Chemistry of the copper(I)-water bond. Some new observations

Article abstract of DOI:10.1039/b105443k

One-dimensional polymeric copper(I) complexes of the type {[CuL(H₂O)]BF₄ · H₂O}_n, where L = 2,3- diphenylquinoxaline, and {[CuL′(H₂O)]X}_n, where L′ = 2,3-dimethylquinoxaline and X^- = ClO₄^- or BF₄^-, containing a rare copper(I)-water bond were synthesised. From the X-ray crystal structures of two of them, the copper(I) centres in these complexes are found to have a planar T-shaped N₂O coordination sphere. It is concluded from the observed Cu(I)-O(water) bond lengths [2.167(7)-2.307(14) A] that the copper(I)-water bonds in these complexes are rather weak. With L, a monomeric complex of the type CuL₂ClO₄ has also been synthesised. But it has not been possible to obtain such a monomeric copper(I) complex with the BF₄^- anion or the ligand L′. In CuL₂ClO₄ the metal is also found to have, from the X-ray crystal structure, a planar T-shaped N₂O coordination sphere with the perchlorate anion very weakly bound to the metal through an oxygen atom [Cu(I)-O(perchlorate) = 2.442(8) A]. While in the solid state electronic spectra, CuL₂ClO₄ displays a band at 346 nm, the aqua complexes show additional band(s) in the 400-480 nm range. CuL₂ClO₄ reacts with water in dichloromethane to yield an aqua copper(I) complex: CuL₂ClO₄ + H₂O → [CuL₂(H₂O)]ClO₄ → 1/n{[CuL- (H₂O)]ClO₄}n + L. In cyclic voltammetry at a glassy carbon electrode in anhydrous dichloromethane under N₂ atmosphere, CuL₂ClO₄ shows a quasi-reversible Cu^II/I couple with a very high redox potential of 0.91 V vs. SCE, which is lowered to 0.79 V vs. SCE upon addition of water. This indicates that binding of water destabilises copper(I), a result expected on the basis of Pearson's HSAB Principle.

Full text of DOI:10.1039/b105443k

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Chemistry of the copper(I)–water bond. Some new observations
Jnan P. Naskar,a Shubhamoy Chowdhury,a Michael G. B. Drewb and Dipankar Datta*a
a
Department of Inorganic Chemistry, Indian Association for the Cultivation of Science,
Calcutta 700 032, India. E-mail: icdd@mahendra.iacs.res.in
Department of Chemistry, University of Reading, Whiteknights, Reading, UK RG6 6AD
b
Received (in Montpellier, France) 15th June 2001, Accepted 21st September 2001
First published as an Advance Article on the web 8th January2002
One-dimensional polymeric copper(I) complexes of the type {[CuL(H2O)]BF4 ꢀ H2O}n , where L ¼ 2,3-
À
À
,
diphenylquinoxaline, and {[CuL0(H2O)]X}n , where L0 ¼ 2,3-dimethylquinoxaline and XÀ ¼ ClO4 or BF4
containing a rare copper(I)–water bond were synthesised. From the X-ray crystal structures of two of them, the
copper(I) centres in these complexes are found to have a planar T-shaped N2O coordination sphere. It is
concluded from the observed Cu(I)–O(water) bond lengths [2.167(7)–2.307(14) A]that the copper( I)–water
+
bonds in these complexes are rather weak. With L, a monomeric complex of the type CuL2ClO4 has also been
synthesised. But it has not been possible to obtain such a monomeric copper(I) complex with the BF4À anion or
the ligand L0. In CuL2ClO4 the metal is also found to have, from the X-ray crystal structure, a planar T-shaped
N2O coordination sphere with the perchlorate anion very weakly bound to the metal through an oxygen atom
[Cu(I)–O(perchlorate) ¼ 2.442(8) A]. While in the solid state electronic spectra, CuL2ClO4 displays a band at
+
346 nm, the aqua complexes show additional band(s) in the 400–480 nm range. CuL2ClO4 reacts with water in
dichloromethane to yield an aqua copper(I) complex: CuL2ClO4 + H2O ! [CuL2(H2O)]ClO4 ! 1=n{[CuL-
(H2O)]ClO4}n + L. In cyclic voltammetry at a glassy carbon electrode in anhydrous dichloromethane under N2
=
II
I
atmosphere, CuL2ClO4 shows a quasi-reversible Cu couple with a very high redox potential of 0.91 V vs.
SCE, which is lowered to 0.79 V vs. SCE upon addition of water. This indicates that binding of water
destabilises copper(I), a result expected on the basis of Pearson’s HSAB Principle.
Water is a very common ligand in the chemistry of copper(II),
but this is not so for copper(I). The reason can be understood
when Pearson’s Hard and Soft Acids and Bases Principle1 is
considered. In terms of Pearson’s hard-soft classification of
Lewis acids and bases, copper(I) is a ‘‘soft’’ acid while water is
a ‘‘hard’’ base and copper(II) a ‘‘borderline’’ acid. Since hard-
soft interactions are not particularly favoured, copper(I) and
water are not compatible with each other. This is manifested in
the experimental observation that in water2 simple Cu+ ion
disproportionates into Cu2+ ion and metallic copper with a
disproportionation constant of ꢁ 106. In an earlier commu-
nication,3 we have reported the synthesis and structure of a
polymeric copper(I) complex {[CuL(H2O)]ClO4 ꢀ 0.5C2H5OH}n
(1a, L ¼ 2,3-diphenylquinoxaline) containing a discrete cop-
per(I)–water bond. Herein we expound the chemistry of the
copper(I)–water bond further.
presence of hydrazine hydrate yields the reddish orange
{[CuL(H2O)]BF4 ꢀ H2O}n (1b). Reaction of 2,3-dimethylqui-
noxaline (L0) with Cu(ClO4)2 ꢀ 6H2O and Cu(BF4)2 ꢀ xH2O
carried out in a similar fashion yields {[CuL0(H2O)]ClO4}n (2a)
and {[CuL0(H2O)]BF4}n (2b), respectively, which are brownish
yellow in colour. Of the three new complexes, the structures of
1b and 2b have been determined by X-ray crystallography; the
structures of the cations in 1b and 2b are shown in Fig. 1 and 2,
respectively. The cations in both the complexes are one-
dimensional polymers, 1b along the y axis and 2b along the
z axis, both aligned around the 21 screw axis. In both the
cations, the copper atoms have essentially a planar, T-shaped
N2O coordination sphere with the N atoms coming from two
adjacent ligands and the oxygen atom from a water molecule.
The dimensions of the two coordination spheres are slightly
different with the two Cu–N distances in 1b being significantly
longer at 1.957(6) and 1.988(6) A than those [1.917(10) and
+
+
1.894(9) A]in 2b. This might well be due to the larger steric
effects of the substituent phenyl rings in 1b compared to the
methyl groups in 2b. However, as if to compensate, the longer
Cu–N distances in 1b are accompanied by a shorter Cu–O
+
+
distance as this is 2.167(7) A in 1b compared to 2.307(14) A in
2b. For comparison, we mention that the Cu(I)–water bond
+
length in 1a is 2.154(6) A.3 From the bond valence sum (BVS)
Results and discussion
models,4,5 which correlate the bond lengths around a metal
with its oxidation state, the ideal Cu(I)–O bond length in a
Earlier, by reacting Cu(ClO4)2 ꢀ 6H2O with L in ethanol in
equimolar proportions in the presence of a slight excess of the
reducing agent hydrazine hydrate, we have obtained 1a. Now
we report that when the anion is changed to tetrafluoroborate
or the substituents at the 2 and 3 positions of the quinoxaline
fragment are changed from phenyl to methyl, the copper(I)–
water bond is retained. Reaction of L with Cu(BF4)2 ꢀ xH2O in
ethanol in equimolar proportions (assuming x ¼ 4.5) in the
+
I
symmetric Cu O3 chromophore is calculated to be 1.99 A.
Thus, the copper(I)–water bond in all the complexes 1a, 1b and
2b is longer than expected and hence somewhat weak. Bond
length considerations indicate that of all the three complexes
the copper(I)–water bond is weakest in 2b. Incidentally, from
BVS calculations, the ideal Cu(I)–N bond length in a sym-
+
I
metric Cu N3 moiety is expected to be 1.98 A.
170
New J. Chem., 2002, 26, 170–175
DOI: 10.1039/b105443k
This journal is # The Royal Society of Chemistry and the Centre National de la Recherche Scientifique 2002

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Fig. 1 The structure of the one-dimensional polymeric cation in 1b with ellipsoids at 20% probability. The hydrogen atoms on the water molecule
+
were not located. Selected bond distances (A) and angles (ꢂ): Cu(1)–N(1) 1.957(6), Cu(1)–N(4)a 1.988(6), Cu(1)–O(100) 2.167(7), N(1)–Cu(1)–N(4)a
150.6(2), N(1)–Cu(1)–O(100) 109.9(3), N(4)a –Cu(1)–O(100) 99.4(3). Symmetry element (a): À x, y À 0.5, z À 0.5.
Earlier, when we reported 1a we indicated that it was the
first example of a copper(I) complex containing a discrete
copper(I)–water bond. However, while pursuing this problem,
we have come across two copper(I) complexes containing such
a bond.6,7 These are monomeric [Cu(phenazine)2H2O]ClO4
reported by Munakata et al.6 in 1993 and [Cu(1,4-thiox-
ane)3H2O]BF4 reported by Olmstead et al.7 in 1982. Of these
two complexes, that of Munakata et al. is directly comparable
with our complexes. In [Cu(phenazine)2H2O]ClO4 , the metal
When the reaction between L and Cu(ClO4)2 ꢀ 6H2O is car-
ried out in ethanol in the presence of hydrazine hydrate in n : 1
(n ¼ 2–4) molar proportions, we obtain the lemon yellow
CuL2 ꢀ ClO4 (3a). Its structure as determined by X-ray crys-
tallography is shown in Fig. 3. It is a monomer in which the
metal atom is bonded to two nitrogen atoms of two different
ligands in trans positions [1.933(7) and 1.942(8) A]and to the
+
perchlorate anion through one oxygen atom at a distance of
2.442(8) A. From our preceding discussion on Cu(I)–O bond
+
has a T-shaped, planar N2O coordination sphere with an
lengths, it is evident that the perchlorate anion is very weakly
bound to the metal in 3a. This is reflected in the solution
conductivity of 3a also; it behaves as a 1 : 1 electrolyte in
methanol. However, we have not been able to synthesise its
tetrafluoroborate analogue. Use of Cu(BF4)2 ꢀ xH2O in a pro-
cedure similar to that adopted for synthesising 3a has always
led to 1b. Further, all our attempts to obtain a monomer like
3a with the ligand L0 have so far failed. Irrespective of the
proportion of the metal and L0 used in the synthetic procedure,
2a has been obtained consistently when starting with
Cu(ClO4)2 ꢀ 6H2O and 2b with Cu(BF4)2 ꢀ xH2O. The reasons
for this are not understood.
+
average Cu–N bond length of 1.949(3) A and a Cu–O distance
+
of 2.241(3) A. Thus, the Cu(I)–water bond length in
[Cu(phenazine)2H2O]ClO4 falls in-between our complexes of
types 1 and 2. In [Cu(1,4-thioxane)3H2O]BF4 , the metal has a
tetrahedral S3O coordination sphere with the Cu–O bond
length being 2.234(7) A.
+
Planar T-shaped tricoordinate copper(I) complexes seem to
be quite common. For a pioneering example, see ref. 8. Several
mercury(II) complexes are known with somewhat distorted
T-shaped N2O coordination spheres in which the oxygen atom
belongs to a water molecule.9,10
Fig. 2 The structure of the one-dimensional polymeric cation in 2b with ellipsoids at 20% probability. The hydrogen atoms on the water molecule
+
were not located. Selected bond distances (A) and angles (ꢂ): Cu(1)–N(1) 1.917(10), Cu(1)–N(4)a 1.894(9), Cu(1)–O(100) 2.307(14), N(1)–Cu(1)–
N(4)a 164.4(4), N(1)–Cu(1)–O(100) 92.1(5), N(4)a –Cu(1)–O(100) 102.3(4). Symmetry element (a): Àx + 1, y + 0.5, z À 0.5.
New J. Chem., 2002, 26, 170–175
171

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Fig. 3 The structure of 3a with ellipsoids at 20% probability. Only one set of positions of the disordered perchlorate is shown. Selected bond
+
distances (A) and angles (ꢂ): Cu(1)–N(1B) 1.933(7), Cu(1)–N(1A) 1.942(8), Cu(1)–O(11) 2.442(8), N(1B)–Cu(1)–O(11) 100.0(3), N(1A)–Cu(1)–
O(11) 93.6(3), N(1B)–Cu(1)–N(1A) 166.3(3).
The complexes of type 1 and 2 are indefinitely stable in air
in the solid state. But in solution, almost immediately they
undergo aerial oxidation giving rise to faint green solutions.
Complex 3a is also indefinitely stable in air in the solid state; its
stability in solution in the presence of air is somewhat better
than that of complexes of type 1 and 2. An interesting feature
is revealed when the solid state electronic spectrum of 3a is
compared with those of the complexes of types 1 and 2.
Complex 3a shows a band at 346 nm; from its intensity in
methanol, it is recognised as being a charge transfer band that
may arise from a metal to quinoxaline charge transfer. In the
solid state electronic spectra of the complexes of types 1 and 2,
this band is present with minor variations in the position;
additionally they display band(s) at wavelengths longer than
400 nm (Fig. 4) assignable to some kind of charge transfer
between copper(I) and water. The complex 2a, which we have
not been able to characterise structurally, shows an electronic
spectra essentially similar to that of 2b. [Cu(phenazine)2-
H2O]ClO4 , the complex reported by Munakata et al.,6 also
displays two bands in the solid state at comparable positions,
372 and 420 nm.
The reaction is irreversible and occurs smoothly and almost
instantaneously in dichloromethane. Based on the change in
colour of a dichloromethane solution of 3a, from lemon yellow
to orange with the addition of water, it appears that first a
monomeric species of the type [CuL2(H2O)]ClO4 is formed,
but the compound that precipitates from the reaction mixture
by the addition of diethyl ether (see Experimental) is the
polymer {[CuL(H2O)]ClO4}n . Thus reaction (1) can be
decomposed into:
n CuL2ClO4 þ n H2O ! n ½CuL2ðH2OފClO4
ð1aÞ
ð1bÞ
n ½CuL2ðH2OފClO4 ! f½CuLðH2OފClO4gn # þ n L
Incidentally, the polymers 1a and 1b are insoluble in dichlor-
omethane. It is noted that reaction (1) does not proceed in
methanol, in which 3a is poorly soluble. In any case, reaction
(1) shows that the copper(I) centre in 3a has a considerable
affinity for water. The electrochemistry of the various com-
plexes described here has been examined by cyclic voltammetry
at a glassy carbon electrode under dry N2 atmosphere. While
in methanol complexes of types 1 and 2 show voltammograms
characteristic of adsorbed species, complex 3a displays an
oxidative, quasi-reversible voltammogram with an E1=2 of
0.91 V vs. SCE (saturated calomel electrode) in dry dichlor-
omethane (Fig. 5). Comparison of the voltammetric peak
currents with those of the ferrocene–ferrocenium couple under
the same experimental conditions establishes that the oxidative
response in 3a involves only one electron. Since L is not elec-
troactive in the potential range of interest here, the electrode
process observed for 3a is metal based:
We have found that 3a can be converted to an aqua complex
by adding water to a solution of 3a under N2 atmosphere:
n CuL2ClO4 þ n H2O ! f½CuLðH2OފClO4gn þ n L ð1Þ
CuII þ eÀ Ð CuI
The redox potential of the couple in eqn. (2) is quite high. To
ð2Þ
I
our knowledge, no electrochemical data on a Cu N2O chro-
mophore (though in a polar solvent like methanol CuL2ClO4
behaves as a 1 : 1 electrolyte, it is very likely that the ClO4
anion in 3a remains bound to the metal in the non-polar
dichloromethane solvent used for cyclic voltammetry) are
available with which we can compare the potential. For
=
II
comparison, we mention that the Cu potential for the Cu N4
I
I
moiety in [Cu(pyridine)4]+ is ca. 0.06 V vs. SCE, for that in
[Cu(bipyridine)2]+ it is ca. 0.01 V vs. SCE in 50% dioxane–
Fig. 4 The room temperature solid state electronic spectra of (a) 1a,
(b) 2a and (c) 3a in the region 800–300 nm.
=
water11 and the highest potential of the Cu couple for a
II
I
172
New J. Chem., 2002, 26, 170–175

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=
II
I
demonstrated in the lowering of the Cu potential in 3a upon
addition of water. We now mention that earlier work with
unsubstituted quinoxaline (Q), a polymeric copper(I) complex
of the type {[Cu(Q)3ClO4]ClO4}n , in which the bridging Q
moieties give rise to a three dimensional network, has been
reported.19 Some structurally uncharacterised polymeriÀc cop-
per(II) complexes of the type CuX2(L or L0) with X ¼ Cl , BrÀ
À
and NO3 are also known.20
Experimental
General
Fig. 5 Cyclic voltammograms of 3a in anhydrous dichloromethane
(—–) and in dichloromethane saturated with water (-----) at a glassy
carbon electrode under N2 atmosphere. Supporting electrolyte: 0.1 mol
dmÀ3 tetrabutylammonium perchlorate; concentration of 3a: 1.4 mmol
dmÀ3; scan rate: 50 mV sÀ1. Under the same experimental conditions,
the ferrocene–ferrocenium couple appears at 0.47 V vs. SCE with a
peak-to-peak separation of 95 mV.
All reagents were procured commercially. The melting points
reported here are uncorrected. Copper was estimated grav-
imetrically as CuSCN. Microanalyses were performed by a
Perkin–Elmer 2400II elemental analyser. FTIR spectra (KBr)
were recorded on a Nicolet Magna-IR spectrophotometer
(Series II) and UV-VIS spectra on a Shimadzu UV-160A
spectrophotometer. Cyclic voltammetry was performed at a
planar EG&G PARC G0229 glassy carbon milli electrode
using an EG&G PARC electrochemical analysis system
(model 250=5=0) in conventional three-electrode configura-
tions.
II
Cu N2O2 chromophore recorded so far is À 0.30 V vs. SCE in
acetonitrile.12 For a tetracoordinate copper complex, the
=
I
II
potential of the Cu couple is believed to increase with the p-
acidity of the ligand and the extent of tetrahedral distortion in
=
I
the corresponding copper(II) species.12,13 The highest Cu
potential hitherto recorded for any copper complex is 1.55 V
II
vs. SCE;14 the relevant copper complex has a Cu N4 chromo-
I
Syntheses
phore. Since the p-acidity of quinoxaline is likely to be com-
2,3-Diphenylquinoxaline (L) . Recrystallised o-phenylene-
diamine (OPDA; 0.54 g, 5 mmol and benzil (1.05 g, 5 mmol)
were taken in 25 ml of ethanol and refluxed for 5 h. Then
the yellow reaction mixture was left in air overnight. The
white needles that deposited were filtered off, washed with a
few drops of ethanol and dried in vacuo over fused CaCl2 .
Yield 1.15 g (80%), mp 120–122 ꢂC. Anal. found (calcd): C,
85.05 (85.07); H, 5.17 (5.00); N, 9.90 (9.92)%.
parable to that of pyridine,15,16 the high magnitude of the Cu
II=I
couple in 3a is attributable to the tricoordination of the metal
in 3a, which preferentially stabilises copper(I) over copper(II)
to a great extent. It should be noted that the electrode process
in eqn. (2) generates a tricoordinate copper(II) species in
solution. Until now, apart from biological systems, only two
structurally characterised tricoordinate copper(II) complexes
II
are known; in one case the metal ion has a Cu N2Cl moiety
17
=
I
II
II
and in the other Cu N2S. The potential of the Cu couple in
2,3-Dimethylquinoxaline (L0) . Recrystallised OPDA (0.54 g,
5 mmol) and 0.434 ml (5 mmol) of diacetyl were taken in 25 ml
of ethanol and refluxed for 5 h. Then the intense brown
reaction mixture was left in air for 4 days. The shiny white
crystals that deposited were filtered off, washed with a few
drops of ethanol and dried in vacuo over fused CaCl2 . Yield
0.25 g (30%), mp 102–104 ꢂC. Anal. found (calcd): C, 74.00
(74.02); H, 8.62 (8.70); N, 17.37 (17.27)%.
II
the Cu N2S chromophore was found to be À 0.42 V vs. SCE in
tetrahydrofuran. A type 1 copper site in human ceruloplasmin,
in which the coordination sphere around the metal is pre-
=
I
II
sumably trigonal N2S, has a Cu potential of ca. 0.78 V
vs. SCE.18
When water is added to the dichloromethane solution of 3a
used for cyclic voltammetry, the potential of the couple in eqn.
(2) is lowered. Fig. 5 shows that in dichloromethane saturated
=
II
I
with water, the Cu potential in 3a becomes 0.79 V vs. SCE.
In reaction (1) we have indicated that water replaces the
coordinated perchlorate anion in 3a. Thus, from Fig. 5 it is
apparent that coordination of water destabilises copper(I).
This is an expected result.
.
{[CuL(H2O)]BF4 H2O}n (1b) . Hydrated Cu(BF4)2 (0.24 g),
dissolved in 5 ml of ethanol, was added to 0.28 g (1 mmol) of L
dissolved in 20 ml of ethanol. The resulting bluish green
solution was heated to boiling. To this hot solution, 0.05 ml
of hydrazine hydrate (4 mmol) was added dropwise with
constant stirring. The resulting red solution was stirred for
5 min and the reaction mixture was left in air. After 10 min,
the brownish yellow crystalline compound that precipitated
was filtered, washed with 5 ml of ethanol and dried in vacuo
over fused CaCl2 . Yield 0.19 g (40%). Direct diffusion of
diethyl ether into the filtrate afforded shiny reddish orange
single crystals suitable for X-ray crystallography. Anal.
found (calcd): C, 51.09 (51.23); H, 3.77 (3.87); N, 6.09 (5.98);
Cu, 13.50 (13.56)%. FTIR (KBr) n=cmÀ1: 1642w (C¼N);
1131w, 1102w, 1066s, 1025s, 981m (BF4). UV-VIS (nujol)
Concluding remarks
Herein we have reported three new one-dimensional copper(I)
polymers containing a copper(I)–water bond synthesized using
2,3-diphenylquinoxaline (L) and 2,3-dimethylquinoxaline (L0)
as ligands. The structures of two of them have been determined
by X-ray crystallography. From the observed Cu(I)–O(water)
bond lengths, it is apparent that the bonding between copper(I)
and water is rather weak in these copper(I) aqua complexes.
With L, we have been able to isolate a monomer of the type
CuL2ClO4 (3a), which does not have a copper(I)–water bond.
We have demonstrated that it is possible to convert 3a to a
polymeric aqua complex of copper(I) by reacting it with water
in dichloromethane. However, it has not been possible
to synthesise the tetrafluoroborate or L0 analogue of 3a.
Pearson’s HSAB Principle predicts that bonding between
copper(I) and water is not particularly favoured. This we have
l
max=nm: 480, 402, 354.
{[CuL0(H2O)]ClO4}n (2a) . 2a was prepared in a manner
similar to that used for 1b by starting with 0.18 g (0.5 mmol)
of Cu(ClO4)2 ꢀ 6H2O and 0.08 g (0.5 mmol) of L0. The
resulting compound was brownish yellow. Yield 0.05 g
(30%). Anal. found (calcd): C, 35.46 (35.38); H, 3.50 (3.57);
New J. Chem., 2002, 26, 170–175
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Table 1 Crystallographic data for complexes 1b, 2b and 3a
1b
2b
3a
Chemical formula
Formula weight
Crystal system
C
20H18BCuF4N2O2
C10H12BCuF4N2O
C40H28ClCuN4O4
727.26
Triclinic
468.48
Orthorhombic
P22121
9.166(12)
13.014(17)
17.57(2)
—
—
—
2096(5)
4
1.095
326.44
Orthorhombic
Pn21a
15.08(2)
9.895(14)
8.700(14)
—
—
—
1298(3)
4
1.722
ꢀ
Space group
P1
+
a=A
+
11.052(14)
11.696(14)
14.049(17)
78.445(10)
74.361(10)
73.854(10)
1664(4)
2
b=A
+
c=A
a=ꢂ
b=ꢂ
g=ꢂ
+
U=A3
Z
m=mmÀ1
No. data collected
0.786
5463
5463
6913
3874
0.0539
1248
1248
0.065
No. unique data
Rint
R indices [I > 2s(I)]
R1
wR2
R indices (all data)
0.0726
0.1863
0.0721
0.1833
0.1162
0.3027
R1
wR2
0.1221
0.2103
0.0948
0.1949
0.1829
0.3402
N, 8.16 (8.26); Cu, 18.76 (18.74)%. FTIR (KBr) n=cmÀ1
:
CuL(H2O)ClO4 . Anal. found (calcd): C, 51.79 (51.82); H,
3.41 (3.48); N, 6.15 (6.05); Cu, 13.69 (13.72). FTIR (KBr)
n=cmÀ1: 1639w (C¼N); 1138m, 1115s, 1089vs, 1079vs, 625s
(ClO4). UV-VIS (nujol) lmax=nm: 470, 402, 360.
1632w (C¼N); 1105vs, 1055m, 999w, 620vs (ClO4). UV-VIS
lmax=nm (nujol): 425, 350.
{[CuL0(H2O)]BF4}n (2b) . 2b was prepared in a manner
Caution! Though we have not met with any incident while
working with the perchlorate compounds described here, care
should be taken in handling them as perchlorates are poten-
tially explosive. They should not be prepared and stored in
large quantities.
similar to that used for 1b by starting with 0.16 g (1 mmol)
of L0 and 0.24 g of hydrated Cu(BF4)2 . The resulting
compound was brownish yellow. Yield 0.06 g (20%). Direct
diffusion of diethyl ether into the filtrate gave shiny orange
single crystals suitable for X-ray crystallography. Anal.
found (calcd): C, 36.83 (36.76); H, 3.64 (3.71); N, 8.53 (8.58);
Cu, 19.50 (19.46)%. FTIR (KBr) n=cmÀ1: 1630w (C¼N);
1196m, 1178m, 1137m, 1002s, 980w (BF4). UV-VIS (nujol)
X-Ray crystallography
Details for the structure determinations of 1b, 2b and 3a are
given in Table 1. Intensity data were collected with Mo-Ka
radiation using a MAR research image plate system at
293(2) K. The crystals were positioned at 70 mm from the
image plate. One hundred frames were measured at 2ꢂ intervals
with a counting time of 2 min. Data analysis was carried out
with the XDS program.21 The structures were solved using
direct methods with the SHELXS-86 program.22 1b contains
three water molecules, each given 33.3% occupancy. In 3a the
perchlorate is disordered; the oxygen atom bonded to the
copper atom is ordered but the other three oxygen atoms are
disordered over two sites. In the three structures the non-
hydrogen atoms were refined with anisotropic thermal para-
meters. The hydrogen atoms bonded to carbon were included
in geometric positions and given thermal parameters equiva-
lent to 1.2 times those of the atom to which they were attached.
Hydrogen atoms of water molecules were not located.
An empirical absorption correction was carried out using
DIFABS.23 The structures were refined on F2 using
SHELXL.24
l
max=nm: 425, 359.
CuL2ClO4 (3a) . L (0.11 g, 0.4 mmol), dissolved in 20 ml of
ethanol, was added to 0.07 g (0.2 mmol) of Cu(ClO4)2 ꢀ 6H2O
dissolved in 5 ml of ethanol. The resulting bluish green
solution was heated to boiling. To this hot solution, 0.04 ml
(3.2 mmol) of hydrazine hydrate taken in 3 ml of ethanol
was added dropwise with constant stirring. Within a minute,
a shiny lemon yellow compound started precipitating from
the orange solution. After 5 min, the compound was filtered,
washed with 5 ml of diethyl ether and dried in vacuo over
fused CaCl2 . Yield 0.11 g (75%). Direct diffusion of diethyl
ether into the filtrate afforded lemon yellow crystals suitable
for X-ray crystallography. Anal. found (calcd): C, 65.97
(66.00); H, 3.95 (3.88); N, 7.68 (7.70); Cu, 8.71 (8.74)%.
FTIR (KBr) n=cmÀ1: 1642w (C¼N); 1145m, 1067vs, 1060m,
1040m, 620s, split (ClO4). LM (CH3OH): 97 OÀ1 cm2 molÀ1
.
UV-VIS (nujol)
l
max=nm (e=dm3 molÀ1 cmÀ1): 346;
(anhydrous, degassed CH3OH): 345 (1100), 255 (3000).
CCDC reference number 174164–174166. See http:==
www. rsc.org=suppdata=nj=b1=b105443k= for crystallographic
data in CIF or other electronic format.
Conversion of 3a to {[CuL(H2O)]ClO4}n . Solid lemon
yellow 3a (0.22 g, 0.30 mmol) was added to a mixture of
25 ml of dichloromethane and 0.075 ml (4.2 mmol) of water
with constant stirring under nitrogen atmosphere. Almost
immediately an orange turbid solution was obtained, which
was stirred for 5 min. Then 10 ml of deaerated diethyl ether
was added and stirring continued for another 5 min. The
light orange precipitate that appeared was then filtered off,
washed with 5 ml of diethyl ether and dried in vacuo over
fused CaCl2 . Yield 0.11 g (75%). The solid analysed as
Acknowledgements
M. G. B. D. thanks EPSRC and the University of Reading for
funds for the Image Plate System. D. D. thanks the Depart-
ment of Science and Technology, New Delhi, India for
financial support.
174
New J. Chem., 2002, 26, 170–175

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