Synthesis and evaluation of conformationally restricted N 4-tetradentate ligands for implementation in An(III)/Ln(III) separations

Article abstract of DOI:10.1007/s10953-011-9784-1

The previous literature demonstrates that donor atoms softer than oxygen are effective for separating trivalent lanthanides (Ln(III)) from trivalent actinides (An(III)) (Nash, K.L., in: Gschneider, K.A. Jr., et al. (eds.) Handbook on the Physics and Chemistry of Rare Earths, vol. 18-Lanthanides/ Actinides Chemistry, pp. 197-238. Elsevier Science, Amsterdam, 1994). It has also been shown that ligands that "restrict" their donor groups in a favorable geometry, appropriate to the steric demands of the cation, have an increased binding affinity. A series of tetradentate nitrogen containing ligands have been synthesized with increased steric "limits". The pK _a values for these ligands have been determined using potentiometric titration methods and the formation of the colored copper(II) complex has been used as a method to determine ligand partitioning between the organic and aqueous phases. The results for the 2-methylpyridyl-substituted amine ligands are encouraging, but the results for the 2-methylpyridyl-substituted diimines indicate that these ligands are unsuitable for implementation in a solvent extraction system due to hydrolysis. Springer Science+Business Media, LLC 2012.

Full text of DOI:10.1007/s10953-011-9784-1

J Solution Chem (2012) 41:1–16
DOI 10.1007/s10953-011-9784-1
Synthesis and Evaluation of Conformationally Restricted
N4-Tetradentate Ligands for Implementation
in An(III)/Ln(III) Separations
Mark D. Ogden · G. Patrick Meier · Kenneth L. Nash
Received: 10 November 2010 / Accepted: 19 April 2011 / Published online: 25 January 2012
© Springer Science+Business Media, LLC 2012
Abstract The previous literature demonstrates that donor atoms softer than oxygen are ef-
fective for separating trivalent lanthanides (Ln(III)) from trivalent actinides (An(III)) (Nash,
K.L., in: Gschneider, K.A. Jr., et al. (eds.) Handbook on the Physics and Chemistry of Rare
Earths, vol. 18—Lanthanides/Actinides Chemistry, pp. 197–238. Elsevier Science, Ams-
terdam, 1994). It has also been shown that ligands that “restrict” their donor groups in a
favorable geometry, appropriate to the steric demands of the cation, have an increased bind-
ing affinity. A series of tetradentate nitrogen containing ligands have been synthesized with
increased steric “limits”. The pKa values for these ligands have been determined using po-
tentiometric titration methods and the formation of the colored copper(II) complex has been
used as a method to determine ligand partitioning between the organic and aqueous phases.
The results for the 2-methylpyridyl-substituted amine ligands are encouraging, but the re-
sults for the 2-methylpyridyl-substituted diimines indicate that these ligands are unsuitable
for implementation in a solvent extraction system due to hydrolysis.
Keywords Copper(II) · Stability constants · Acid dissociation constants · Phase
distribution · Ligand preorganization · Nitrogen donor complexants
1 Introduction
The separation of lanthanides from transplutonium actinides is critical to the efficient man-
agement of spent nuclear fuel. Transplutonium actinides are found predominantly in the
M.D. Ogden · G.P. Meier · K.L. Nash (
)
ꢀ
Department of Chemistry, Washington State University, Pullman, WA 99164-4630, USA
e-mail: knash@wsu.edu
M.D. Ogden (
)
ꢀ
Idaho National Laboratory, P.O. Box 1625, Idaho Falls, ID 83514-6152, USA
e-mail: mark.ogden@inl.gov
Present address:
M.D. Ogden
ANSTO Minerals, Locked Bag 2001, Kirrawee DC, NSW 2232, Australia
e-mail: mark.ogden@ansto.gov.au

2
J Solution Chem (2012) 41:1–16
+3 oxidation state, as are the lanthanides. Therefore, the separation of these elements must
exploit the small difference between the solution chemistry of lanthanides and transpluto-
nium actinides in the trivalent oxidation state [1].
Actinides bind the softer bases such as N, S, and Cl− more strongly than the trivalent
lanthanides. The difference in binding affinity can be exploited for successful separations.
The 2-methylpyridyl structural analogue of EDTA, N,N,Nꢀ,Nꢀ-tetrakis(2-pyridylmethyl)-
ethylenediamine (TPEN), shows a preference for An3+ over Ln3+ [2]. The stability constant
for the americium (Am) TPEN complex is 100 times greater than for the samarium (Sm)
complex, resulting in an overall difference in the ꢀG for complexation of 11.6 kJ·mol−1 [2].
The effect of donor group preorganization is clearly seen in the difference between the
formation constants for a 1:1 complex with Am(III) of N,Nꢀ-dimethyl bicyclic diamide
(DMBCDA) and tetramethylmalonamide (TMMA) at 0.1 mol·L−1 HNO3 and 0.9 mol·L−1
NaNO3 is ꢀlog10 β = 1.98, which represents a difference in the Gibbs energy of formation
between TMMA and DMBCDA of −11.3 kJ·mol−1 [3].
The preorganizing effect of a cyclohexane backbone can be seen in results of calorimet-
ric studies on the complexation of Am(III) and Cm(III) with polyaminocarboxylate ligands
[4]. The amine groups in the hexadentate ligand trans-1,2-cyclohexanedinitrilotetraacetic
acid (CDTA) are prearranged trans to each other in the 1,2 positions on a cyclohexane ring;
the trans- arrangement on the cyclohexane ring eliminates free rotation which can occur
around the ethylene bridge joining the iminodiacetate groups in ethylendiaminetetraacetic
acid (EDTA). In the Cm3+ data, the CDTA stability constant is larger than that for EDTA
by ꢀlog10 K101 = 1.23, representing ꢀG being more favorable by 7.1 kJ·mol−1 [4]. The
improved ꢀG is opposed by an endothermic contribution to ꢀH that is overcome by in-
creased ꢀS. This difference in T ꢀS results from the CDTA not having to rearrange to any
great degree before complexing the Cm3+; changing the hydration sphere around the metal
center also contributes to this increased entropic effect [4, 5].
To a first approximation, a significant portion of the entropy difference in the formation of
metal complexes must be attributed to molecular rearrangement of the ligand. To rearrange
EDTA into an appropriate configuration for binding a metal center, the molecule must, in an
acidic solution (both nitrogen atoms are protonated), reconfigure from a conformer in which
the nitrogens are as far apart as possible (dihedral angle approx. 180°) to one in which the
nitrogens are in close proximity to one another. For CDTA this molecular rearrangement
is constrained due to the cyclohexane ring hindering the rotation around the C–C bond of
the C’s bonded to nitrogen atoms. Cleavage of a ring C–C bond would have to occur to
allow free rotation. This high energy process hinders the rotation. The trans orientation of
the N donor atoms is adopted to minimize steric repulsions between the substituents on the
ring (dihedral angle approx. 60°).
The energy difference required for a separation factor of 100 in a solvent extraction sys-
tem is approximately 11.4 kJ·mol−1, which is comparable to the energy of a single hydrogen
bond [1]. Even though “soft” donor-containing ligands in general have an inherent selectiv-
ity for An(III) over Ln(III), the binding constants in general are low. To increase the binding
constant of the complexant for metal ions, the ligand can be structurally preorganized or
preoriented. In this investigation, ligands containing only nitrogen donor atoms have been
synthesized. The impact of structural prearrangements of the donor atoms is to be evaluated.
The ligands synthesized for this study and their acronyms are presented in Fig. 1. The
sterically restricted 2-methylpyridyl ligands synthesized in this study are N,Nꢀ-bis(2-meth-
ylpyridyl)-1,2-diaminoethane (BPMDAE), N,Nꢀ-bis-[pyridin-2-ylmethylene]-ethane-1,2-
diamine (BPMDIE), N,Nꢀ-bis(2-pyridylmethyl)piperazine (BPMPIP), N,Nꢀ-bis(2-meth-
ylpyridyl)-1,3-diaminopropane (BPMDAP), trans-N,Nꢀ-bis-[pyridin-2-ylmethylene]cyclo-

J Solution Chem (2012) 41:1–16
3
Fig. 1 Conformationally restricted 2-methylpyridyl ligands synthesized in this study: BPMDIE =
ꢀ
ꢀ
N,N -bis-[pyridin-2-ylmethylene]ethane-1,2-diamine; BPMDAE = N,N -bis(2-methylpyridyl)-1,2-diami-
ꢀ
ꢀ
noethane; BPMPIP = N,N -bis(2-pyridylmethyl)piperazine; BPMDAP = N,N -bis(2-methylpyridyl)-1,3-di-
ꢀ
aminopropane; BPMDIC = trans-N,N -bis-[pyridin-2-ylmethylene]-cyclohexane-1,2-diamine; BPMDAC =
ꢀ
trans-N,N -bis(2-pyridylmethyl)-1,2-diaminocyclohexane
hexane-1,2-diamine (BPMDIC), and trans-N,Nꢀ-bis(2-pyridylmethyl)-1,2-diaminocyclo-
hexane (BPMDAC).
The acid dissociation constants for the ligands are reported for the first time at I =
0.5 mol·L−1, along with their copper complex stability constants and partitioning constants
for the resulting complexes between aqueous phase (I = 0.5 mol·L−1) and 1-octanol as a
function of pH. The screening of these compounds using Cu(II) was initially pursued to
explore the aqueous stability of these ligands as well as investigating using the Cu(II) chro-
mophore formation as an indicator to explore ligand phase distribution as a function of pH.
The stability of the ligands in solution and how they distribute between an organic and
aqueous phase will be used to screen these ligands as to how effectively they may be for
implementation in an Am(III) and Ln(III) separation.
The variation in the arrangement of nitrogen donor atoms in the ligands addresses two
important issues, the effect of prearrangement of ligand donor atoms and the impact of con-
verting imine into amine groups. The backbone of the ligand has been changed, in order to
preorganize the amine or imine donor groups into configurations that maximize metal inter-
actions, from the flexible alkyl backbones of BPMDAE, BPMDIE, BPMDAP and BPMDIE,
to the trans-arrangement using a cyclohexane backbone (BPMDIC and BPMDAC), to the
more rigid boat/chair cyclohexane backbone of BPMPIP. The change from diimine to di-
amine decreases the planarity of the 2-methyl pendant arms and was hypothesized to de-
crease the possibility of metal–ligand interaction. In the diimine bonds the C and N atoms
are sp2 hybridized and conjugated to the pyridyl ring, giving the pendant arm-pyridine por-
tion of the ligand a planar configuration that does not readily undergo free rotation. The
effect of chelate ring size can be investigated in the chemistry of BPMDAP when compared
with BPMDAE. The binding of the backbone of BPMDAP will form a 6 membered chelate
ring with a metal center, whereas BPMDAE will form a 5 membered chelate ring. In general
it is considered that with large metal ions, 5 membered chelate rings are more stable than 6.
2 Experimental
2.1 Ligand Synthesis
The ligands were synthesized following modifications of established synthetic procedures
(BPMDAE, BPMDIE, and BPMDAP [6], BPMDIC [7], and BPMDAC [8]). In general, the

4
J Solution Chem (2012) 41:1–16
diimines were synthesized by condensation of 2-pyridinecarboxaldehyde with the appropri-
ate diamine in methanol over 5 Å molecular sieves. The addition of 5 Å molecular sieves
only marginally increased yields above those of previous syntheses. The tetradentate di-
amines were then prepared by reduction of the diimine with sodium borohydride (NaBH4)
in ice cold methanol (MeOH). After purification by solvent extraction, the tetradentate di-
amines were precipitated as the HCl salts from diethyl ether. Before precipitation, the struc-
ture and purity of the final compounds were confirmed by comparison of the NMR spectra
with literature data [6–9]. BPMPIP was synthesized by the reaction of piperazine with a
slight stoichiometric excess of 2-chloromethylpyridine in the presence of triethylamine in
acetonitrile following an established procedure [9].
2.1.1 Synthesis of N,Nꢀ-bis-[Pyridin-2-ylmethylene]ethane-1,2-diamine (BPMDIE)
2-Pyridinecarboxaldehyde (6.492 g, 60 mmol) was added dropwise to a solution of 1,2-di-
aminoethane (1.803 g, 30 mmol) in anhydrous methanol (MeOH, 50 mL) with stirring under
a nitrogen balloon. Activated 5 Å molecular sieves were added to this reaction mixture and
the reaction allowed to stand for 2 h at room temperature. The reaction mixture was then
refluxed for 3 h. The solution was filtered and solvent evaporated to yield crude product.
The orange product was recrystallized from hot MeOH and subsequently from hot hexanes.
The product was isolated as a yellow crystalline material (60.8% yield). 1H NMR (CDCl3):
δ 4.07 (s, 4H), 7.31 (m, 2H), 7.73 (m, 2H), 8.42 (s, 2H), 8.62 (m, 2H). 13C NMR (CDCl3):
δ 61.56, 121.58, 125.00, 149.63, 154.58, 163.67.
2.1.2 Synthesis of N,Nꢀ-bis(2-Pyridylmethyl)ethane-1,2-diamine (BPMDAE)
Sodium borohydride (NaBH4) (1.579 g, 41.75 mmol) was added slowly with stirring over 2 h
to a solution of BPMDIE (B) (4.003 g, 16.7 mmol) in anhydrous MeOH (100 mL), under a
nitrogen atmosphere and held at 0 °C using an ice bath. After the addition was completed, the
reaction mixture was refluxed for 15–20 minutes, and then left to stir at room temperature
overnight. Sodium hydroxide (1.0 mol·L−1, 150 mL) was added to the reaction mixture
and the compound was then extracted into methylene chloride (CH2Cl2) (4 × 15 mL). The
CH2Cl2 solution was dried using anhydrous potassium carbonate (K2CO3), filtered, and
the solvent evaporated to leave behind a light yellow oil (96.4% yield). The light yellow
oil was taken up in 25 mL of diethyl ether and the ligand precipitated with 3 equivalents
of HCl (48.6 mL, 1.0 mol·L−1 HCl in diethyl ether). The precipitate was washed (diethyl
ether) and dried by multiple evaporations of diethyl ether on a rotary evaporator (care must
be taken as the compound is very hydroscopic before the excess HCl is removed). 1H NMR
(CDCl3): δ 2.71 (s, 4H), 3.80 (s, 4H), 7.02 (m, 2H), 7.20 (m, 2H), 7.50 (m, 2H), 8.42 (m, 2H).
13C NMR (CDCl3): δ 49.13, 55.21, 122.20, 122.58, 136.77, 149.27, 159.66.
2.1.3 Synthesis of N,Nꢀ-bis(2-Pyridylmethyl)piperazine (BPMPIP)
A solution of piperazine (1.010 g, 11.73 mmol) in CH3CN (40 mL) was treated with 2-pic-
olyl chloride hydrochloride (3.943 g, 24.04 mmol) and triethylamine (6.5 mL). The resultant
mixture was stirred in a loosely sealed flask at room temperature. After 48 h, the mixture
was poured into 1 mol·L−1 NaOH (100 mL) and extracted with CH2Cl2 (3 × 50 mL). The
combined extracts were dried over anhydrous K2CO3, filtered, then the solvent was evapo-
rated to yield the crude product as a brown solid. The solid was dissolved in diethyl ether
(50 mL) and heated to reflux, resulting in the deposition of an insoluble brown oil. The ether

J Solution Chem (2012) 41:1–16
5
solution decanted from the oil was evaporated to yield the pure product as a very light yel-
1
low crystalline solid (45.7% yield). H NMR (CDCl3): δ 2.58 (s, 8H), 3.68 (s, 4H), 7.18
(m, 2H), 7.39 (m, 2H), 7.64 (m, 2H), 8.56 (m, 2H). 13C NMR (CDCl3): δ 39.75, 122.45,
126.45, 137.53, 148.39, 149.90, 165.26.
2.1.4 Synthesis of N,Nꢀ-bis(2-Methylpyridyl)-1,3-diaminopropane (BPMDAP)
2-Pyridinecarboxaldehyde (6.492 g, 60 mmol) was added dropwise to a solution of 1,3-
diaminopropane (2.224 g, 30 mmol) in anhydrous MeOH (50 mL) with stirring under a
nitrogen atmosphere. To this reaction mixture, activated 5 Å molecular sieves were added
and the reaction allowed to stand for 2 h at room temperature after which time the solution
was refluxed for 3 h. The reaction was filtered and the solvent evaporated to yield the product
as a yellow oil (84.6% yield). The oil product (6.406 g, 16.7 mmol) was redissolved in
anhydrous MeOH (100 mL), held at 0 °C using an ice bath under a nitrogen atmosphere, and
NaBH4 (2.401 g, 63.47 mmol) was added slowly with stirring over 2 h. The reaction mixture
was refluxed for 20 minutes and left to stir at room temperature overnight. To this solution
1.0 mol·L−1 NaOH (150 mL) was added and the product was then extracted with CH2Cl2
(4 × 15 mL). The CH2Cl2 solution was dried using anhydrous K2CO3, filtered, and the
solvent evaporated leaving a light yellow oil (96.5% yield). The light yellow oil was taken
up in 25 mL of diethyl ether and the tetradentate diamine precipitated with 3 equivalents of
HCl (72.3 mL of 1.0 mol·L−1 HCl in diethyl ether). The precipitate was washed with diethyl
ether and dried by multiple evaporations of diethyl ether on a rotary evaporator (care must
1
be taken as the compound is very hydroscopic before excess HCl is removed). H NMR
(D2O): δ 2.15 (m, 2H), 3.20 (t, 4H), 4.55 (s, 4H), 7.84 (m, 2H), 7.95 (m, 2H), 8.28 (m, 2H),
8.65 (m, 2H). 13C NMR (CDCl3): δ 22.83, 45.26, 48.35, 127.67, 127.95, 144.42, 145.77,
146.37.
2.1.5 Synthesis of trans-N,Nꢀ-bis-[Pyridin-2-ylmethylene]-cyclohexane-1,2-diamine
(BPMDIC)
2-Pyridinecarboxaldehyde (6.492 g, 60 mmol) was added drop wise to a solution of trans-
1,2-diaminocyclohexane (3.426 g, 30 mmol) following the same method used for making
BPMDIE. The crude yellow product was recrystallized from hot methanol. The product was
isolated as a yellow crystalline material (80.9% yield). 1H NMR (CDCl3): δ 1.50 (m, 4H),
1.85 (m, 4H), 3.52 (m, 2H), 7.20 (m, 2H), 7.62 (m, 2H), 7.86 (m, 2H), 8.30 (s, 2H), 8.53
(m, 2H). 13C NMR (CDCl3): δ 24.56, 32.93, 73.77, 121.57, 124.70, 136.67, 149.47, 154.80,
161.66.
2.1.6 Synthesis of trans-N,Nꢀ-bis(2-Pyridylmethyl)-1,2-diaminocyclohexane (BPMDAC)
A solution of BPMDIC (4.065 g, 13.9 mmol) in anhydrous MeOH (100 mL) was cooled
to 0 °C using an ice bath, NaBH4 (1.461 g, 34.76 mmol) was added slowly with stirring
over 2 h. The solution was then worked up following the previous method for the synthesis
of BPMDAE resulting in a light yellow crystalline material (98.4% yield). The crystals
were taken up in 25 mL of diethyl ether and the diamine precipitated with 3 equivalents of
HCl (42.0 mL of 1.0 mol·L−1 HCl in diethyl ether). 1H NMR (D2O): δ 1.50 (m, 4H), 1.17
(m, 4H), 1.67 (m, 2H), 2.17 (m, 2H), 2.85 (m, 2H), 4.20 (d, 2H), 4.40 (d, 2H) 7.62 (m, 2H),
7.64 (d, 2H), 8.16 (m, 2H), 8.43 (m, 2H). 13C NMR (CDCl3): δ 25.43, 32.54, 53.40, 63.27,
123.95, 124.70, 137.84, 149.58, 161.68.

6
J Solution Chem (2012) 41:1–16
2.2 Potentiometric Measurements of Acid Dissociation Constants
All measurements were made using a Mettler Toledo DL50 Autotitrator fitted with an Orion
Ross ultra semi-micro glass electrode with 3 mol·L−1 NaCl as the filling solution. The ex-
perimental and titration solutions were made up to 0.5 mol·L−1 ionic strength (NaClO4),
temperature was maintained at 25 °C and a slow stream of nitrogen was passed through the
titration vessel to reduce the interference from CO2. All titration were corrected (pCH) back
to theoretical pH after the titration of standardized strong acid (0.01 mol·L−1 HClO4) and
strong base (0.01 mol·L−1 NaOH).
All the ligand solutions were titrated with standardized 0.01 mol·L−1 NaOH. Solutions
of BPMDIE, BPMPIP, and BPMDIC (1 × 10−3 mol·L−1) were prepared by dissolving the
crystalline material in 18 Mꢁ water with an excess of acid (8 equivalents 0.01 mol·L−1
HClO4) resulting in a solution of approximately pH = 2. Solutions of BPMDAE, BPMDAP,
and BPMDAC (1 × 10−3 mol·L−1) were prepared from the neutralized HCl salt. The HCl
salt was weighed out and the combined with 1.0 mol·L−1 NaOH (20 mL). The free amine
was then solvent extracted into CH2Cl2 (5 × 5 mL) and transferred to a weighed round
bottom flask. The solvent was evaporated and the mass of the free amine determined. The
stock solution of the free amine for potentiometric titration was made up following a similar
procedure to that described for the imines. All potentiometric titrations were performed
in triplicate, in both static and dynamic modes. The potentiometric curves were fit using
PSEQUAD [10] and the most acidic pKa fit by minimization of the standard error in the
data.
2.3 Spectrophotometric Measurements and Determination of Stability Constants
Ligand–metal complex solutions were made using Cu(ClO4)2·6H2O, 99.99% purity
(0.001 mol·L−1) and an excess of the neutralized ligand (0.01 mol·L−1). The solution was
acidified using 0.01 mol·L−1 HClO4 (approximately pH = 2) and NaClO4 was added to
maintain the ionic strength at 0.5 mol·L−1. The ligand–metal complex solution was then
titrated with standardized 0.1 mol·L−1 NaOH. The absorption spectra of the copper com-
plexes with the ligands (structures shown in Fig. 1) were recorded as a function of pH at
23 ( 2) °C in the range of 350–950 nm using a Varian Cary 50 Bio UV–vis or Cary 13
UV–vis spectrophotometer. For the ligands that underwent hydrolysis, SQUAD fitting was
only carried out on the spectra of the hydrolyzed Cu(II) species.
A different methodology was developed to determine the stability constants of ligands
that have copper complexes forming at pH = 1 (BPMDAE, BPMDAP, and BPMDAC).
At a constant acid concentration (0.25 mol·L−1), ionic strength (I = 0.5 mol·L−1 us-
ing NaClO4), and constant copper (0.001 mol·L−1), the metal:ligand ratio was changed
(0.0004–0.004 mol·L−1 ligand concentration). The resulting spectra for the spectrophoto-
metric titrations were fit using SQUAD [11] and the log10 K101 value determined.
2.4 Phase Distribution Study
Partitioning constants for ligand distribution between n-octanol and water were determined
as follows: 1 cm3 of octanol was spiked with 50 µL of a methanolic solution containing
4 × 10−2 mol·L−1 of ligand creating a 5% MeOH/octanol (v/v) phase. Th−is1organic solution
was equilibrated with 1 cm3 of aqueous solution containing 0.49 mol·L of NaClO4 and
1 × 10−2 mol·L−1 of a buffer (dichloroacetic acid, chloroacetic acid, acetic acid, metanilic
acid, 2-(N-morpholino)ethanesulfonic acid (MES) or N-(2-hydroxyethyl)piperazine-Nꢀ-2-
ethanesulfonic acid (HEPES)) by shaking for 12 h. After centrifugation the aqueous phase

J Solution Chem (2012) 41:1–16
7
−1
Fig. 2 Potentiometric titrations of fully protonated N -tetradentate ligands with NaOH at 0.5 mol·L
4
NaClO . PSEQUAD fit shown by solid grey line: (A) E, BPMDAE; !, BPMDAP; ×, BPMDAC; (B) P,
4
BPMPIP; ∗, BPMDIE; 1, BPMDIC
was separated and a 400 µL sample taken and 400 µL of indicator solution added. The indica-
tor solution consisted of excess copper (0.006 mol·L−1) at an ionic strength of 0.5 mol·L−1
using NaClO4, at pH = 5.5 set by sodium acetate buffer (approximately 0.4 mol·L−1) made
up in a 50% (v/v) methanol/water solution. Using the experimentally measured extinction
coefficients, the aqueous concentration of ligand was determined and the concentration of
ligand in the organic phase was determined assuming mass balance. All phase distribution
samples were performed in triplicate and the pH measured before and after contact. Ligand
concentration of the stock solution was determined spectrophotometrically using the forma-
tion of the copper complex chromophore by addition of a copper containing solution, as an
indicator, buffered at pH = 5.5 using (0.2 mol·L−1) sodium acetate containing 50/50 (v/v)
H2O/MeOH and (0.008 mol·L−1) Cu(ClO4)2 at 0.5 mol·L−1 ionic strength (NaClO4).
3 Results and Discussion
3.1 Determination of Acid Dissociation Constants
The results of potentiometric titrations of the fully protonated forms of the ligands with
NaOH are shown in Fig. 2. The protonation constants for BPMDAE and BPMDAP have
been determined previously at 0.1 mol·L−1 ionic strength [12], but acid dissociation con-
stants for all these ligands have not been determined previously at I = 0.5 mol·L−1. The
resolved acid dissociation constants determined by PSEQUAD using non-negative least-
square analysis (NNLS) are given in Table 1. The PSEQUAD fit for each potentiometric
ligand titration is shown by the solid dark grey line in each of the titration runs (Fig. 2).
Modeling of the aqueous speciation of BPMDAE, BPMDAP, BPMDAC and BPMPIP were
carried out using the hyperquad simulation program Hyss2009 and the resulting diagrams
are shown in Fig. 3.
The error in pK1 for BPMDAE, BPMDAP, BPMPIP and BPMDAC are considerably
large. This error is due to the fact the pKa is not directly measurable under these exper-
imental conditions and the values were thus determined by minimization of the standard

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J Solution Chem (2012) 41:1–16
Table 1 Acid dissociation constants for 2-methylpyridyl based ligands at 25 °C
Compound BPMDAE BPMDAP BPMDAC BPMPIP BPMDIE
BPMDIC
pK
pK
pK
pK
8.65 ( 0.02) 9.29 ( 0.02) 9.25 ( 0.02) 7.26 ( 0.01) 9.09 ( 0.01) 9.39 ( 0.01)
5.79 ( 0.01) 7.31 ( 0.01) 5.55 ( 0.01) 5.12 ( 0.01) 7.38 ( 0.02) 6.33 ( 0.01)
2.14 ( 0.03) 2.29 ( 0.02) 1.55 ( 0.08) 2.57 ( 0.03) 4.80 ( 0.02) 4.35 ( 0.01)
4
3
2
1
1.6 ( 0.7)
1.2 ( 0.7)
1.0 ( 0.7)
1.8 ( 0.7)
3.84 ( 0.02) 3.78 ( 0.01)
Fig.
3 Speciation diagrams for (A) BPMDAE; (B) BPMDAP; (C) BPMDAC; (D) BPMPIP at
−1
−1
0.001 mol·L ligand concentration, 25 °C and 0.5 mol·L ionic strength (NaClO ). In all speciation dia-
4
4+
3+
2+
2
+
grams: H L , black dashed line; H L , grey solid line; H L , grey dotted line; HL , grey dashed line,
4
3
L
, black solid line
free
deviation in the pK1 value. The lowest two pKa values for BPMDAE, BPMDAP, BPMPIP,
and BPMDAC (shown in Table 1) indicate the more acidic nature of the pyridinium moieties
when compared with the completely aliphatic analogue triethylenetetramine (trien) [13].
The higher pKa values for BPMDIE and BPMDIC appear counterintuitive based on the
effect of conjugation on the overall acidity of a ligand. The pK1 values for both BPMDIE
and BPMDIC match that of 2-pyridine carboxaldehyde at 3.8 [14] and indicate that both of
these ligands undergo hydrolysis in aqueous solution to produce a complex mixture of both
the synthetic precursors pyridine carboxaldehyde and the bridging amine, as well as the
mono-substituted and bis-substituted pyridylmethylene amine. Based on this data BPMDIE
and BPMDIC are not suitable for use in a solvent extraction system.
The values for BPMDAE are in good agreement with the earlier literature, in which the
pKa values are reported at 0.1 mol·L−1 ionic strength as 1.62, 1.81, 5.45, and 8.23, respec-

J Solution Chem (2012) 41:1–16
9
2+
Table 2 Absorption data for the complexes of with Cu
Compound
λ
Molar absorptivity ε
(L·mol ·cm
Lowest observable conc.
max
−1
−1
−1
)
(nm)
of ligand (mol·L
)
−5
BPMDAE
BPMDAP
BPMDAC
BPMPIP
606
598
589
650
606
606
149 ( 6)
158 ( 3)
188 ( 4)
288 ( 8)
71 ( 3)
1.34×10
−5
1.18×10
−5
0.80×10
−5
0.69×10
−5
BPMDIC
BPMDIE
2.61×10
−5
156 ( 6)
1.28×10
tively [12]. The increase in pKa values for BPMDAE is due to the increase in ionic strength
(0.5 versus 0.1 mol·L−1) [13]. In previous studies it has been asserted that the relative affinity
of the ligand BPMDAE for protons does not correlate with its affinity for selected divalent
metal ions [12]. The pKa values of these ligands may not reflect their distribution between
an organic and aqueous phase as this is also a function of the lipophilicity of the amine
backbone.
The pKa values for BPMDAP are similar to the values reported in the NIST database at
1.0 mol·L−1 ionic strength [13], which are reported as 8.99, 7.28, 2.34 and 1.3, respectively.
We do see that pK3 and pK4 are more acidic at lower ionic strength, as would be expected,
at 2.29 and 1.2. The pK3 of BPMDAP at 7.31 is within the literature error of 7.28; the values
for pK3 and pK4 at 7.31 and 9.29 do not follow the expected trend of increased acidity with
increasing ionic strength and pK4 is more basic than expected. The effect could be due to
a degree of impurity in the prepared BPMDAP compound (residual presence of unreacted
1,3-diaminopropane) as this was the only ligand made via a one-pot synthesis methodology.
ꢀ
Overall the trend in the total basicity ( pKa) for these ligands is a follows:
BPMPIP < BPMDAC < BPMDAE < BPMDAP < BPMDIC < BPMDIE
3.2 Determination of Copper–Ligand Stability Constants
The spectral details for the ligands screened with copper are given in Table 2. The lowest
observable amount of ligand (as the Cu2+ complex) is calculated from the observed spectra
of the 1:1 complex greater than 5σ (limit of quantification) of the baseline at the λmax. It is
noteworthy that the Cu-BPMPIP complex has a molar absorptivity nearly double the values
of the other complexes. This is probably due to the distortion of the metal center from ideal
symmetry when complexed to BPMPIP ligand [9].
In all of the UV–vis spectrophotometric pH titrations, the growth of a single peak is ob-
served with pH increase or increased ligand concentration. In Fig. 4 the spectrophotometric
titrations under high acidic conditions are shown (BPMDAE, BPMDAP, and BPMDAC).
These experiments were carried out as metal–ligand titrations at high acidity so that pro-
tons could effectively compete with the strength of the metal ligand complex. In Fig. 5 the
spectrophotometric titrations of the diimine based ligands and BPMPIP are presented. These
experiments were carried out as complex pH titrations from pH = 2.00 to neutral conditions
to prevent precipitation of hydrolyzed species. The stoichiometry of the copper species was
determined for all ligands as 1:1 by using Job plot analysis.
The calculated stability constants for the Cu(II) complexes at 0.5 mol·L−1 ionic strength
and 21 °C are given in Table 3. The stability constants were calculated using SQUAD from

10
J Solution Chem (2012) 41:1–16
−1
−1
Fig. 4 Spectrophotometric titrations of 0.001 mol·L
Cu(ClO ) , 0.15–0.25 mol·L
HClO , I =
4 2
4
−1
0.5 mol·L
(NaClO ) and 21 °C, by incremental addition of ligand, black line—M:L ratio 1:0, grey
line—M:L ratio 1:4; (A) BPMDAE, (B) BPMDAP, and (C) BPMDAC with Cu(II)
4
2+
Table 3 Stability constants for the complexation of Cu by N -tetradentate ligands determined at 21 °C
4
−1
and I = 0.5 mol·L (NaClO )
4
−1
Compound
log
K
ꢀG (kJ·mol
−99
)
101
10
BPMDAE
BPMDAP
BPMDAC
BPMPIP
17.7 0.9
5
a
a
18.9 0.2 (18.7 0.2 )
19.5 0.6
−106 1 (−105 1 )
−109
3
7.41 0.06
−41.7 0.3
−66.1 0.1
−79.6 0.1
BPMDIE
BPMDIC
11.73 0.02
14.13 0.03
a
−1
ionic strength from the NIST data-
Data fitted using selected pK values for BPMDAP at 1.0 mol·L
a
base [13]
the changes in the absorption spectra as a function of changing ligand concentration or
changing pH. The high errors associated with fitting the amine based ligands (BPMDAE,
BPMDAP, and BPMDAC) is predominantly due to the uncertainty in pK1 when operat-
ing at high acidic conditions. The stability constant values determined for BPMDIE and
BPMDIC are of questionable value due to the fact that these ligands undergo rapid hydrol-
ysis in solution and the nature of these mixtures is highly uncertain.
The complex formed between BPMDAE and Cu2+ is strong, with log10 K101 = 16.3 [12]
at 0.1 mol·L−1 ionic strength. The Cu2+ complex with BPMDAP is approximately 100 times
stronger (log10 K101 = 18.9). This difference in the strength of these complexes is consistent
with previous reports, as the steric demands imposed by smaller Cu2+ cation lead to a pref-
erence for six-membered chelate rings (over five-membered rings) due to the reduced steric
strain on the ligand backbone with the longer propyl spacer of BPMDAP [18].
At 0.5 mol·L−1 ionic strength we see an increase in the Cu(II) log10 K101 for BPMDAE
and BPMDAP (17.7 and 18.9 respectively) when compared to those obtained at 0.1 mol·L−1
as is seen for triethylenetetramine (trien) [13]. Using the pKa values determined for BPM-
DAP (Table 1) in fitting the Cu(II) log10 K101, we get an increased value more typically
expected for an increase in ionic strength. Using values for pK1 and pK2 between the val-
ues available in the NIST database for 0.1 and 1.0 mol·L−1 ionic strength and the pK3 and
pK4 for 1.0 mol·L−1 ionic strength in the SQUAD fitting gives a predicted log10 K101 value
approximately equal to the literature value of 18.35 at 0.1 mol·L−1 ionic strength for cop-

J Solution Chem (2012) 41:1–16
11
Table 4 Cu(II) stability
a
Ligand
Cu(II) log
K
101
10
−1
constants at 0.1 mol·L ionic
strength with selected ligands
[13]
PipDA
EDDA
7.37
16.2
Ethylenediamine
10.49
11.09
16.3
a
PipDA =
trans-Cyclohexane-1,2-diamine
BPMDAE
piperazine-1,4-diacetic acid;
EDDA =
ethylenedinitrilodiacetic acid
per complexation (log10 K101 = 18.4 0.2). Fitting the data using the 1.0 mol·L−1 ionic
strength pKa values gives a result of log10 K101 = 18.7 0.2, which is nearer to the ex-
pected value. The highest Cu(II) log10 K101 value (19.5 0.6) for all the ligands tested is
seen for the Cu2+ complex with BPMDAC and is evidence of the effect of constraining
the N-donor groups using a cyclohexane backbone when compared with BPMDAE. A par-
allel comparison can be made for the complexes of Cu2+ with ethylene diamine (en) and
trans-cyclohexane-1,2-diamine (trans-cn) [13] (Table 4).
Results from the potentiometric titration clearly indicate that BPMDIE and BPMDIC
undergo rapid hydrolysis and since the nature of these hydrolysis mixtures is highly uncer-
tain, the numerical value for any equilibrium constant obtained from them is of question-
able value. One clear result from the spectrophotometric titration of BPMDIC with Cu2+
is the indication that the ligand is more stable to hydrolysis than the BPMDIE. At the start
of the titration, at acidic pHs, we see a λmax at 634 nm which may be indicative of the
Cu(BPMDIC)2+ complex and, as the titration continues, hydrolysis of the ligand occurs and
λmax shifts down to 606 nm, similar to the λmax observed in the BPMDIE and BPMDAE
complexes.
The log10 K101 value for BPMPIP is approximately 10 orders of magnitude smaller than
the value for BPMDAE (Table 3), indicating that protons successfully compete with the
metal at much lower concentrations; these complexes become more important at pH val-
ues more basic than pH = 2.5. This energy deficit is interpreted as indicating that the free
BPMPIP ligand, rather than being preoriented for cation binding, probably maintains the
backbone piperazine ring in the chair configuration. For tetradentate bonding of the cation,
the piperazine ring would have to “flip” into the boat configuration. The energy penalty as-
sociated with this conformational reorganization would be considerable. A parallel quantita-
tion of this effect can be seen when comparing the binding of copper by ethylenediaminedi-
acetic acid (EDDA) and piperazine-1,4-diacetic acid (PipDA) [13] (Table 4). Similarly, the
energy required to flip the piperazine from the boat to the chair conformer in the tetradentate
nitrogen donor ligand 1,4-bis(2-aminoethyl)-1,4-diazacyclohexane (BAE-P) has been calcu-
lated as approximately 29.3 kJ·mol−1, which would equate to a difference in log10 K101 of
5.20 at 21 °C [14]. This is approximately half of the ꢀlog10 K101 value between BPMDAE
and BPMPIP (ꢀlog10 K101 = 10.3). It is clear that flipping the piperazine ring accounts
for only part of the energy difference in the binding comparison. The rest of the energy
difference could be attributed to steric interactions such as those between the methylene hy-
drogens on the pendant arm (N-CH2-Pyr) and the hydrogens meta to the pyridyl nitrogen,
and interactions of the piperazine backbone with other species in the inner coordination
sphere. Taking into account the effect of the polarity and size of the pyridyl ring could also
be a factor, but this is probably not the case based on the comparison of the complexation of
EDDA and BPMDAE with Cu(II) (Table 4).
The crystal structure of the complex [Cu(BPMPIP)(NO3)]NO3 is a 5 coordinate square
pyramidal complex with the ligand occupying the equatorial plane with the piperazine in

12
J Solution Chem (2012) 41:1–16
−1
−1
Fig. 5 Spectrophotometric pH titrations of ligands with 0.001 mol·L Cu(ClO ) , 0.01 mol·L ligand,
4 2
−1
I = 0.5 mol·L
and 21 °C; (A) BPMPIP—black line starting p H = 2.34, dark grey line is final pH of
C
7.45; (B) BPMDIE—black line starting p H 2.19, dark grey line is final pH of 6.73; (C) BPMDIC—black
C
line starting p H = 2.26, dark grey line is final pH of 7.23
C
the boat configuration [15]. A counter ion is held within the inner coordinating sphere in an
axial position and the final counter anion is in the second coordination sphere with respect
to the nitrate (no distances are given for the second nitrate) [15]. The steric distortion due
to the interaction of the piperazine backbone with species coordinated in the axial positions
pushes the second anion into the second coordination sphere. A steric interaction of this sort
would create a further distortion away from ideal symmetry and thus lead to a much higher
molar absorptivity (Table 2).
3.3 Phase Distribution Study
The distribution ratio, DL, of a ligand between an organic phase and an aqueous phase
is given by the partitioning coefficient, KD:
[L]org
KD =
[L]free
The expression for DL, which is the ratio of the concentration of the ligand in the organic
phase to the total concentration of all ligand containing species in the aqueous phase, may
be rewritten as follows:
[L]org
DL =
[H4L4+] + [H3L3+] + [H2L2+] + [HL+] + [L]free
This equation can be rearranged to give an expression that shows the dependence of KD on
pH (H+ concentration):
KD
DL =
+
4
4
+
3
3
+
2
2
+
[H
]
[H
]
[H
]
[H ]
β
1
(
+
+
+
+ 1)
β
β
β
where [H+] is the hydrogen ion concentration determined by pH measurement, KD is the
distribution equilibrium constant, and β is the overall acid dissociation constant for acid
dissociation (e.g. β4 = K1 × K2 × K3 × K4). This expression can be rearranged as
ꢁ
ꢂ
[H+]4 [H+]3 [H+]2 [H+]
DL ·
+
+
+
+ 1 = KD
β4
β3
β2
β1
For the ultimate application of these complexants in solvent extraction, it is important to
evaluate the solubility of the reagent in aqueous and acceptable organic solutions. Because

J Solution Chem (2012) 41:1–16
13
a
Fig. 6 Distribution of N-donor
ligands between an aqueous
−1
phase (0.5 mol·L NaClO )
4
and 1-octanol as function of p H:
C
E, BPMDAE; F, BPMDIE; 2,
a
BPMDAP. Distribution fit
shown by dotted lines
a
Fig. 7 Distribution of N-donor
ligands between an aqueous
−1
phase (0.5 mol·L NaClO )
4
and 1-octanol as function of
p H: Q, BPMPIP; !, BPMDAC;
C
a
", BPMDIC. Distribution fit
shown by doted lines
these tetraaza complexants are moderately polar, it is clear that a polar organic solvent will
be needed for any separations application. The distribution of each ligand in 1-octanol as a
function of pCH is shown in Figs. 6 and 7.
The acid dissociation equilibria in aqueous solution also must be considered in develop-
ing the phase partitioning equilibria. The calculated values for KD for each ligand are shown
in Table 5. The KD of the ligand was determined using the pKa values from the potentio-
metric titrations at 0.5 mol·L−1 ionic strength. The KD value for all the complexes were
fitted using SOLVER program in Excel. In most instances a distribution of the HL+ had to
be included into the fitting to get a closer representation of the experimentally determined
distribution value.
The KD for BPMDIE, trans-BPMDIC, and BPMDIB have been determined previously
between nitrobenzene and aqueous phases [16, 17] as 0.98, 0.80, and 1.87, respectively.
From these distribution constants and that calculated for TPEN (KD = 0.82) [18] it was de-
termined that the distribution equilibria are primarily governed by the lipophilicity of the
2-pyridyl moiety rather than the aliphatic backbone of the ligand. The caveat to these pre-
vious distribution studies of BPMDIE, trans-BPMDIC and BPMDIB is that the tendency
for the ligands to hydrolyze was never taken into account. For the distribution of the amine
complexants into octanol, the opposite effect is seen, in that the lipophilicity of the back-
bone directly correlates to the phase distribution. BPMDAC has the highest KD whereas
BPMDIC and BPMDIE fall exactly as would be expected considering that the ligand under-
goes hydrolysis.

14
J Solution Chem (2012) 41:1–16
Table 5 Distribution coefficients
+
Compound
K
(HL
)
K (L)
D
D
of N -tetradentate ligands from
4
1-octanol arranged by K
D
BPMDAE
BPMDAP
BPMDAC
BPMPIP
0.32 0.09
0.26 0.09
2.9 0.1
0.71 0.06
1.89 0.02
7.7 0.6
a
NC
4.4 0.1
a
BPMDIE
BPMDIC
NC
0.36 0.04
2.0 0.1
0.50 0.05
a
NC = not calculated
The KD values for BPMDIE, and BPMDIC must be treated with some caution, since
these ligands readily undergo hydrolysis leading to the composition of the organic and aque-
ous phase changing as a function of time and pH. The chromophores used to identify the
phase transfer of these ligands are the hydrolysis products and not the ligand. The clearly
observable double step in the distribution of BPMDAC (Fig. 7) is noteworthy. In the distri-
bution curve, an initial partitioning of the ligand into the organic phase around a pH of 5.7 is
observed, which coincides with the third pKa of BPMDAC as shown in Table 1. It appears
that both H(BPMDAC)+ and the neutral BPMDAC ligand partition to the 1-octanol phase.
More data points at lower pHs are needed to fit the KD accurately but it appears that the KD
for HBPMDAC+ is calculated as 2.9 0.1 (Table 4); if taken directly from the level part of
the distribution curve a value of 2.6 0.2 can be determined.
4 Conclusions
The determination of the individual pKa values for the ligands with amine backbones
(BPMDAE, BPMDAP, BPMDAC and BPMPIP) follows the corresponding trend in the over-
all basicity of the free amine [13], in that:
BPMDAP > BPMDAE > BPMDAC > BPMPIP
propane-1,3-diamine > ethane-1,2-diamine > trans-cyclohexane-1,2-diamine > piperazine
The values of the last two pKa values for BPMDAE, BPMDAP, BPMPIP, and BPMDAC
indicate the more acidic nature of the pyridinium moieties when compared with the aliphatic
analogue triethylenetetramine (trien) [13]. The pK1 for both BPMDIE and BPMDIC match
that of 2-pyridine carboxaldehyde at 3.8 [19] and suggests that both of these ligands undergo
rapid hydrolysis in aqueous solution and are not compatible with a solvent extraction system.
From these studies it has been ascertained that the relative affinity of the N-based donor
ligands for protons does not correlate with its affinity for the divalent copper ion. The stabil-
ity constants for the tetradentate ligands, containing secondary amines, are exceedingly high.
The most preorganized ligand, BPMDAC, has the highest value of log10 K101 (19.5) with
BPMDAP and BPMDAE running up with values of 18.9 and 17.7, respectively. BPMDAP
has a higher Cu(II) log10 K101 than BPMDAE which is indicative of the greater stability of
Cu(II) for 6-membered chelate rings over 5 [20]. The log10 K101 value for BPMPIP is ap-
proximately 10 orders of magnitude smaller than the value for BPMDAE. This deficit is at
least partially due to the energy penalty associated with the reorganization of the piperazine
ring from the chair to the boat configuration. The lowest value for the Cu(II) log10 K101 is
for the bidentate ligand PBIm (4.4).

J Solution Chem (2012) 41:1–16
15
The formation of a Cu(II) chromophore allows the direct detection of the ligand and the
measurement ligand distribution ratio as a function of pH. The distribution coefficient for
the ligand between 1-octanol and an aqueous phase as a function of pH is as follows:
BPMDAC > BPMPIP > BPMDAP > BPMDAE
Contrary to previous distribution studies in nitrobenzene [16, 17] the distribution of the
ligands in 1-octanol directly correlates to the lipophilicity of the backbone. What is clear
from the distribution studies of all the tested ligands is that extraction experiments must be
carried out at higher pH to prevent loss of ligand to the aqueous phase. It can be argued that
the KD values for these ligands are too low to be useful in conventional solvent extraction
but they may still prove to be effective at separating Am(III) from Ln(III) much like TPEN,
under the following conditions, 0.001 mol·L−1 TPEN, 0.1 mol·L−1 NH4NO3, pH = 4.16
and nitrobenzene as the organic phase can still achieve as separation factor of 200 [21] even
though the TPEN KD is 0.82 [18].
The work presented here is part of an on-going project to determine the structural features
of complexing ligands that are important in increasing the binding affinity and selectivity for
An(III) over Ln(III). With a better understanding of the factors that are required for affinity
and selectivity, more effective and efficient ligands can be designed.
Acknowledgements The authors would like to thank Dr. Mikael Nilsson, Dr. Sarah Pepper, Dr. Peter
Zalupski and Dr. Syouhei Nishihama for their support and insight in this project. This research was conducted
at WSU with funding provided by the U.S. Department of Energy, Division of Nuclear Energy Science and
Technology, Nuclear Energy Research Initiative Consortium (NERI-C) program under project number DE-
FG07-07ID14896.
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