis.
Theoretical. C 51.57
Found. C 51.24
H 4.81
H 5.05
N 6.68
N 6.77
S 15.30
S 15.87
(c.) When N,N-dimethylformamide (IMF) was employed as a solvent
56
in place of the acetonitrile, the reaction proceeded as directly above
(b) but the precipitate was much slower in forming and the elemental
analysis indicated a molecule of solvent associated with the complex
Ni(pta)2'00',
Elemental Analysisl
Theoretical for Ni(pta)2'DMFI
c 51,23 H 5,53
Found I c 51.24 H 5,71
N 8,.54
N 8,70
S 13.03
S 13,14
Nitrogen analyses indicated a similar solvation when DMF is em-
ployed as a solvent for the preparation of Ni(tv)2 and Ni(ti)2'
Nitrogen Analysisl
TheoreticalI
Founds
Ni(tV)2'IMF
N 10,55
N 10.40
Ni(ti)2'I:MF
N 9,85
N 9.79
57
5. Preparations of Bis(thiometbionine)nickel(II) I
[Ni(CH3SCH2CH2CH(NH2)COS)2J, Ni(tm)2
* -2(a.) Two grams (1.2 x 10 moles) of thiomethionine were sus-pended in, without being wetted by, 150 mle of water. To this was
added, with stirring, 0.75 grams (3.02 x 10-3 moles) of Ni(C2H302)2'4H20
which had. been dissolved. in 50 ml. of ~O. The solution was filtered to
remove the undissolved acid. A red-orange precipitate began to form im-
mediately. The solution was filtered and the precipitate washed with
absolute ethanol and then ether, and dried !!l vacuo over ~S04' Yield
- 0.6 g. (51%) [m.p. 191 (193) d] [Ni(CH3
SCH2CH2CH(NH2)COS)2] (M.W.
387.26 grams).
Elemental Analysis.
Theoretical. C 31.02
Found. C 32.30
H 5.21
H 5.39
(b.) Two grams* (1.2 x 10-2 moles) of thiomethionine were sus-
pended in 150 ml. acetonitrile. To this suspension was· added, with
stirring, 0.75 grams (3.02 x 10-3 moles) of Ni(C2H302)2'4H20 which had
been dissolved in 50 ml. of acetonitrile. A red-orange precipitate
formed immediately. The solution was filtered and the precipitate was
washed with absolute ethanol and then ether, and dried in vacuo over
H2S04 , Yield - 1.0 g. (86%) [m.p. 191 (193) d]
[Ni(CH3
SCH2CH2CH(NH2)COS)2] , (M.W. 387.26 grams).
* This is twice the amount that would be needed in the thiomethioninewas pure. The impurity of the thiomethionine should not cause anydifficulty as long as the solubility of the metal-thiomethionine com-plex is very small compared with other possible products.
58
Elemental Analysis.
Theoretical. C 31.02
Found. C 32.06
H 5.21
H 5.12
N 7.23
N 7.23
S 33.12
S 31.85
In all of the above preparations of the nicke1-n-aminothioacid com-
p1ex, the amounts of the reactants were varied in attempts to produce
other than the two to one products obtained. In no case was more than
one product obtained nor was there any indications that any products,
other than the ones given above, were formed.
6. Preparation of Bis(thiog1ycinato)coba1t(n), [Co(NH2CH2COS)2J,
Co(tgh
Five tenths of a gram (5.5 x 10-3 moles) of thiog1ycine was dis-
