Bis(hydroxyamino)triazines: Versatile and high-affinity tridentate hydroxylamine ligands for selective iron(III) chelation

Article abstract of DOI:10.1039/b513719e

A new versatile family of chelating agents based on bis(hydroxyamino)-1,3, 5-triazines, BHTs, is described. The properties of different BHT ligands are determined by electrochemistry, spectroscopy and titrimetry revealing high redox stability, transparenc

Full text of DOI:10.1039/b513719e

View Article Online / Journal Homepage / Table of Contents for this issue
PAPER
www.rsc.org/dalton | Dalton Transactions
Bis(hydroxyamino)triazines: versatile and high-affinity tridentate
hydroxylamine ligands for selective iron(III) chelation
Irina Ekeltchik, Jenny Gun, Ovadia Lev,* Rimma Shelkov and Artem Melman*
Received 27th September 2005, Accepted 27th October 2005
First published as an Advance Article on the web 22nd November 2005
DOI: 10.1039/b513719e
A new versatile family of chelating agents based on bis(hydroxyamino)-1,3,5-triazines, BHTs, is
described. The properties of different BHT ligands are determined by electrochemistry, spectroscopy
and titrimetry revealing high redox stability, transparency in the visible range, and diprotic acid-like
behaviour in the 5–9 pH range. The iron(III) and iron(II)–BHT complexes were studied revealing high
affinity of BHTs to iron(III). Electrochemical studies show exceptional preference of the BHT ligands to
iron(III) over iron(II), this, in addition to their small size and their fast and reversible electrochemistry
makes them potentially useful electrochemical redox couples for the low end of the aqueous potential
window (<0.6 V, vs. NHE). The synthetic versatility of the new ligands allows easy tuning of the
hydrophobicity, redox potential, and to some extent the stability constant of the complexes by
alteration of the peripheral groups appended to the BHTs.
which in combination with low bioavailability prevents oral
administration of DFO and other hydroxamate drugs.
Introduction
In a recent communication14 we demonstrated that bis-
(hydroxyamino)-1,3,5-triazines (BHTs) constitute a potent class of
chelators, whose chelating properties were, as far as we know, never
reported before. Bis(hydroxyamino)-1,3,5-triazines (BHTs) had
been previously synthesized and screened as potential herbicides,16
antioxidants,17 and additives for photographic emulsions.18
The basic properties of amino substituents in 1,3,5-
triazines resemble the properties of amido derivatives. Therefore
hydroxyamino-1,3,5-triazines were expected to possess chemical
properties similar to hydroxamates without suffering from the
drawbacks associated with amide groups. The nitrogen atoms of
1,3,5-triazine cycle efficiently coordinate the iron cation thereby
fulfilling the coordinative functionality of the hydroxamate oxy-
gen. Consequently, unlike other arylhydroxylamines BHTs are not
oxidized by iron(III) and instead form strong distorted octahedral
tridentate complexes with iron(III).
In this article we provide a detailed description of the synthesis
and the fundamental properties of different model BHT ligands.
The synthetic versatility of this class of reagents is demonstrated
in the Experimental section by a large range of different amino
substituents that can be synthesized by a relatively simple reaction
scheme. Electrochemical, spectroscopic, mass spectrometric, and
acid–base titrimetric methods illuminate the ability to tune the
complexation constants, redox potential, and hydrophobicity of
the BHT–iron complexes.
Iron is an abundant element, essential for virtually all life
forms. Considerable research activity is devoted to iron overload
abatement and control of its cellular uptake.1 A large range
of medical disorders including b-thalassemia, cardiac disorders,2
malaria,3 renal failure,4 tumor cell growth,5 and Parkinson’s6 and
Alzheimer’s7 diseases are either associated with failure of iron
regulation or can be treated by administration of iron ligands.
In parallel with these activities stimulated by the biological role of
iron ligands there is now considerable siderophore research aiming
at the development of analytical8,9 and catalytic10,11 applications
and bioapplications.
Surprisingly, the massive scientific activity in this field still
concentrates only on a handful of iron(III) chelating groups. In
a recent comprehensive review Hider12a noted the importance of
catechols, hydroxamates, hydroxypyridinones, carboxylates, and
dihydroxyphenyl-1,2,4-triazoles13 as iron(III) selective chelators;
though additional groups of chelators do exist they are much less
useful. To date, despite the very intensive research activity, hydrox-
amates and most notably desferrioxamine-B (DFO), which was
introduced over three decades ago for b-thalassemia treatment, are
still the most important medical chelating agents. Hydroxamates
exhibit excellent chelating power and selectivity for iron(III) over
iron(II) and other biologically abundant cations. Desferrioxamine-
B hexacoordinates the iron centre and prevents water hydrolysis
and formation of oxo complexes which may endow undesirable
catalytic activity. The success of the hydroxamates stem, at least
partly, from the high acidity of the hydroxylamine which allows
iron binding rather than oxidation of the hydroxyamino group
and from the carbonyl coordination of the iron atom. Thus each
hydroxamate exhibits bidentate coordination. However, the amide
bond of the hydroxamates is prone to enzymatic transformation,14
Results and discussion
Synthesis
Scheme 1 depicts the general synthesis pathways for obtaining
BHT ligands. Ligands were synthesized through the monosubsti-
tution of trichloro-1,3,5-triazine with a secondary amine R1R2NH
followed by the replacement of two remaining chlorine atoms
in dichlorotriazines of type 1 with R3NHOR4. This reaction
The Institute of Chemistry, The Hebrew University of Jerusalem, Givat Ram,
Jerusalem 91904, Israel. E-mail: amelman@chem.ch.huji.ac.il; Fax: +972
26585345; Tel: +972 26585279
This journal is
The Royal Society of Chemistry 2006
Dalton Trans., 2006, 1285–1293 | 1285
©

View Article Online
Scheme 1
allows easy access also to non-symmetric derivatives such as 2j
through substitution of one chlorine atom in the corresponding
dichlorotriazine of type 1 followed by a substitution of the second
chlorine atom with a different hydroxyamino function.
For the current studies we selected a range of compounds, 2a–i,
k with R1, R2 representing different peripheral functional groups
that do not interfere with the bis[hydroxy(methyl)amino]1,3,5-
triazine backbone. As a control we used compounds 2j and 2h,
for which one or two of the hydroxylamine coordination sites are
blocked by methyl or benzyl groups.
ions of most of the ligands. Positive mode MS was not available
for 2d due to the negative charge imposed by its two acetic acid
substituents, and for 2h, which is very hydrophobic, and also had
very weak negative ion signal. 2e, g & j were not observed in the
negative ion mode. The MS-MS fragments and the corresponding
neutral losses are shown in columns 6 and 7 of Table 1. The
Chemically Induced Dissociation (CID) analysis was conducted
in positive mode at 35% residual base peak with 2 m/z isolation
width. Almost all the BHTs lost the labile OH group of the
hydroxylamine, except for 2g, whose cyclohexyl groups on the
amino head group were even more labile. Similarly, non-symmetric
BHT 2j, which has a –N(CH3)OCH3 substituent in place of the
hydroxylamine lost the methoxy group instead. Compounds 2a–
c and 2k lost a CH3N group (i.e. without losing the OH group
attached to the same nitrogen). All the BHTs except for the 2a
and 2i lost one or two of the substituents by CID (compound 2k
lost a propyl; 2c lost an acetyl and ethyl; and 2b and 2g lost their
Properties of BHT ligands
ESI-MS analysis of BHT ligands
Table 1 presents the results of the MS analysis in both positive
and negative modes. ESI-MS discerned the positive and negative
Table 1 ESI-MS of BHT ligands
Precursor ion
Elemental
composition
m/z of the positive mode fragmentation
Ligand
2a
Exact mass (M + H)+
(M − H)−
255.2
ions (relative intensity,%)
Assignment
C9H16N6O3
C17H30N6O2
256.13
350.24
257.2
351.2
240.2 (38%), 228.2(100%), 199.3 (17%)
(2a-OH + H)+; (2a-CH3N + H)+;
(2a-2CH3N + H)+
2b
349.3
334.2 (12%), 322.2 (10%), 269.3 (100%),
252.3 (10%), 240.3 (20%)
(2b-OH + H)+; (2b-CH3N + H)+;
(2b-C6H11 + 2H)+; (2b-C6H11
+
2H-OH)+ (2b-C6H11 + 2H-CH3N)+
(2c-OH + H)+; (2c-CH3N +H)+;
(2c-CH3N–OH + H)+; (2c-2C2H5
+3H)+; (2c-2C2H5 + 3H-H2O)+;
(2c-3C2H5 + 3H-H2O)+; (2c-3C2H5 +
3H-H2O–COO)+
2c
C13H22N6O6
358.16
359.2
357.2
342.2 (5%), 331.2 (15%), 313.2(100%),
303.1(22%), 285.2 (58%), 256.3 (8%),
213.3 (20%)
2d
2e
C9H14N6O6
C10H18N6O4
302.10
286.14
NA
287.2
301.2
NA
NA
NA
270.2 (25%), 259.3 (57%), 241.2(83%),
213.3(100%), 184.2 (45%)
(2e-OH + H)+; (2e-C2H5 + 2H)+;
(2e-C2H5–H2O + 2H)+;
(2e-C2H5–H2O–CH3N + 3H)+;
(2e-C2H5–H2O–CH3N + 3H)+;
NA
2f
2g
2h
2i
2j
2k
C8H14N6O4
C15H26N6O2
C29H38N6O2
C12H16N6O2
C10H18N6O3
C8H16N6O2
258.11
322.21
502.31
276.13
270.14
228.13
259.2
323.3
NA
277.2
271.2
229.2
257.1
NA
501.2
275.3
NA
NA
241.3 (100%), 159.2 (18%)
NA
260.2(100%), 248.2 (94%)
240.1(100%), 223.3 (22%),
200.3 (45%), 183.2 (20%), 171.2(100%),
157.2 (20%), 129.1 (37%)
(2g-C6H11 + 2H)+; (2g-2C6H11 + 3H)+;
NA
(2i-OH + H)+; (2i-CH3N + H)+
(2j-OCH3 + H)+; (2i-OCH3–OH + H)+
(2k-CH3N + H)+; (2k-CH3N-OH
+H)+; (2k-2CH3N +H)+;
(2k-2CH3N–CH2 + H)+;
(2k-2CH3N–(CH2)2CH3 + 2H)+;
227.4
1286 | Dalton Trans., 2006, 1285–1293
This journal is
The Royal Society of Chemistry 2006
©

View Article Online
cyclohexyl groups). This can be explained by the two bonds that
should be broken in order to lose the morpholine substituent in
2a, and the low stability of methyl or phenyl radicals for 2i.
Acid dissociation constants of the BHT ligands
As a first step in the characterization of the new set of ligands
we measured the pKa values of the different ligands by acid–base
titrimetry. In the following discussion we assign LH2 to the neutral
BHT, except for the acetic acid appended BHTs (2f and 2d) where
LH2 refers to the singly and doubly charged moieties, respectively.
Table 2 depicts the pKa,1 and pKa,2 of the BHT ligand cations,
+
LH3 . The table shows that it is possible to fine tune the basicity
of BHT ligands—which corresponds to the first deprotonation
+
constant of LH3 —by over a pH unit by appropriate selection
of the substituents of the NR1R2 group. Placing two acetic acid
residues at position R1, R2 decreases the basicity by more than
10 times compared to BHTs with dialkylamino substituents R1,
R2 such as morpholine (e.g., the pKa,1 of 2d is 6.4 whereas that
of 2a is 5.2). In contrast, the pKa,2 constants corresponding to
deprotonation of a hydroxyamino group are almost unaffected
by the peripheral group appended to NR1R2. In all cases, we
received approximately pKa,2 = 8.7 which is also very similar to
the dissociation constants of hydroxamic acids.12a,b
Fig. 1 Optical spectra of 0.3 mM of 2a taken at different pH. (a) Low pH
range: A–pH 6.2, B–pH 5.1, C–pH 3.6. (b) Moderate pH range: A–pH 9.1,
B–pH 8.3, C–pH 6.2.
Fig. 1a & b depict typical UV/Vis spectra of BHT 2a as a
function of the pH. Similar observations were recorded also for
other BHTs. For clarity the spectra at the pH range corresponding
to pKa,1 are shown in frame (a) and the spectra around pKa,2 are
depicted in frame (b). Frame (a) clearly shows the existence of
two isosbestic points at 213 and 240 nm, which separate the two
peaks of the low pH range (252 and 205 nm) from the high pH
absorbance (at 231 nm), the latter becoming dominant at pH >
pKa,1. At higher pH, Fig. 1b shows a transition from a spectrum
with an absorbance peak at 231 nm (for pKa,1 < pH < pKa,2) to a
spectrum with a shallow peak around 261 nm (for pH > pKa,2) with
an isosbestic point at 242 nm. The isosbestic points confirm the
presence of two conservative (i.e. without byproduct formation)
acid–base transitions within the environmentally and medically
relevant pH range. The pKa’s based on spectrophotometry fit the
titrimetric studies within less than 0.2 pH units.
geometry of this complex, at least as crystallized from methanol.
This information was reported in our recent communication15 and
will not be repeated here. We conducted spectrophotometric titra-
tion of all the different BHTs with iron acetate (at pH 7) in order
to determine the stoichiometry at which the absorbance reaches
saturation. Finally, electrospray mass spectrometry confirmed the
observed stoichiometry of the complex. Typical electrospray mass
spectra of the 2a–iron(III) complex are depicted in Fig. 2. The
positive ion mode reveals a base mass at m/z = 566 corresponding
to the molecular mass of [Fe(LH)2]+. The CID MS-MS study
yields a rather dull fragmentation with predominant neutral loss
of one water molecule. The negative ion mode revealed a base peak
at m/z 564 corresponding to the mass of the deprotonated anion,
FeL2−. MS analysis of all other ligands showed practically the same
pattern—instead of the OH loss from the hydroxylamine exhibited
by the ligand we observed loss of water and the products of the
Properties of iron–BHT complexes
Stoichiometry and mass spectrometry. The biligand nature
of iron ligation by BHTs was established by three independent
methods. First we conducted single crystal X-ray crystallography
of 2a which confirmed the 2 : 1 structure and distorted octahedral
Table 2 Acid dissociation constants of water soluble BHT iron ligand
+
cations, LH3
pKa,1
pKa,2
2a
2c
2d
2e
2f
5.2
5.0
8.8
8.6a
9.2
6.4
5.4a
5.8
8.7a
8.8
2i
5.6a
5.9a
8.7a
8.7
2k
a Titration in water : methanol (60 : 40% volume).
Fig. 2 Positive and negative mode ESI-MS of morpholino-BHT–iron(III)
complex.
This journal is
The Royal Society of Chemistry 2006
Dalton Trans., 2006, 1285–1293 | 1287
©

View Article Online
Table 3 Cumulative formation constants and the formal potentials of
dissociation of the groups attached to the head amine. In contrast,
solutions of 2h and 2j, i.e. the capped hydroxylamine BHTs, did
not show even a trace of the parent mass of the complex and
instead exhibited the mass spectrum of the free 2h and 2j ligands.
It should be mentioned that the exclusive 2 : 1 ligation is quite
different from the behaviour of some of the more successful iron
ligands such as the hydroxypyridinones which can accommodate
different stoichiometry depending on the precursor composition.19
BHT–iron complexes
p*bFe(III)
Eo, Va
-p*bFe(II)
pFe3+
2a
2.73
2.10
3.07
2.16
2.97
2.97
—
3.06
−1.21
−3.48
−0.80
−0.80
−0.97
−0.79
−0.88
−0.87
−0.80
−0.90
−25.96
−25.32
−29.17
−25.38
−27.55
−27.37
22.63
23.25
22.24
23.20
22.38
22.39
2c
2d
2e
2f
2i
Stability constants of iron–BHT complexes. The stability
constants of the Fe(III)–BHT complexes were determined by
competition tests against EDTA. In all cases EDTA was added
to the BHT complexes in large (usually 10 times) excess, and the
fractional decrease (x) of the peak optical density of the BHT
complex was measured. The solutions were allowed to equilibrate
for 24 h before the spectrophotometric analysis was conducted.
Since we could not resolve the pKa,3 of the ligands by acid
based titrations, we provide the cumulative formation constants
*bBHT based on the LH2 form, the most stable form under near
neutral conditions. *bBHT corresponds to the equilibrium constant
of reaction 1.
2j
2k
−27.97
−17.5
22.28
21.3
26
EDTAb
DFOc
−0.47
a E0, V vs. Ag/AgCl sat’d electrode. b Calculated based on ref. 20 and pKa’s
= 2.2, 2.54, 7.2 and 9.05. c Calculated based on ref.12 and pKa’s = 8.39,
9.03 and 9.7. E0 was taken from ref. 12a.
10−6 M Fe cations and 10−5 M of the ligand (irrespective of the
stoichiometry of the complex). The comparison is most frequently
used for medical applications since it represents physiological
conditions, but it also forms a good basis for comparison in other
fields. Table 3 shows that based on pFe3+ values, all the BHT
ligands are somewhat inferior to DFO. The level of the residual
iron(III) is approximately 3.3 pFe units lower than DFO. Setting
a constant 0.01 mM ligand level, gives an artificial preference
for the hexadentates over the tridentate ligands. Actually all the
tridentate ligands in Hider’s comprehensive review on iron(III)
chelators12a have pFe3+ values lying in the range 15–22.5. Novartis’
lead compound, Deferasirox, 4-[3,5-bis(2-hydroxyphenyl)-1,2,4-
triazol-1-yl]benzoic acid has a pFe3+ of value of 22.5.21 Bidentate
ligands have usually even lower pFe3+. For example the promising
deferiprone, 1,2-dimethyl-3-hydroxy-pyridine-4-one has a pFe3+
value of 19.12
2LH2 + Fe3+ → FeL− + 4H+
(1)
2
The cumulative stability constant was calculated based on the
stability constant of EDTA
[EDTA4−][H+]4 (1 − x)
∗bBHT = bEDTA
(2)
[LH2]2
x
or.
C
T,EDTAa4,EDTA[H+]4 (1 − x)
∗bBHT = bEDTA
(3)
2
2
CT,LH2 a1,L
x
where x is the conversion of the initial complex, i.e. the fractional
reduction of the initial optical density of the BHT–iron peaks;
square brackets denote concentrations; CT, EDTA and CT, LH2 denote
the total free, uncomplexed concentrations of the EDTA and the
BHT ligands (sum of all their acid and base forms). The different
ionization fractions, a-values, were calculated based on the pKa
values of Table 1 and reported pKa values of EDTA).20
A comparison of the calculated pFe3+–pH dependence of several
BHT complexes to the behaviour of DFO and EDTA complexes is
depicted in Fig. 3. Since the behaviour of all the BHT complexes
that were studied were rather similar, three representative BHT
ligands (2a, 2c and 2d) were chosen for clarity. Fig. 3 shows that the
Table 3 depicts the *b-values of the examined BHTs and the
corresponding *b-values of EDTA and DFO. *bEDTA corresponds
to the reaction with EDTAH4:
[EDTA H4] + Fe3+ → [FeEDTA]− + 4H+
(4)
* bEDTA differs from the cumulative stability constant bEDTA which
corresponds to the reaction
EDTA4− + Fe3+ → [FeEDTA]−
(5)
The peripheral groups at position R1, R2 have only a moderate
effect on p*bFe(III) the maximal difference between all the seven
different BHTs, which bear widely different substituents on the
head amine group, ranged between p*bFe(III) = 3.07 for the diacetic
acid modified BHT to around 2.1 for two EtO2CCH2 modifiers.
Although the changes in p*bFe(III) values are limited, the trend
is apparent, electron donor substituents at positions R1 and R2
decrease the complex formation constant.
Since the stoichiometry of the BHT complexes differs from
that of DFO and EDTA complexes, the comparison between the
affinities can better be based on pFe3+ values at pH 7.4. pFe3+
is defined as −log[Fe3+] in pH 7.4 aqueous solution containing
Fig. 3 pFe3+–pH dependence: BHT concentration 10 lM, total iron(III)
concentration 1 lM.
1288 | Dalton Trans., 2006, 1285–1293
This journal is
The Royal Society of Chemistry 2006
©

View Article Online
pFe3+ values of the BHT complexes lie for the most part between
the values corresponding to DFO and EDTA, except at the
high pH range where the BHT complexes also outperform DFO
chelation. The figure also illuminates the high pH dependence of
the chelation power and the residual iron concentration. High pH
dependence implies that iron regulation is pH sensitive. The pFe3+
vs. pH slope of the BHT–iron complexes in the near neutral range
is about 4 reflecting the dominant chelation reaction (in the pKa,1 >
pH > Ka,2 range):
2LH2 + Fe3+ ꢀ FeL2 + 4H+
(6)
−
The slope of the 2d moiety is somewhat less steep than the other
BHTs due to its higher pKa,1 value (6.4).
The formal potential of the BHT iron complexes. The formal
potential of the complexes is thermodynamically informative,
and it also has practical consequences for electrocatalysis and
electrosensing. The range of the available charge mediators at the
very low end of the water potential window is rather limited, and
an additional class of charge mediators at such low potentials
increases the opportunities for charge shuttling between remote
redox centres and enzymatic cofactors and conductive surfaces.
This may become particularly useful if the formal potential and
the properties of the mediators can be fine-tuned by appropriate
selection of ligands. We used cyclic voltammetry in order to
determine the pH dependent formal potential of BHT–iron
complexes. When the Fe(II) and Fe(III) complexes have the same
coordination then the formal redox potential is given by eqn (7):
Fig. 4 Cyclic voltammetry of 1 mM 2a–iron(III) complex. Frame (a):
(1) blank; (2) pH 12.3; (3) pH 4.6; (4) pH 3.6; (5) 2 mM solution of 2a
pH 3.6. Frame (b): (1) pH 12.3; (3) pH 3.6; (4) pH 4.6. Frame (c): pH
dependence of E0 of the complex in 0.5 M KCl water solution.
The second irreversible peak (of Fig. 4a) appears only when
the potential is lowered and it gradually increases as the pH is
decreased. We attribute this peak to the destruction of the complex,
release of the free BHT, and its reduction on the electrode. Indeed,
curves 3 and 4 in Fig. 4a show that the CV of the free ligand is
almost identical to the CV of the destroyed complex.
The BHT–iron complexes of 2a, i, j, e and 2c exhibited well
defined anodic and cathodic peaks in cyclic voltammetry studies
(Fig. 5), which allowed determination of their formal potential
over a fairly large pH range. In all cases the formal potential
was constant at least in the pH range 6–10. Table 3 reports the
formal potentials of the different BHT complexes. It shows that in
ꢁ
ꢂ
∗bBHT,Fe
∗bBHT,Fe
2+
3+
ꢀ
ꢀ
0
Ec0omplex = E
 + 59.16 log
(7)
ꢀ
2+
Fe


3+
Fe
where the different terms have their intuitive meaning. The redox
potential should be below −440 mV at pH 7.4 in order to guarantee
that the complex will be precluded from participation in Fenton
reactions under physiological conditions. A redox potential below
hydrogen evolution practically guarantees that iron will not cycle
between its di- and trivalent state and will not catalyse the
formation of potent oxo radicals and hydrogen peroxide during the
redox cycle. Indeed, immediately after the addition of iron(II) to
an aqueous solution containing the different ligands the solution
turned purple due to the formation of BHT–iron(III) complexes.
The formal potential and the selectivity of iron(III) over iron(II)
were determined by electrochemical studies using a hanging-
mercury drop electrode in 1 M KCl adjusted to the set pH with
concentrated HCl or NaOH solutions. The general features of the
cyclic voltammetry are depicted in Fig. 4, using Fe(III)–2a as a
model complex.
Taking pH 4.6 as an example, one may notice two peaks in
Fig. 4a. The first, smaller peak is attributed to the reversible
reduction of the iron complex, and the second to the reduction
of the ligand. The redox peaks of the complex start at 0.85 V at
the high pH end (Fig. 4c). As the pH is decreased the voltammetry
becomes almost reversible, and the formal potential is slightly
shifted to more positive values, until it reaches −0.80 V and
remains constant for the entire pH range, 5.0–10 (Frame c of
Fig. 4). At pH < 5.0 the redox peaks of the complex start to shift
to lower potentials the reduction peak height is decreased until it
disappears at pH 4.2.
Fig. 5 CV of different BHT iron complexes using a hanging mercury
drop electrode. Test conditions: Sat’d. Ag/AgCl reference; pH 7.5; 1 M
KCl; 50 mV s−1 scan rate.
This journal is
The Royal Society of Chemistry 2006
Dalton Trans., 2006, 1285–1293 | 1289
©

View Article Online
all cases the formal potential ranged between −0.78 and −1.0 V
vs. Ag/AgCl, substantially lower compared to DFO and EDTA
complexes. Table 3 and Fig. 5 show that the E0 of the redox pair
can be tuned by appropriate selection of peripheral modifiers.
Incorporation of electron donor R1 and R2 (e.g., when R1 = R2 =
CH2CO2−, as for 2d) shifts the E0 of the complex by over 200 mV
compared to BHTs with more electron withdrawing groups such
R1 = R2 = CH2CO2Et or a morpholine cycle.
Based on the measured E0 and eqn (7) we could calculate the
formation constants of the Fe(II) complex as depicted in column
4 of Table 3. It can be seen that Fe(II) affinity is substantially
decreased for harder electron donor substituents (e.g., 2d, 2k and 2f
corresponding to two acetates, hydrogen and monoacetate amine
substituents). The observed pecking order is 2d > 2k > 2f > 2i >
2a > 2e > 2c, which largely follow the redox potential trend and
the Fe(III) affinity ranking.
Octanol–water partition coefficients of the complexes. The
hydrophobicity of the complexes is of vital importance for
therapeutics, and to some extent it is also important for sensing
and other analytical applications. The synthetic versatility of
BHTs allows easy manipulation of the complex hydrophobicity.
In order to demonstrate this property we examined the octanol–
water partition coefficients of several BHT–iron(III) complexes.
Table 4 shows that it is possible to tune the hydrophobicity of
the complexes by appending hydrophobic groups or hydrophilic
groups onto the triazine backbone. Over three orders of magnitude
difference in the octanol–water partition coefficient is achieved by
replacing the two cyclohexyl groups in position R1 and R2 of 2b
by two acetic acid residues (2d).
Fig. 6 (a) Spectrum of 2a–iron(III) complex at different pH curve
(1) pH 4–11; curve (2) pH 2.4; Visible range spectra (A) pH 2.9; (B) pH 3.6;
(C) pH 4.6–11. (b) pH dependence of the absorbance at the peak
wavelength (520 nM). The concentration of 2a was in all cases 0.3 mM.
absorptivity were insensitive to the pH, and a significant change
of the absorbance took place only at pH < 5. This deterioration
of the optical signal coincided with the deterioration of the
electrochemical response, which indicates the instability of the
complex at the very low pH range.
The absorption wavelength was almost invariant to the hy-
drophobicity of the solvent, and the change from water to octanol
solutions had little effect on the optical characteristics. Again, 2d
and 2f behaved differently, protonation of the acetates increased
the absorptivity substantially, in accordance with the postulate
that was proposed before.
Spectrophotometric attributes. An important property of the
BHT–iron(III) complexes is their intense purple color which is
favourable for colorimetric assays, since natural samples have
much less color interferences in this range. The maximum molar
absorptivity and the peak wavelengths of BHT complexes are
presented in Table 4. All the BHT–iron(III) complexes, except for
the acetic acid modified ones, exhibited very similar behaviour
with molar absorptivity, e = 2240–2500 mol−1 cm−1 and a broad
absorbance band, typical of charge transfer complexes. The only
exception, from the optical point of view, are the acetic acid
containing complexes (2d and 2f) which exhibited substantially
lower molar absorptivities (1290 and 920, respectively). In the UV
region, the complexes retained the optical band of the ligands
(see Fig. 6a). The visible absorbance peak wavelength and molar
Competing metal ions
We examined the effect of competing metal ions on iron(III)
ligation by introduction of 50 mM of different metal chloride
salts, as indicated in Table 5 to a 0.5 mM solution of 2a–iron(III)
complex. The tests were conducted in 0.2 M acetic acid buffer. The
first and second rows of Table 5 describe the residual percentage of
the iron complex after addition of the different metal chloride salts.
The symbol “a” marks cases were precipitate formation precluded
analysis. Since the acetate complexes, interfere and reduce the
concentration of the free metal interferences we have added in
rows 3 and 4 of the Table 5 the free residual concentration of the
respective ions as calculated by MinteqTM22 (a US Environmental
Protection Agency program for thermodynamic speciation). The
most dominant forms of the competing ions are also introduced.
These values must put into the right perspective as the low
observed competition from the various ions stems from that (at
least at pH 7.4) many of these ions are bound and not free
to interact with iron. Nevertheless the table demonstrates that
Cu2+ and Cr3+ are the most dominant possible interfering ions.
Colorimetric determination of iron(III) by 2a ligands is almost
indifferent (<11%) to all other ions except for cobalt at least at
pH 5.
Table 4 Physicochemical properties of ferric complexesa
Ligand
kw/nm
ew
ko/nm
eo
logKow
2a
2b
2c
2d
2e
2f
2g
2i
2k
530
NA
530
541
530
592
NA
520
523
1710
NA
2240
1290
2250
920
NA
2400
2450
535
530
542
538
530
548
502
544
519
2830
2640
2710
2080
2080
2500
870
2.45
3.17
2.26
−0.18
2.55
1.00
3.17
1.85
2.74
2170
2060
a NA = not available; w and o denote aqueous octanol solutions, Kow
=
octanol water partition coefficient.
1290 | Dalton Trans., 2006, 1285–1293
This journal is
The Royal Society of Chemistry 2006
©

View Article Online
Concluding remarks
BHTs provide a new general group of iron ligands. Their high
iron(III) affinity, low redox potential, tridentate character, small
size, and above all the versatility of their synthesis which allows
tuning of their physico-chemical properties open the door for
widely different potential applications in medicine, plant nutrition,
analysis and bioanalysis.
Experimental
Materials
Zinc chloride dihydrate was purchased from Merck (Darm-
stadt, Germany). Potassium chloride, ethanol and iron(II) acetate
(99.99%) were from Aldrich (Milwaukee WI). Iron(III) chloride
anhydrous (98%) and ethylenediaminetetraacetate disodium salt
dehydrate (EDTA) were from BDH (Poole, England). Sulfu-
ric acid (98%), hydrochloric acid (37%), and methanol were
from J. T. Baker (Deventer, Holland). Sodium hydroxide was
from Frutarom (Haifa, Israel). Ammonium acetate was from
Mallinckrodt (Phillipsburg, NJ, USA). N-octyl alcohol was from
Riedel-de Hae¨n (Seelze, Germany). 2,4,6-Trichloro-1,3,5-triazine
and amines were from Aldrich, N-methylhydroxylamine was from
Fluka, Unless otherwise stated all reagents were of analytical
grade and were used as received. We used deionized water
(conductivity < 0.1 mS cm−1) purified by a Seradest SD 2000
system.
Analytical procedures and instruments. The electrochemistry
of the ligands was studied on a glassy carbon electrode. A
featureless cyclic voltammetry between hydrogen evolution in the
cathodic side and the ligand oxidation starting at approx. 750 mV
vs. Ag/AgCl reference was observed at pH 7.2 in 0.5 M KCl
supporting electrolyte (not shown). A very similar response was
obtained for hydroxylamine solution under identical conditions.
Electrochemical characterization of BHT–iron complexes was
carried out using a three-electrode electrochemical cell equipped
with a hanging mercury drop electrode, sat’d. Ag/AgCl reference
electrode, and glassy carbon counter electrode. For the electroox-
idation studies we used a PARC 273 potentiostat/galvanostat
(EG&G, Princeton) and platinum or GC disk working electrodes.
All potentials are reported vs. KCl sat’d. Ag/AgCl reference.
Unless mentioned specifically, 1H NMR spectra were recorded
on Bruker AMX-300 and DRX-400 spectrometers. Elemental
analysis was made in the Microanalysis Laboratory of the Hebrew
University.
Electrospray Mass Spectrometry was conducted using a LCQ
(Thermo-Quest, Finnigan, San Jose´, CA, USA) ion trap mass
spectrometer equipped with an electrospray ionization interface
(ESI). The ESI was operated either in negative ion mode using
4.0 kV spray voltage and −20 V cone voltage or in positive mode
with 4.5 kV spray voltage 33 V capillary voltage.
Titration was conducted using Titroprocessor 672 from
Metrohm. Typically we used 0.5 mM of the ligand or the complex
in distilled water and 10 mM sulfuric acid or soda caustic as a
titrant. When it was impossible to discern the inflection points
based on purely aqueous solution due to the low solubility of
some of the analytes, we used 50% (v) water methanol solution for
titrimetry. The relevant data are indicated by * symbols in Table 2.
This journal is
The Royal Society of Chemistry 2006
Dalton Trans., 2006, 1285–1293 | 1291
©

View Article Online
General procedure for preparation of ligands 2a–c, 2e, 2g–h. To
a solution of 2,4,6-trichloro-1,3,5-triazine (18.4 g, 100 mmol) in
ethyl acetate (300 mL) a solution of a corresponding secondary
amine R1R2NH (200 mmol) in THF (100 ml) was added dropwise
at −2 ◦C under vigorous stirring. The reaction was continued
15 min after completing the addition and filtered. The filtrate
was evaporated to afford crystalline 2,6-dichloro-4-R1,R2N-1,3,5-
triazines that are purified by recrystallization from isopropanol.
To a suspension of resultant recrystallized 2,6-dichloro-4-R1,R2N-
1,3,5-triazines (30 mmol) in dioxane at 0 ◦C (90 ml) was added
a solution of R3NHOR4 hydrochloride (120 mmol) and NaOH
(5.4 g, 108 mmol) in water (20 mL). The reaction mixture was
stirred for 14 h, diluted with water and the precipitated BHT
ligands 2a–c, 2e, 2g–h were collected by filtration and recrystallized
from isopropanol.
6H); 7.43 (d, J = 6.8 Hz, 4H); 7.53 (br. s, 2H). 13C NMR (CDCl3,
75 MHz): 25.5, 26.3, 30.4, 55.5, 78.1, 128.1, 128.3, 128.8, 136.3,
164.8, 167.9. Calcd. for C29H38N6O2: C, 69.29; H, 7.62; N, 16.72.
Found: C, 69.24; H, 7.79; N, 16.42.
2,6-Bis[hydroxy(methyl)amino]-4-(N-methyl, N-phenyl)-1,3,5-
1
triazine (2i). Yield 52%. H NMR (CDCl3, 400 MHz): 3.50 (s,
6H); 3.67 (s, 3H); 7.42 (t, J = 7.3 Hz; 1H); 7.47 (d, J = 7.4 Hz, 2H),
7.56 (t, J = 7.3 Hz, 2H). 13C NMR (CDCl3, 75 MHz): 36.9; 37.8;
126.0; 126.5; 128.5; 144.0; 162.9; 163.3. Calcd. for C12H16N6O2: C,
52.16; H, 5.84; N, 30.45. Found: C, 51.98; H, 5.83; N, 30.35.
2,6-Bis[hydroxy(methyl)amino]-4-(N-propylamino)-1,3,5-triazine
(2h). Yield 48%. 1H NMR (CDCl3, 400 MHz): 0.95 (t, J =
7.3 Hz, 3H); 1.59 (sextet, J = 7.2 Hz, 2H); 3.36 (br. approx. s,
8H). 13C NMR (CDCl3, 75 MHz): 11.3; 22.6; 37.4ꢀ 42.6; 162.7;
163.1. Calcd. for C8H16N6O2: C, 42.10; H, 7.07; N, 36.82. Found:
C, 42.21; H, 6.97; N, 36.96.
2,6 - Bis[hydroxy(methyl)amino]) - 4 - morpholino - 1,3,5 - triazine
◦
1
(2a). Yield 40%. Mp 198–199 C. H NMR (300 MHz, D2O):
3.27(s, 2H); 3.42 (s, 6H); 3.67 (t, J = 4.5 Hz, 4H); 3.84 (t, J =
4.5 Hz, 4H). 13C NMR (CDCl3, 75 MHz): 36.8, 44.0, 66.6, 175.7,
177.2. Calcd. for C9H16N6O3: C, 42.18; H, 6.28; N, 32.79. Found:
C, 42.34; H, 6.28; N, 33.04.
2-(Methoxyamino)-6-(hydroxyamino)-4-morpholino-1,3,5-triazine
(2j). To a solution of 2,6-dichloro-4-morpholino-1,3,5-triazine
(1.43 g, 5 mmol) in dioxane (15 mL was added to a solution
of NaOH (0.40 g, 10 mmol) and N,O-dimethylhydroxylamine
hydrochloride (0.98 g, 10 mmol) in water (1 ml). The reaction
mixture was stirred for 2 h, and diluted with 1 : 1 EtOAc–hexane
mixture (50 ml). Organic layer was separated, washed with
water, dried, and evaporated. The residue was purified by
flash chromatography (20 to 30% EtOAc–hexane) to afford
2-(methoxyamino)-6-chloro-4-morpholino-1,3,5-triazine which
was dissolved in dioxane (10 mL) and treated with a solution
of NaOH (0.40 g, 10 mmol) and N-methylhydroxylamine
hydrochloride (0.84 g, 10 mmol) in water (1 mL). The reaction
mixture was stirred for 2.5 h at room temperature, evaporated,
dissolved in EtOAc, washed with water and brine, dried, and
2,6 - Bis[hydroxy(methyl)amino] - 4 - dicyclohexylamino - 1,3,5-
triazine (2b). Yield 46%. Mp 202–203 ◦C. 1H NMR (300 MHz):
1.11 (br. approx. quint, 2H); 1.28 (br. app. q., 4 H); 1.57 (br. app.
t, 6H); 1.80 (br. approx. d, 4H); 1.97 (br. s, 4H); 3.36 (s, 6H); 3.9
(very br. s, 2H); 9.62 (br. s, 2H). 13C NMR (CDCl3, 75 MHz):
25.7; 26.3; 37.3; 55.6; 161.9. NCH proton of cyclohexyl rings are
1
not observed in H NMR but can be indirectly seen in 13C non-
decoupled NMR providing doublet for NCH carbon of cyclohexyl
ring. Reasons for this behaviour will be reported elsewhere. Calcd.
for C17H30N6O2: C, 58.26; H, 8.63; N, 23.98. Found: C, 58.31; H,
8.70; N, 24.28.
1
evaporated to give the title compound (0.88 g, 61%). H NMR
2,6-Bis[hydroxy(methyl)amino]-4-bis(N-ethoxycarbonylmethyl)-
amino-1,3,5-triazine (2c). Yield 87%. Mp 131–132 ◦C. 1H NMR
(300 MHz): 1.27 (t, J = 7.1 Hz, 6H); 3.34 (s, 6H); 4.20 (q, J =
7.1 Hz, 4H); 4.32 (s, 4H). 13C NMR (CDCl3, 75 MHz): 14.1, 36.8,
49.9, 61.1, 161.8, 163.6, 169.8. Calcd. for C13H22N6O6: C, 43.57;
H, 6.19; N, 23.45. Found: C, 43.52; H, 6.29; N, 23.47.
(CDCl3, 400 MHz): 3.26 (s, 3H); 3.28 (s, 3H); 3.66 (s, 3H), 3.74
(br. s, 8H). 13C NMR (CDCl3, 75 MHz): 35.8; 37.0; 43.6; 61.0;
66.6; 164.2; 166.7; 167.3. Calcd. for C10H18N6O3: C, 44.44; H,
6.71; N, 31.09. Found: C, 44.43; H, 6.46; N, 30.99.
General procedure for preparation of ligands 2d, f. A suspen-
sion of BHT ligands 2c or 2e (0.1 mmol) in ethanol (2 mL) and
2M aq. NaOH (2 mmol, 1 mL) was stirred for 16 h. The reaction
mixture was acidified to pH 3 with 5% aq. HCl and evaporated
to dryness. The residue was dissolved in boiling chloroform, the
chloroform solution was evaporated, the residue is triturated with
ether, and the resultant crystals were isolated by filtration.
2,6-Bis[hydroxy(methyl)amino]-4-(N-methyl, N-ethoxycarbonyl)-
amino-1,3,5-triazine (2e). Yield 86%. Mp 122–124 ◦C. 1H NMR
(300 MHz): 1.27 (t, J = 7.1 Hz, 6H); 3.34 (s, 6H); 3.21 (s, 3H);
3.43 (s, 3H); 3.52 (s, 3H); 4.20 (q, J = 7.1 Hz, 2H); 4.24 (s, 2H). 13
C
NMR (CDCl3, 75 MHz): 14.2, 36.0, 36.8, 51.0, 61.0, 161.7, 163.6,
169.8. Calcd. for C10H18N6O4: C, 41.95; H, 6.34; N, 29.36. Found:
C, 41.72; H, 6.23; N, 29.70.
2,6-Bis[hydroxy(methyl)amino]-4-bis(hydroxycarbonylmethyl)-
amino-1,3,5-triazine (2d). Yield 42%. Mp 210–212 ◦C decomp.
1H NMR (300 MHz, D2O): 3.42 (s, 6H); 4.25 (s, 4H). 13C
NMR (75 MHz, D2O): 37.5, 51.3, 162.7, 167.1, 173.8. Calcd. for
C9H14N6O6·xH2O: C, 33.75; H, 5.04; N, 26.24. Found: C, 33.34;
H, 5.00; N, 26.43.
2,6-Bis(hydroxyamino)-4-dicyclohexylamino-1,3,5-triazine (2g).
◦
1
Yield 80%. Mp 217–218 C decomp. H NMR (300 MHz): 1.1
(br. app. quint, 2H); 1.32 (br. app. q., 4 H); 1.58 (br. app. d., 6H);
1.69 (br. app. d, 2H); 1.85 (app.d, 4H); 2.04 (br. s, 4H); 3.7 (br. s).
13C NMR (CDCl3, 75 MHz): 24.9, 25.7, 29.9, 55.6, 159.5, 161.9.
Calcd. for C15H26N6O2: C, 55.88; H, 8.13; N, 26.08. Found: C,
56.14; H, 8.28.
2,6-Bis[hydroxy(methyl)amino]-4-(N-methyl,
N-hydroxycar-
bony◦lmethyl)amino-1,3,5-triazine (2f). Yield 14%. Mp 178–
1
2,6 - Bis(benzyloxyamino) - 4 -◦dicyclohexylamino - 1,3,5 - triazine
(2h). Yield 68%. Mp 217–218 C. 1H NMR (CDCl3, 400 MHz):
1.1 (br. app. q, 2H); 1.32 (br. app. q., 4 H); 1.58 (br. s, 8H); 1.78 (br.
approx. d, 4H); 2.1 (very br. s, 2H); 4.94(s, 4H); 7.35 (br. app.q,
182 C decomp. H NMR (300 MHz, D2O): 3.22 (s, 3H); 3.38
(s, 3H); 3.47 (s, 3H); 4.35 (s, 2H). 13C NMR (75 MHz, D2O):
35.9, 37.0, 37.1, 51.4, 163.4, 167.9, 173.8. Calcd. HRMS (ESI) for
C8H14N6O4H+: 259.1149. Found: 259.1146.
1292 | Dalton Trans., 2006, 1285–1293
This journal is
The Royal Society of Chemistry 2006
©

View Article Online
J. Tanaka, E. Wada and D. P. Curran, J. Am. Chem. Soc., 1998, 120,
3074–3088.
Acknowledgements
11 For example,J. Y. Chen, O. Ikeda, T. Hatasa, A. Kitajima, M.
Miyake and A. Yamatodani, Electrochem. Commun., 1999, 1, 274–277,
1999.
12 See reviews,(a) Z. D. Liu and R. C. Hider, Coord. Chem. Rev., 2002,
232, 151–171; (b) Z. D. Liu and R. C. Hider, Med. Res. Rev., 2002,
22, 26–64; (c) G. Faa and G. Crisponi, Coord. Chem. Rev., 1999, 184,
291–310; (d) T. B. Chaston and D. R. Richardson, Am. J. Hematol.,
2003, 73, 200–210; (e) D. R. Richardson and P. Ponka, Am. J. Hematol.,
1998, 58, 299–305.
13 (a) U. Heinz, K. Hegetschweiler, P. Acklin, B. Faller, R. Lattmann and
H. P. Schnebli, Angew. Chem., Int. Ed., 1999, 38, 2568–2570; (b) H.
Nick, P. Acklin, R. Lattmann, P. Buchlmayer, S. Hauffe, J. Schupp and
D. Alberti, Curr. Med. Chem., 2003, 10, 1065–1076.
14 H. Keberle, Ann. N.Y. Acad. Sci., 1964, 119, 758.
15 I. Ekeltchik, J. Gun, O. Lev, R. Shelkov and A. Melman, Chem.
Commun., 2005, 5319–5321.
We are grateful to the European Union for partial funding under
the Aquachem network (contract MRTN-CT-2003–503864) and
for the financial support of the ISF-Israel Science Foundation.
We thank Prof. Emilia Kirowa-Eisner from Tel-Aviv University
for the fruitful discussion.
References
1 R. D. Hancock and A. E. Martell, Chem. Rev., 1989, 89, 1875–1914.
2 See reviews (a) L. D. Horwitz and E. A. Rosenthal, Vasc. Med., 1999,
4, 93–99; (b) J. L. Sullivan, J. Lab. Clin. Med., 2004, 144, 280–284.
3 C. R. J. C. Newton, T. T. Hien and N. White, J. Neurol. Neurosurg.
Psychiatry, 2000, 69, 433–441.
4 R. Baliga, N. Ueda, P. D. Walker and S. V. Shah, Drug Metabol. Rev.,
1999, 31, 971–997.
5 E. D. Weinberg, Eur. J. Cancer, 1996, 5, 19–36.
6 (a) E. C. Hirsch and B. A. Faucheux, Mov. Disord., 1998, 13, 39–45
Suppl. 1; (b) W. Linert, E. Herlinger, R. F. Jameson, E. Kienzl, K.
Jellinger and M. B. H. Youdim, Biochim. Biophys. Acta, 1996, 1316,
160–168.
7 M. A. Smith, P. L. R. Harris, L. M. Sayre and G. Perry, Proc. Natl.
Acad. Sci. USA, 1997, 94, 9866–9868.
16 J. T. Shaw, E. R. Nicottra and R. K. Madison, J. Org. Chem., 1962, 27,
4054–4056.
17 (a) P. P. Klemschuk, (J. R. Geigy, A.-G.). Ger. Offen. DE 1910759 (1969)
and US Patent 3867445, (1975).
18 (a) M. Uchida, (Fuji Photo Film Co.). JP 09189973, (1997); (b) Y. Ishii
and M. Taniguchi, (Fuji Photo Film Co.). JP 09304898, (1997).
19 M. Merkor, R. Kissner, R. C. Hider and W. H. Koppenol, Helv. Chim.
Acta, 2004, 87, 3021–3034.
20 R. M. Smith and A. E. Martell, Critical Stability Constants, Plenum
Press, NY, USA, 1989.
21 C. Hershko, A. M. Konijn, H. P. Nick, W. Breuer, Z. I. Cabantchik and
G. Link, Blood, 2001, 97, 1115–1122.
22 http://www.lwr.kth.se/English/OurSoftware/vminteq/.
8 W. Stumm and J. J. Morgan, Aquatic Chemistry, 3rd edn, ch. 10,
pp. 646–648 and ch. 12.5 pp. 744–748, J. Wiley & Sons, NY, 1996.
9 J. Almog and B. Glattstein, J. Forensic Sci., 1997, 42, 993–996.
10 (a) M. Costas, K. Chen and L. Que, Coord. Chem. Rev., 2000, 200, 517–
544; (b) S. Kanemasa, Y. Oderaotoshi, S. Sakaguchi, H. Yamamoto,
This journal is
The Royal Society of Chemistry 2006
Dalton Trans., 2006, 1285–1293 | 1293
©