Salicylaldehyde Hydrazones: Buttressing of Outer-Sphere Hydrogen-Bonding and Copper Extraction Properties

Article abstract of DOI:10.1071/CH16639

Salicylaldehyde hydrazones are weaker copper extractants than their oxime derivatives, which are used in hydrometallurgical processes to recover ～20% of the world's copper. Their strength, based on the extraction equilibrium constant K_e, can be increased by nearly three orders of magnitude by incorporating electron-withdrawing or hydrogen-bond acceptor groups (X) ortho to the phenolic OH group of the salicylaldehyde unit. Density functional theory calculations suggest that the effects of the 3-X substituents arise from a combination of their influence on the acidity of the phenol in the pH-dependent equilibrium, Cu²⁺+2L_org→[Cu(L-H)₂]_org+2H⁺, and on their ability to 'buttress' interligand hydrogen bonding by interacting with the hydrazone N-H donor group. X-ray crystal structure determination and computed structures indicate that in both the solid state and the gas phase, coordinated hydrazone groups are less planar than coordinated oximes and this has an adverse effect on intramolecular hydrogen-bond formation to the neighbouring phenolate oxygen atoms.

Full text of DOI:10.1071/CH16639

RESEARCH FRONT
CSIRO PUBLISHING
Aust. J. Chem.
Full Paper
http://dx.doi.org/10.1071/CH16639
Salicylaldehyde Hydrazones: Buttressing of Outer-Sphere
Hydrogen-Bonding and Copper Extraction Properties
A
A
B
A
Benjamin D. Roach, Tai Lin, Heiko Bauer, Ross S. Forgan,
A
A
A
Simon Parsons, David M. Rogers, Fraser J. White,
A C
,
and Peter A. Tasker
A
School of Chemistry, University of Edinburgh, Edinburgh EH9 3FJ, UK.
B
Fachbereich Chemie, Technische Universitat Kaiserslautern, 67663 Kaiserslautern,
¨
Germany.
C
Corresponding author. Email: p.a.tasker@ed.ac.uk
Salicylaldehyde hydrazones are weaker copper extractants than their oxime derivatives, which are used in hydrometal-
lurgical processes to recover ,20 % of the world’s copper. Their strength, based on the extraction equilibrium constant Ke,
can be increased by nearly three orders of magnitude by incorporating electron-withdrawing or hydrogen-bond acceptor
groups (X) ortho to the phenolic OH group of the salicylaldehyde unit. Density functional theory calculations suggest that
the effects of the 3-X substituents arise from a combination of their influence on the acidity of the phenol in the
pH-dependent equilibrium, Cu2þ þ 2Lorg Ð [Cu(L–H)2]org þ 2Hþ, and on their ability to ‘buttress’ interligand hydrogen
bonding by interacting with the hydrazone N–H donor group. X-ray crystal structure determination and computed
structures indicate that in both the solid state and the gas phase, coordinated hydrazone groups are less planar than
coordinated oximes and this has an adverse effect on intramolecular hydrogen-bond formation to the neighbouring
phenolate oxygen atoms.
Manuscript received: 11 November 2016.
Manuscript accepted: 3 January 2017.
Published online: 1 February 2017.
Introduction
shown in Eqn 2, whereas electron-donating substituents such as
alkyl and methoxy groups have the opposite effect.[3,8] Although
proligands with higher phenol acidity form conjugate phenolate
ions with lower basicity, which are poorer s-donors[9] and will
Approximately 20 % of the world’s copper is produced hydro-
metallurgically using phenolic oxime solvent extractants[1,2] of
the types[3,4] shown in Fig. 1. The acidity of the phenol allows
metal uptake and release to be controlled by varying the pH of
the aqueous phase (see Fig. 1).
The high selectivity for CuII over other first transition series
metal cations has been assumed to be due, at least in part, to the
goodness-of-fit[5] for the cavity in these pseudomacrocyclic
ligands. The head-to-tail hydrogen bonding in the complexes
is also often found in the proligands.[5]
In principle, variation of R, R0 or R00 groups in the structure in
Fig. 1 can be used to tune the extractant strength, which is
usually defined for these systems by the pH0.5 values (the pH at
which 50 % of the theoretical loading is observed). Variation of
the nature of an alkyl group para to the phenol (R in Fig. 1) does
not change pH0.5 values greatly,[3,6,7] but more highly branched
alkyl groups impart higher solubility in the hydrocarbon diluents
used in industrial applications.
In general, salicylaldoximes with R0 ¼ H are stronger extrac-
tants than ketoximes or benzophenone oximes where R0 groups
are alkyl or aromatic groups.[3] Groups present in the 3-position,
ortho- to the phenolic oxygen atom (R00 in Fig. 1), have very
significant effects on extractant strength.[3,6–8] Electron-
withdrawing substituents such as halide and nitro groups make
the proligands more acidic, favouring the extraction equilibrium
RЈ
H
O
H
N
O
Cu2ϩ
2Hϩ
2 R
RЉ
RЈ
O
H
H
RЉ
N
O
O
N
R
Cu
R
O
RЈ
RЉ
Cu2ϩ
2Hϩ
RЉ
RЈ
O
H
H
O
N
O
O
N
R
H
H
R
RЈ
RЉ
Fig. 1. The pH-dependent complexation of copper(II) by phenolic oxime
extractants.
Journal compilation ꢀ CSIRO 2017
www.publish.csiro.au/journals/ajc

B
B. D. Roach et al.
give a smaller overall formation constant b2, the square depen-
dence of Ke on Ka:
descriptor[11] and thus yield pseudomacrocyclic complexes
similar to those formed by the analogous oximes. In the present
paper, we consider whether 3-substitution of the phenol ring or
N-substitution of the hydrazone group can be used to tune
extractant strength to meet the requirements of commercial
operations. It has been reported that both the unsubstituted
ligand with R ¼ R0 ¼ X ¼ H (Fig. 2) and its N-phenyl analogue
(R ¼ X ¼ H, R0 ¼ C6H5) form neutral complexes of the type
[Cu(L–H)2].[12–16]
No work has been reported on the use of salicylaldehyde
hydrazones as copper extractants, but other types of hydrazones
have been used as spectrophotometric and gravimetric
reagents for the detection and analysis of transition metals,
including copper, nickel, and iron.[12,13,17] Some show unusual
magnetic[18,19] and electronic[20] properties, or are non-linearly
optically active[21–23] or fluorescent[24,25] materials. They can
also be used as optical chemosensors for analytical pur-
poses.[26,27] Salicylaldehyde benzoylhydrazone and analogous
aroyl hydrazones have been studied as chelating agents to
mobilize iron for iron-overload therapy,[28,29] and their transi-
tion metal complexes possess antitumour properties.[30–33]
Many hydrazones and their transition metal complexes are used
asantibacterial, antiviral, antitubercular,[34] antimycobacterial,[35]
and antifungal agents,[36,37] and in industry, they are employed as
plasticizers[38] and catalysts.[39–41]
N-acyl and N-aroyl hydrazones of salicylaldehyde[42–46]
(Fig. 3) can form 2 : 2 copper(II) complexes as tridentate
ligands with the phenolate oxygen, azomethine nitrogen, and
enolimide oxygen atoms defining the binding site.[42–46] In the
N-aroylhydrazones shown in Fig. 3, it is understood[42–46] that
the phenolate oxygen atoms act as the bridges between the two
copper atoms and, in some cases, a solvent molecule acts as a
fifth ligand for each copper atom.
The 2Cu : 2L stoichiometry and 3Cu : 2L stoichiometry of
related systems have been exploited to enhance the mass-
transport efficiency of copper extraction over that associated
with the 1Cu : 2L system currently used (Fig. 1) in commercial
processes.[47,48]
Ke Ð b2Ka2
ð1Þ
where Ke is the equilibrium constant for the formation of the
neutral complex in a single phase,
ꢀ
ꢁ
Cu2þ þ 2L Ð CuðLꢀHÞ2 þ 2Hþ
ð2Þ
b2 represents the overall formation constant for the 2 : 1
(ligand : CuII) complex,
ꢀ
ꢁ
Cu2þ þ 2ðLꢀHÞꢀ Ð CuðLꢀHÞ2
ð3Þ
and Ka represents the acid dissociation constant of the proligand
LH,
L Ð ½LꢀHꢁꢀ þ Hþ
ð4Þ
indicates that the phenol acidity has a greater effect on Ke than
the phenolate basicity.
Substituents in the 3-position also influence the interligand
hydrogen bonding shown in Fig. 1. Bulky t-butyl groups disrupt
this stabilizing motif, whereas hydrogen-bond acceptors such as
halide, nitro, and methoxy groups favour formation of bifurcat-
ed hydrogen bonds,[10] ‘buttressing’ the stabilizing motif,[6,7]
and thus increasing extractant strength.
The combination of electronic, steric, and hydrogen-bond
buttressing effects underpins the observed variation by more
than two orders of magnitudes of the distribution coefficients for
copper in the series of 3-substituted 3-tert-butylsalicylaldoxime
extractants: Br . NO2 . Cl . OMe . Me $ H . tBu.[6,7]
Salicylaldehyde hydrazones (Fig. 2) could also form N–H?O
hydrogen bonds to generate dimers with an R44 (10) graph set
RЈ
RЈ
X
H
H
N
H
N
N
O
R
N
O
Results and Discussion
Cu2ϩ
2Hϩ
Cu
R
O
N
R
L1–L10 were readily obtained in high yields (see Supplemen-
tary Material) by reaction of methyl or phenyl hydrazine with
the substituted salicylaldehydes 1–5, which were prepared by
formylation and 3-substitution of 4-tert-butylphenol
(Scheme 1). A stock of 5-tert-butyl-2-hydroxybenzaldehyde (1)
was synthesized by the Levin method,[49] and the nitro-substituted
N
H
X
RЈ
X
L1
L2
L3
L4
L5
OMe
L6
L7
L8
L9 L10
X
H
Me NO2 Br
H
Me NO2 Br
OMe
RЈ Me Me Me Me Me Ph Ph Ph Ph Ph
tBu tBu tBu tBu tBu tBu tBu tBu tBu tBu
R
H
N
O
RЈ
N
H
H
O
H
O
N
O
2Cu2ϩ
2 R
RЈ
O
O
N
R
4Hϩ
N
N
O
O
X
Cu
Cu
R
R
O1
O2
O3
O4
O5
X
R
H
Me NO2 Br
OMe
tBu tBu tBu tBu tBu
N
RЈ
Fig. 2. The 3- and N-substituted salicylaldehyde hydrazones L1–L10
and their pseudomacrocyclic copper complexes [Cu(L–H)2], and related
salicylaldoxime extractants O1–O5 discussed in this paper.
Fig. 3. Dinuclear complexes formed by N-acyl or N-aroyl hydrazones of
salicylaldehyde derivatives.[42–48]

Salicylaldehyde Hydrazone Copper Extractants
C
(3) and bromo-substituted (4) salicylaldehydes were produced by
optimizing literature conditions for electrophilic substitution.[50]
The methoxy-substituted salicylaldehyde (5) was prepared from
its bromo analogue 4 by a modification of the literature meth-
od.[51] In the case of 3-methyl-salicylaldehyde (2), the synthesis
exploited the modified Duff reaction[52] starting with 4-tert-butyl-
2-methylphenol, because it was commercially available and it
would be difficult to methylate 1 by Friedel–Crafts or related
reactions.
X-ray structure determinations of L6 and L10 (Fig. 4) reveal
that, as expected, the phenolic group acts as an intramolecular
bond donor, O–H?N, to the imine nitrogen atom and as an
intermolecular hydrogen-bond acceptor from a hydrazone NH
group in a neighbouring molecule. In both structures, the
hydrogen bonding involving the NH groups leads to the forma-
tion of linear polymers rather than pseudomacrocyclic dimers of
the type commonly found[5] in the structures of phenolic oximes
(see Fig. 1), which are preorganized for metal binding. In the
case of L10, the formation of the linear polymer is clearly
favoured by the hydrazone group being able to make an
additional bonding contact with the 3-methoxy group.
NMR methods were used to probe how tautomerism, phenol
acidity, hydrogen bonding, and intermolecular association in
solution (Fig. 2) are affected by the substituents on the nitrogen
atom or ortho to the phenol (Y or X respectively in Fig. 5)
because these are likely to have a significant influence on
1
strength as metal extractants. The H NMR spectra confirm
(see Supplementary Material) that the proligands exist in solu-
tion as the imino-enol tautomer (Fig. 5) as expected from the
substantial advantage of aromaticity.[53,54]
OH
X
X
H
Me
All the resonances in the 1H NMR spectra of the N-methyl-
hydrazones L1–L5, and the N-phenylhydrazones L6–L10, can
be unambiguously assigned. In the spectra of the N-methylhy-
drazones, the hydrazone proton is coupled with the NCH3
protons and appears as a quartet at d 7–8 ppm (Supplementary
Material). For the N-phenylhydrazones L6–L10, both the phe-
nol and hydrazone protons appear as singlets at d .10 ppm.
These can be distinguished by examining the COSY (correlation
spectroscopy) spectra, which indicate a weak coupling between
the azomethine proton and one of the very-low-field protons
(Fig. S1, Supplementary Material), which was therefore
assigned as the hydrazone.
tBu
(i) or (ii)
O
O
O
OH
X
X
H
Me
1
2
(vi)
tBu
(iii) or (iv)
OH
RHN
For the oximes O1–O5, the COSY spectra did not show spin–
spin coupling between the azomethine and oxime protons, which
suggests that the oxime oxygen, in contrast with the methyl
hydrazone NH group, is not much involved in the conjugation
of the molecule. Heteronuclear multiple-bond connectivity
OH
N
X
X
X
(vi)
3
4
NO2
Br
tBu
tBu
L1-L10
4
(v)
Y
H
Y
H
(vi)
N
O
N
O
6
1
OH
X
OMe
5
tBu
tBu
X
5
2
4
3
X
X
Imino-enol tautomer
Amino-keto tautomer
tBu
X ϭ H, Me, NO2, Br, OMe and Y ϭ OH, NHMe, NHPh
Scheme 1. Synthetic routes to salicylaldehyde hydrazones L1–L10:
(i) (CH2O)n/[Mg(OMe)2]/MeOH; (ii) hexamethylenetetramine/CF3CO2H;
(iii) HNO3/CH3CO2H; (iv) Br2/CH3CO2H; (v) NaOMe/MeOH;
(vi) MeNHNH2 or PhNHNH2/EtOH.
Fig. 5. Imino-enol and amino-keto tautomers of the structurally
related salicylaldoximes (O1–O5), N-methylhydrazones (L1–L5), and
N-phenylhydrazones (L6–L10).
2.11(3)
2.32(3)
2.22(2)
1.84(3)
1.81(3)
1.81(3)
2.06(3)
2.80(3)
˚
Fig. 4. X-ray crystal structures of the free ligands L6 and L10 with hydrogen-bond distances (A).

D
B. D. Roach et al.
(HMBC, also known as long-range 1H–13C COSY) experiments
were carried out to distinguish the phenol and oxime protons.
The HMBC spectrum of the unsubstituted oxime O1 in DMSO-d6
(Fig. S2, Supplementary Material) reveals correlation between
the azomethine 13C at 148.95 ppm and the proton at 11.25 ppm,
which confirms the latter to be the oxime proton.
shown[55,56] that there is a linear correlation between the pKa
values of phenols in aqueous solution and the chemical shift of
the phenolic protons at infinite dilution in DMSO. On this basis,
the chemical shift data for L1–L10 and O1–O5 in DMSO-d6
(Table 1) suggest that the acidity of the phenolic proton follows
the electron-withdrawing properties of the imine component
with oxime . phenylhydrazone . methylhydrazone. The mag-
nitude of the 3-substituent effects on acidity: NO2 . Br . Me .
H . OMe also follows that expected from their electron-
withdrawing properties.
Hydrogen bonding between proligands in solution, to form
cyclic dimers or linear oligomers, is of great relevance to their
preassembly and strength as metal extractants. The temperature
and concentration dependence of 1H NMR experiments (Fig. S3,
Supplementary Material) shows that, in contrast to the CH
protons, the shifts for the phenolic and oximic protons are
markedly temperature-dependent. As temperature is increased,
the larger shift of the hydrazone signal to high field compared
with that of the phenol is consistent with the former forming
inter- and the latter intramolecular hydrogen bonds.[57–61] The
preference for the phenolic proton to form an intramolecular
hydrogen bond to the imino nitrogen accords with solid-state
structures (see Fig. 4). An exception is the 3-nitro-substituted
proligands where the nitro group competes effectively for the
phenolic proton to give structures such as those shown in Fig. 6
(see also computational results below).
An order of electron-withdrawing effect, oxime . phenyl-
hydrazone . methylhydrazone, is supported by the comparison
1
of H NMR spectra of the unsubstituted proligands L1, L6,
and O1 with 4-tert-butylphenol, and those of their 3-methyl-
substituted analogues L2, L7, and O2 with 4-tert-butyl-2-
methyl-phenol (see Supplementary Material). It has been
Table 1. Chemical shifts (ppm) of the phenolic hydrazone and oxime
protons of L1–L10 and O1–O5 in DMSO-d6
3-Substituent
Methylhydrazones
Phenylhydrazones
Oximes
O1 9.88
H
L1
11.19
11.57
12.94
12.26
11.08
L6
10.40
11.09
11.76
11.64
9.99
Me
NO2
Br
L2
L3
L4
L5
L7
O2 10.12
O3 11.13
O4 10.69
L8
L9
MeO
L10
O5
9.48
ZH
N
N
O
Z
The conformation and aggregation of the methylhydrazones
L1–L5 were investigated by nuclear Overhauser effect spectros-
copy (NOESY). The spectrum of L2 in CDCl3 (Fig. 7) demon-
stratesthattheazomethineprotonisclosetoboththeN-substituted
methyl group and the benzene C6 proton. This is consistent with
the presence of a phenol OH to imino N intramolecular hydrogen
bond in the E conformation (see Fig. 8).
H
tBu
O
H
O
tBu
H
O
N
N
O
O
E-isomer
Z-isomer
Fig. 6. The E and Z isomers of the 3-nitrosubstituted ligands L3, L8, and
O3 in which ZH ¼ MeNH, PhNH, and OH.
The E conformation implied by the NOESY spectrum allows
the pseudomacrocyclic dimer to be formed as shown in Fig. 8.
H
H
N
N
1
6
4
Azomethine 6-position
NHCH3
5
2
tBu
OH
3
1
2
3
4
5
6
7
8
9
10
11
11
10
9
8
7
6
5
4
3
2
1
Chemical shift [ppm]
Fig. 7. NOESY spectrum of L2 in CDCl3.

Salicylaldehyde Hydrazone Copper Extractants
E
X
the related oxime complexes.[5] As the structures all have the
copper atom on a crystallographic inversion centre, the
bonding cavity radius[5] defined by the N2O2 donor set is
the mean of the Cu–O and Cu–N lengths. The smaller value
for the oxime complex [Cu(O1–H)2] implies tighter binding
than in the hydrazone complexes but caution needs to be
exercized in using cavity radii determined from solid-state
structures to compare the equatorial planar fields defined by
the N2O2 donor sets because axial contacts vary considerably
between structures;[5] for example, the phenolate oxygen
atoms of adjacent complexes make close contacts (2.670
3
N
H
H
H
H
N
N
H
H
H
H
tBu
N
O
O
N
N
O
H
H
6
6
6
tBu
2 t
Bu
H
H
3
3
X
X
E conformation
E,E dimer
Fig. 8. Formation of cyclic dimers from the E conformers of the salicy-
laldehyde methylhydrazones showing possible hydrogen-bond buttressing
by the 3-X groups.
˚
(2) A) in the structure of [Cu(L1–H)2] (see Fig. 9) but not in
[Cu(L5–H)2].
An interesting feature of the structures of [Cu(L1–H)2] and
[Cu(L5–H)2] is that the hydrazone N–N bonds bend away from
the central CuN2O2 unit to a greater extent than the N–O bonds
in the related oxime complexes. This is manifest (Table 2) by
larger deviations of the terminal hydrazone nitrogen atoms from
the least-squares planes defined by the copper atom and the
donor set (CuN2O2) than the oximic oxygen atoms and by the
related torsion angles. The implications of this are that inter-
ligand hydrogen bonding is weaker in the hydrazones and will
reduce complex stability and extractant strength. This and the
effects on the hydrogen bonding by the 3-X substituents are
considered further in density functional theory (DFT) calcula-
tions (see below).
A preliminary investigation of the strength of the hydrazones
as solvent extractants by studying the pH-dependence of the
reaction
ꢀ
ꢁ
CuSO4 þ 2L
Ð CuðLꢀHÞ
þ H2SO4
ð5Þ
ðCHCl3Þ
2
ðCHCl3Þ
indicated that they are intrinsically weaker extractants than their
oxime analogues. Two problems were encountered in obtaining
solvent-extraction data. For some of the weaker extractants,
loading is only observed at pH values at which precipitation of
copper(II) hydroxide occurs and several of the ligands and their
CuII complexes were insufficiently soluble in chloroform to
allow extraction experiments to be conducted. To circumvent
these problems, strength data were obtained by determining
the pH values for 50 % loading of copper from stripping
experiments in which chloroform solutions of preformed
[Cu(L–H)2] were put in contact with aqueous solutions having
different acidities but a constant sulfate (0.01 M) concentration.
The pH-dependence of copper loading curves and pH0.5 values
are shown in Fig. 10.
Fig. 9. The solid-state structures of [Cu(L1–H)2] (top); and [Cu(L5–H)2]
(below).
Clearly, 3-substitution significantly affects the extraction
strength, and the distribution coefficient for copper transfer
varies by more than three orders of magnitude across the series:
NO2 . Br . H $ OMe $ Me. It is apparent that the hydrazones
are intrinsically weaker extractants than their oxime analogues
(pH0.5 data[6] for the latter, O1–O5, are included for comparison
in the table inset in Fig. 10). The unsubstituted hydrazone (L1)
and its 3-methyl and 3-methoxy derivatives (L2 and L5) show
distribution coefficients for CuII approximately three orders of
magnitude smaller than their oxime analogues (O1, O2, and
O5). The differences in strength between the 3-nitro- and
3-bromo-substituted hydrazones L3 and L4 and their oxime
analogues (O3 and O4), and between all the oximes (O1–O5) is
much smaller.
The NMR results suggest that in solution, the hydrazone
proligands L1–L10 show a similar facility for aggregation and
preorganization to their much more studied oxime analogues.
Consequently, their abilities to function as extractants for CuII
are of interest.
All the hydrazone proligands were found to form 1 : 2 Cu : L
complexes [Cu(L–H)2] on reaction with copper(II) acetate in
methanol. X-ray structure determinations of [Cu(L1–H)2] and
[Cu(L5–H)2] confirm that the methylhydrazones yield planar
CuII complexes with pseudomacrocyclic structures similar to
their oxime analogues (Fig. 9). The intramolecular contacts
between terminal hydrazone nitrogen atoms and the phenolate
oxygen atoms Y1?O2 in the hydrazone complexes are slightly
longer than those defining the intramolecular hydrogen bond in
the oxime complex [Cu(O1–H)2] (see Table 2).
In an attempt to define to what extent the electronic, steric,
and hydrogen-bond buttressing properties of the 3-X substitu-
ents determine the relative strengths of the extractants, DFT
The Cu–N and Cu–O bond lengths in the hydrazone
complexes (Table 2) fall in a similar range to those found in

F
B. D. Roach et al.
Table 2. Bond lengths and angles in the inner coordination spheres of the [Cu(L1–H)2] and [Cu(L5–H)2] compared with those in [Cu(O1–H)2]
X
R
Y2
Compound
[Cu(L1–H)2]
[Cu(L5–H)2]
[Cu(O1–H)2]
X
H
Y
N
N
O
R
H
O1
Cu
tBu
N2
Me
Me
tBu
OMe
H
N1
R
O2
Y
1 H
-
X
A
[Cu(L1–H)2]A
[Cu(L5–H)2]A,B
[Cu(O1–H)2]A,C
˚
Bond lengths (A) and angles (8)
Cu–O1
1.9244(11)
1.9986(13)
180.0
1.8658(10)
2.0020(13)
1.905(2)
1.941(2)
1.900(2)
1.945(2)
180.0
1.907(2)
1.943(2)
Cu–N1
O1–Cu–O2
N1–Cu–N2
O1–Cu–N1
O2–Cu–N1
Y1?O2
180.0
180.0
180.0
180.0
180.0
180.0
180.0
180.0
91.03(5)
88.97(5)
2.7005(18)
1.962(2)
160.55(10)
19.45(10)
0.418(1)
92.24(5)
87.76(5)
2.604(4) 2.614(4)B
92.26(8)
87.74(8)
2.581(2)
1.923(3)
173.8(2)
6.20(16)
0.135(2)
92.13(8)
87.87(8)
2.583(5)
1.923(3)
177.9(2)
2.12(16)
0.046(2)
91.76(8)
88.24(8)
2.584(3)
1.925(3)
175.5(2)
4.54(16)
0.099(2)
Cavity radiusD
O1–Cu–N1–Y1
O2–Cu–N1–Y1
1.934(2)
168.7(2), 162.8(2)B
11.3(2), 17.2(2)B
0.247(2), 0.380(4)B
Distance of Y from CuN2O2 plane
AIn all complexes, the CuII atom lies on an inversion centre.
BSeparate values are given for the two molecules having differently disordered hydrazone groups.
CSee Forgan et al.[6]
DThe mean distance of the N and O donor atoms from their centroid.
100
HydrazoneA
Oxime
pH0.5
80
60
40
20
0
X
pH0.5
H
L1 4.42
L2 4.88
O1 1.68
O2 1.67
Me
O3B 0.70
O4 0.42
O5 1.09
NO2
Br
L3 1.23
L4 2.01
OMe L5 4.54
0
1
2
3
4
5
6
7
Equilibrium pH
Fig. 10. Copper loading of L1–L5 after placing 0.005 M chloroform solutions of [Cu(L–H)2] in contact with an
equal volume 0.01 M Na2SO4 aqueous solutions of various acidities. The pH0.5 values of the oxime analogues
O1–O5 are also listed. AFor [Cu(L3–H)2], 0.0005 M chloroform solutionswere used owing to the limitedsolubility.
BThe 5-t-octyl-substituted analogue as O4 and its copper complex have insufficient solubility to allow the
experiment to be conducted.
calculations were carried out to determine the enthalpies of the
reactions
the N–N bonds in the hydrazone complexes are bent away
from the coordination plane. In contrast, the N–O bonds in
the oxime [Cu(O1–H)2] lie very close to the CuN2O2 plane
(see O–Cu–N–Y torsion angles in Table 3). This is consistent
with the interligand O–H?Ophenolate hydrogen bonds being
shorter and stronger than the N–H?Ophenolate hydrogen
bonds in the hydrazone complexes. The greater strength
of the hydrogen bonds to the phenolate oxygen atoms in
[Cu(O1–H)2] is consistent with the oxime complex having
slightly weaker Cu–O bonds than those in the hydrazone
complexes (see natural bond orbital (NBO) values in
Table 3).
ꢀ
ꢁ
2LðgÞ þ Cu2þ ! CuðLꢀHÞ ðgÞ þ 2Hþ
ð6Þ
ðgÞ
ðgÞ
2
and
ꢀ
ꢁ
½Lꢁ2ðgÞ þ Cu2þ ! CuðLꢀHÞ ðgÞ þ 2Hþ
ð7Þ
ðgÞ
ðgÞ
2
(see Table 3)
The energy-minimized structures of the copper complexes
(see Supplementary Material) compare closely with those
determined by X-ray crystallography (see Fig. 11). As for the
X-ray structures (Table 2), the DFT calculations indicate that
The buttressing of interligand hydrogen bonding by the nitro-,
bromo-, and methoxy-groups in [Cu(L3–H)2], [Cu(L4–H)2],

Salicylaldehyde Hydrazone Copper Extractants
G
Side view
Top view
Fig. 11. The calculated structure of [Cu(L1–H)2] showing the displacement of the N–NHMe units from the coordination
plane.
Table 3. DFT calculated gas-phase deprotonation enthalpies (Edepr, L - [L–H]2 1 H1), ligand dimerization energies (Edimer, 2L - [L]2), Cu
complex formation energies, natural bond orbital (NBO) values, and geometric data for energy-minimized structures of Cu complexes
L1
L3
L4
L5
O1
X ¼ H
X ¼ NO2
X ¼ Br
X ¼ OMe
X ¼ H
Edepr [kcal mol
]
ꢀ1 C for L - [L–H]ꢀ þ HþA
364.48
ꢀ8.46
34.57
26.11
51.23
36.71
170.5
9.5
347.32
ꢀ17.08
30.44
13.36
52.01
37.07
172.5
7.5
357.19
ꢀ11.98
33.08
21.10
52.69
36.05
174.2
5.8
364.84
ꢀ10.92
32.80
21.88
54.70
37.26
171.9
8.1
356.42
ꢀ10.63
27.43
16.80
50.10
47.37
180.0
0.0
Edimer [kcal molꢀ1] [L]2A
Eform [kcal molꢀ1] [Cu(L–H)2] from [L]A2
Eform [kcal molꢀ1] [Cu(L–H)2] from 2LA
NBO for Cu–O [kcal molꢀ1 B
NBO for Cu–N [kcal molꢀ1 B
Torsion angle O–Cu–N–Y [8]B
Torsion angle O0–Cu–N–Y [8]B
NBO for Y–HyO [kcal molꢀ1 B
]
]
]
B
9.13
5.77
5.98
6.08
11.50
1.78
˚
Y–HyO contact distance [A]
Y–HyX contact distance [A]
Distance of Y from the least-squares mean plane CuN2O2 [A]
1.86
1.91
1.88
1.86
B
˚
2.91
2.02
2.81
2.72
2.99
A
˚
0.21
0.16
0.13
0.18
0.00
AEnergies and Cartesian coordinates for energy-minimized structures are provided in the Supplementary Material, Sections 4.1 and 4.2 respectively.
BMean of the two calculated values (no symmetry constraints were imposed during energy minimization).
C1 kcal molꢀ1 ¼ 4.186 kJ molꢀ1
.
and [Cu(L3–H)2] (see the Y–H?X contact distances, Table 3)
is, as might be expected, accompanied by a weakening of the
other interligand hydrogen bond (Y–H?Ophenolate).
The non-planar disposition of bonds about the coordinated
nitrogen atoms in the hydrazone complexes suggests that they
have sp3 character to a greater extent than those in their oxime
analogue. This is associated with them forming significantly
weaker Cu–N bonds than those in the oxime complex [Cu(O1–
H)2] and mirrors the relative Cu–N bond lengths found in the
solid-state structures (see Table 2).
The extent to which the 3-X substituents can buttress the
interligand hydrogen bonds is also revealed by the enthalpies of
formation of the proligand dimers (Edimer in Table 3). Incorpo-
ration of the hydrogen-bond accepting groups, NO2, Br, and
OMe, leads to more favourable enthalpies of dimerization than
that for the unsubstituted compound.
As discussed above, the NO2, Br, and OMe groups all show
bonding interactions in the outer coordination sphere with the
hydrazone hydrogen. Although the nitro-substituent forms a
particularly strong interaction with the hydrazone N–H group,
the dimerization enthalpy of L3 is only slightly more favourable
than that of the unsubstituted ligand L1. This is a consequence of
the monomer adopting a very favourable conformation that
allows the phenolic OH group to hydrogen bond to the adjacent
nitro substituent and the hydrazone N–H group to the phenolic
oxygen atom.
The calculated gas-phase formation energies for the copper
complexes([L]2(g) þ Cu2þ(g) - [Cu(L–H)2](g) þ 2Hþ(g),Table3)
are more favourable for those having the hydrogen-bond acceptor
groups NO2, Br and OMe. This trend is not so marked for the
methoxy-substituted ligand, mainly a consequence of its less
favourable deprotonation enthalpy. Although the predicted order
of extractant strength, NO2 . Br. OMe . H, appears to corre-
late reasonably well with that observed experimentally, NO2 .
Br. H $ OMe, the calculated values are based on the assump-
tion that the reagents are exclusively in the cyclic dimeric form
before copper uptake, and this is known not to be the case for the
nitro ligand L3 (see above). If the monomeric proligands are
taken as the starting forms, the nitro-substituted extractant still
shows the most favourable formation energy.
For L5, the electron-donating properties of the methoxy
substituent lead to a higher deprotonation enthalpy term than

H
B. D. Roach et al.
1
for the unsubstituted ligand L1, but this appears to be more
than compensated for by the more favourable strengths of
bonds to copper and the buttressed hydrogen bonding to the
hydrazone NH.
were purchased from Alfa Aesar. H and 13C NMR spectra
recorded in the experimental section were run on a Bruker
1
ARX250 at ambient temperature. H NMR NOESY spectra
were recorded on a Bruker DPX360 spectrometer at 298 K
(unless stated otherwise), and chemical shifts (d) are reported in
parts per million (ppm) relative to TMS. CHN analytical data
were obtained by the University of St Andrews Microanalytical
Service. Mass spectrometry was performed on a Micromass
ZMD instrument with a z-spray electrospray ionisation (ESI)
source. Melting points were measured on a Gallenkamp melting
point apparatus. ICP-OES analysis was performed on a Perkin
Elmer Optima 5300DV spectrometer. The measurement of pH
was carried out using a Sartorius PP-50 pH meter.
At first sight, the correlation between calculated energies of
formation in the gas phase and the relative strength of extrac-
tants in a two-phase liquid system is quite remarkable. As
discussed recently for other systems where this occurs,[62] this
correlation will only be observed when the process occurring in
the aqueous phase is the same, in this case two protons replacing
one copper dication, for the series of extractions. Consequently,
as in the current work, hydration energies of species do not
contribute to differences in the formation energies of metal
complexes. The solvation energies of the proligands and their
copper complexes in the organic phase will obviously show
some dependence on the nature of the 3-X substituent. However,
the differences in solvation energies of the preorganized dimers
and their copper complexes are likely to be fairly small and,
consequently, it is the effect of the 3-X substituents on the
deprotonation energies of the proligands and the binding ener-
gies of the conjugate anions to copper that largely determine the
energies of formation of the complexes.
X-Ray Crystal Structures
These were obtained by measuring suitably sized crystals on
either a Bruker D8 with a Smart Apex or Apex II charge coupled
device (CCD) or an Oxford Diffraction SuperNova with an Atlas
CCD. Data were collected at 150 K unless otherwise specified
and reduced with the relevant manufacturer’s software. All
structures were solved by direct methods and refined with
ShelXL. Refer to the Supplementary Material for the details for
each structure. In the crystal structure of [Cu(L5–H)2], there was
some disorder in the position of the hydrazone N–H group. This
was modelled as two components, with each fixed at 50 %
occupancy, determined from approximate values gleaned
through refinement. This approach was used as a consequence of
minor instability of the refinement and unreasonable, unequal
atomic displacement parameters when occupancies were refined
with linked occupancies summing to 100 % occupancy.
Conclusions
The approach of using the cavity size in macrocyclic proligands
to tune the strength of binding of base metal ions, established by
Lindoy and coworkers,[63] can be extended to much simpler
reagents such as the salicylaldehyde hydrazones described in the
present work, which assemble in pseudomacrocyclic structures
via interligand hydrogen bonding.
The N-methyl hydrazone derivatives (L1–L5) are analogues
of the commercial phenolic oxime reagents but are significantly
weaker copper extractants. On the basis of the DFT calculations
and X-ray structure determinations, it appears that this arises
from a combination of the weaker Cu–N bonds and the less
favourable interligand hydrogen bonds formed by the hydra-
zones (OH groups are generally better hydrogen bond donors
than NH groups[64,65]) and from their higher deprotonation
enthalpies. The N-phenyl hydrazones (L6–L10) are too weak
to allow solvent extraction to be carried out under conventional
conditions.
NMR Studies
Samples were generally prepared by dissolving 10 mg of the
ligand in 0.6 mL CDCl3 or DMSO-d6. Samples for NOESY and
1D NOE difference spectra required more concentrated solu-
tions, and 40 mg of the ligand in 0.6 mL CDCl3 was used. The
concentration-dependence NMR study was carried out by dilut-
ing an NMR sample containing 40mg of the ligand in 0.6 mL
DMSO-d6 to 1.2, 1.6, and 2.4 mL in the same NMR tube.
Solvent Extraction
The results confirm the importance of understanding outer-
sphere coordination chemistry in designing metal solvent
extractants and, in particular, the efficacy of interligand hydro-
gen bonding in extracted species. As for the oximes,[6,7] the
introduction of substituents adjacent to the phenolic hydroxyl
group with electron-withdrawing and hydrogen-bond acceptor
properties significantly increases the strength as copper extrac-
tants. For the N-methyl hydrazones, the distribution coefficient
for copper extraction is increased by three orders of magnitude
on introducing a 3-bromo or 3-nitro-substituent. The resulting
extractants have strengths comparable with the commercial
oximes, but their lower solubilities in water-immiscible solvents
and the higher costs of synthesis arising from introduction of the
substituent make them poor candidates to replace the tried and
tested commercial reagents.
Preliminary loading experiments were carried out by placing a
chloroform solution of the ligand (5.00 mL, 0.010 M) in contact
with an aqueous solution of copper(II) sulfate (5.00 mL,
0.010 M) at different pH values in a tightly sealed, screw-top
glass jar. The aqueous solution was prepared by adding an
H2SO4 solution (0.10 M for pH 1.5 and above, 1.00 M for
lower pH) or a NaOH solution (0.10 M) to a CuSO4 solution
(3.00 mL, 0.0167 M) and adding water to make up to 5.00 mL.
The two-phase system was stirred at 900 rpm at room temper-
ature for 18 h. A 1.00 mL aliquot was taken from the organic
phase, dried under vacuum, and redissolved in butan-1-ol
(10.00 mL). The residues from L3 and L4 solutions were
dissolved in nitrobenzene. The copper was then analysed by
ICP-OES. The pH of the aqueous phase was measured using a
pH meter. The calculated percentage of the copper(II) taken into
the organic phase was plotted against the measured equilibrium
pH to give the S-curve.
Experimental
Chemicals and Equipment
The data presented in Fig. 10 were obtained by taking a
chloroform solution of the copper(II) complex (5.00 ml,
0.005 M) and an aqueous solution of sodium sulfate (5.00 mL,
0.010 M) with varying H2SO4 content. In some cases, the low
Unless otherwise specified, reagents or solvents were used as
obtained from Aldrich, Fisher or Acros. Standards for induc-
tively coupled plasma optical emission spectroscopy (ICP-OES)

Salicylaldehyde Hydrazone Copper Extractants
I
solubility of the complex required slight modifications of the
procedure (Section 3 in the Supplementary Material).
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For the determination of energies of formation (rows 1–4 in
Table 3), DFT calculations employed the B3LYP hybrid
exchange-correlation functional[66,67] and the 6–31þG(d,p)
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03[73] with the B3LYP functional and 6–31Gþ(d,p) basis set.
More information, including Cartesian coordinates for atom
positions, are provided in Tables S4.1 and S4.2 in the Supple-
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copper complexes are described in detail in the Supplementary
Material (Section 5). L6: CCDC deposition number: 1517105;
L10: CCDC deposition number: 1517104; [Cu(L1–H)2]:
YUPBAO, CCDC deposition number: 1410136; [Cu(L5–H)2]:
YUNZUE, CCDC deposition number: 1410135.
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Supplementary Material
Additional NMR data for the proligands, experimental condi-
tions for solvent extraction experiments, positional parameters
for all DFT energy minimized proligand and copper complex
structures, and full experimental details for the synthesis and
characterization of proligands and complexes are available on
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We thank the Engineering and Physical Sciences Research Council, UK,
for funding, EaStCHEM for the award of a studentship to TL and for
access to the EaStCHEM Research Computing Facility, Cytec Industries,
and Infineum UK for funding of PhD studentships for RSF, BDR, and
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