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TUTORIAL REVIEW

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Anion recognition by coordinative interactions: metal–amine complexes as receptors† Luigi Fabbrizzi* and Antonio Poggi Alfred Werner’s complexes of formula [MIII(NH3)6nXn]X3n involved inert metal centres (M = Cr, Co), and anions X ‘frozen’ in the coordination sphere, a circumstance which allowed the isolation of a variety of isomers. Amine complexes of labile transition metal ions, studied later, do not form isomers, yet they allow the investigation of the fast and reversible interaction of the anion X with the metal–amine core. On these bases, anion receptors of varying degrees of sophistication have been synthesised, which consist of coordinatively unsaturated polyamine metal complexes and whose vacant coordination sites can be occupied by anion donor atoms. A thoughtful design of the polyamine framework may introduce geometrical selectivity, resulting from the matching between anion shape and size and the geometrical features of receptor’s cavity. Compared to their purely organic counterparts, metal containing receptors display several advantages: (i) metal–anion interactions are strong enough to more than compensate anion dehydration energy, which allows recognition

Received 30th July 2012

studies to be carried out in water; (ii) transition metal ions of different electronic configurations exhibit

DOI: 10.1039/c2cs35290g

different geometrical preferences, which addresses anion binding and introduces a further element of selectivity. Chosen examples of polyamine metal complexes, including macrocycles and cages, displaying

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selective binding tendencies towards anions will be illustrated in this tutorial review.

Dipartimento di Chimica, Universita` di Pavia, 27100 Pavia, Italy. E-mail: [email protected]; Fax: +39 0382 528544; Tel: +39 0382 987328 † Part of the centenary issue to celebrate the Nobel Prize in Chemistry awarded to Alfred Werner.

Luigi Fabbrizzi was born in 1946 in Florence, where he had all his education and obtained his degree in Chemistry in 1969. He was a post-doctoral fellow and a lecturer of inorganic chemistry at the University of Florence during the period 1971–1980. Since 1980 he is Professor of Chemistry at the University of Pavia. His research interests span a range of topics at the interface of supramolecular chemistry and inorganic Luigi Fabbrizzi chemistry, with a special regard to anion recognition and sensing, to the design of metal-based molecular devices and to metal promoted self-assembling. He is Honorary Professor of the East China University of Science and Technology in Shanghai. He was the recipient of the 2010 Izatt–Christensen Award on Macrocyclic Chemistry.

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The Royal Society of Chemistry 2013

1. Introduction Alfred Werner and some distinguished inorganic chemists before him1 dedicated their attention to the ammonia

Antonio Poggi graduated in 1979 from the University of Florence; in 1983 he was appointed as a Research Assistant at the Dipartimento di Chimica of the University of Pavia. Since 1987 he is an Associate Professor at the same department. His research activity deals with the coordination chemistry of multidentate ligands, covering in particular the topics of the redox properties of transition Antonio Poggi metal complexes and the design of functionalized ligands for the assembly of multi-centre redox systems. More recently, his research interests addressed the design of supramolecular systems able to act as sensors, especially toward anions.

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Fig. 1

Significant reactions involving [CoIII(NH3)5Cl]Cl2: (1) inner sphere electron transfer;5 (2) base hydrolysis;6 (3) photoinduced electron transfer.7

complexes of CoIII, CrIII, PtIV, because they are inert and their chemistry could be elucidated by preparative methods alone.2 The only physical experiment Werner, with his student Arturo Miolati,3 could carry out was the measurement of electrical conductance in solution. On these simple evidences, Werner developed a complete theory, capable of explaining stoichiometry, bonding, geometry, isomerism and stereochemistry of hundreds of inert metal amine complexes, synthesised by him and by his precursors. In these compounds, a determining role was played by anions, whether bound to the metal (‘acid residues’, contributing with ammonia molecules to the coordination number) or not (‘ionogenic’). Classical Werner’s complexes continued to raise the attention of chemists in the following decades and were instrumental to the development of coordination chemistry. As an example, the cobalt(III) complex salt, CoCl35NH3, called purpureo by Gibb and Genth in view of its red-purple colour,4 to which Werner ascribed the correct constitution [CoIII(NH3)5Cl]Cl2,2 has been the protagonist (or a co-protagonist) of several chemical processes that defined landmarks in the progress of chemistry: (i) inner sphere electron transfer (Taube et al., 1953);5 (ii) hydrolysis through the conjugate base mechanism (Basolo and Pearson, 1956);6 (iii) photoinduced electron transfer (Gafney and Adamson, 1972),7 as outlined in Fig. 1. In his studies, Werner tried also to isolate ammine–halide complexes and isomers with 3d divalent cations (FeII, NiII, CuII, ZnII),8 without success, due to the labile nature of the metal centres, which addressed the formation of thermodynamically stable complexes, in the absence of any kinetic control. Only decades after, stepwise equilibria of complex formation of labile ions with amines (e.g. [NiII(NH3)6]2+, stable complex) and with anions (e.g. [NiIICl4]2, less stable complex) were clearly elucidated, by means of pH-metric and spectrophotometric titration experiments in aqueous solution.9,10 However, the characterisation in the solution of mixed (ternary) complexes, containing in the coordination sphere both amine ligands and anions, seemed a rather elusive or a poorly interesting goal and, in any case, was not systematically pursued. Nevertheless, it was observed that, in crystalline ternary metal amine complexes, anions could exert a profound influence on the electronic properties of the metal, from which substantial differences derived in colour, magnetism, coordination geometry. An example was provided by bis-diamine complexes of nickel(II) of general formula NiIIL2X2. When X is a coordinating anion, (NCS, N3, Cl), the complex shows a more or less distorted octahedral geometry, with the two diamine ligands L occupying the equatorial positions and the two anions positioned in the axial sites.11

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Fig. 2 The crystal structure of complexes of formula NiII(N,N-diethylethylenediamine)2X2 (hydrogen atoms omitted for clatity): (a) X = NCS, NiIIN(amine) = 2.09, 2.30 Å, NiIIN(NCS) = 2.08 Å;12 (b) X = ClO4, NiIIN(amine) = 1.93, 1.97 Å, NiIIO(ClO4) = 3.75 Å.13

As an example, Fig. 2a shows the crystal structure of the NiII(adiEten)2(NCS)2 complex (adiEten = N,N-diethylethylenediamine).12 The salt has a blue-violet colour and is paramagnetic with a magnetic moment m = 3.20 B.M., corresponding to 2 unpaired electrons, a behaviour consistent with the regular octahedral geometry shown in Fig. 2a. On the other hand, when the anion X is poorly coordinating (ClO4, BF4, PF6), NiIIL2X2 complexes are yellow-orange and diamagnetic, as expected for a square coordination geometry. Fig. 2b shows the crystal structure of the orange NiII(adiEten)2(ClO4)2 complex:13 the perchlorate ions are far away from the coordination sphere and the complex is square, a geometry that favours electron pairing. Notice that NiII–N bond distances in the yellow diamagnetic complex are distinctly smaller than in the blue-violet paramagnetic one. During the 1970s, a few equilibrium studies were carried out on the interaction of copper(II) tetramine complexes with halides and pseudohalides, to give five-coordinate species, which disclosed only moderate anion selectivity. Selectivity would require a more elaborated design of the polyamine framework of the metal complex, which should provide geometrical constraints to anion interaction. Indeed, a more pronounced selectivity was later observed with dimetallic complexes with polyamine ligands of more elaborated design, in which the anion could interact with both metal centres. In particular, the amplitude of the cavity could be modulated through synthetic modifications of the cage framework, in order to accommodate anions of different sizes and shapes. These studies have shown that receptors containing coordinatively unsaturated metal centres may establish with anions stronger and more selective interactions than metal-free receptors, operating through electrostatic and/or hydrogen bonding interactions. This tutorial review is intended to illustrate with some selected examples the development of anion receptors

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containing metal ions and operating through metal–ligand interactions.


2. The interaction of anions with copper(II) bisdiamine complexes: the role of DHo and DSo Divalent cations of the first transition series are labile and tend to bind up to six molecules of ammonia, to give octahedral complexes, according to a sequence of stepwise equilibria. As an example, in the case of NiII, in aqueous solution, the following series of log Ki values is observed for the six steps (i = 1,. . .6): 2.80, 2.24, 1.73, 1.19, 0.75, 0.03.9 The decrease in log Ki values depends both (i) on the statistical effect and (ii) on a progressive decrease in the electrostatic contribution to the metal–ligand interaction. CuII is special, in view of its sensitivity to the Jahn–Teller effect: it binds quite strongly four ammonia molecules (log Ki = 4.15, 3.50. 2.89, 2.13), to give a square complex, but shows a very low affinity for the fifth one (log K5 = 0.52).9 In this sense, the [CuII(NH3)4]2+ complex could be an ideal receptor for a given anion X, offering to it a vacant binding site, to form the ternary species [CuII(NH3)4X]+. The constant of the [CuII(NH3)4]2+ + X # [CuII(NH3)4X]+ equilibrium should give a measure of the intensity of the CuIIX interaction. However, an aqueous solution containing Cu2+ and 4 equiv. of ammonia (obtained either by mixing standard solutions of the metal and of the ligand or by dissolving the solid [CuII(NH3)4]SO4H2O complex salt) does not contain only the integral complex [CuII(NH3)4]2+, but a mixture of species at the equilibrium. Fig. 3a shows the abundance of the species which form when a solution 103 M in Cu2+ is titrated with ammonia. It is observed that, on addition of 4 equiv. of NH3, the desired complex Cu(NH3)42+ is only a minor species, being present at 7%, whereas Cu(NH3)2+ is 6%, Cu(NH3)22+ 41% and Cu(NH3)32+ 47%. Thus, in order to have a receptor with a single available binding site for anions, a CuII tetramine complex of distinctly higher stability should be taken into consideration. In this sense, one could profit, for instance, from

Fig. 4 (a) Crystal structure of the [CuII(en)2Cl]PF6 complex salt;15 PF6 and hydrogens omittted for clarity; (b) crystal structure of the [CuII(adimeen)NCS][CrIII(NH3)2(NCS)4]2 complex salt;16 [CrIII(NH3)2(NCS)4]2 counteranions and hydrogens omittted for clarity.

the chelate effect, thus moving from ammonia to diamines. Fig. 3b shows the abundance of the species which form over the course of the titration with N,N-dimethylethylenediamine (adiMeen = L) an aqueous solution 103 M in Cu2+.14 We notice that, on addition of 2 equiv. of the diamine, the [CuIIL2]2+ complex is formed at 98%, which suggests that copper(II) bis-diamine complexes can be good anion receptors in water and in other polar media. The tendency of copper(II) bis-diamine complexes to bind one anion X has been confirmed by the isolation of crystalline salts containing the five-coordinate [CuII(diamine)2X]+ ternary species, whose structure has been elucidated by X-ray diffraction studies. As an example, Fig. 4a shows the structure of the [CuII(en)2Cl]PF6 salt (en = ethylenediamine).15 The five-coordinate [CuII(en)2Cl]+ complex shows a regular square pyramidal geometry. The CuII centre is not exactly coplanar with the plane of the four nitrogen atoms, but stays above it 0.15 Å. On the other hand, alkyl substitution at the amine nitrogen atoms does not seem to affect seriously the coordination geometry. In fact, the complex shown in Fig. 4b, [CuII(adiMeen)2NCS]+, presents

Fig. 3 Relative abundance (with respect to the analytical concentration of CuII) of the copper(II) complexes which form over the course of the titration of a 103 M aqueous solution of Cu2+ with (a) ammonia,9 and (b) N,N-dimethylethylenediamine (L). On addition of 4 equiv. of NH3, the tetramine complex [Cu(NH3)42+] is present only at 6%, whereas, on addition of 2 equiv. of diamine, the [CuIIL2]2+ complex is formed at 98%.

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Fig. 5 Log K values for the equilibria: [CuII(diamine)2]2+ + X # [CuII(L)2X]+ in MeOH at 25 1C.17

a slightly distorted square pyramidal geometry, with the CuII centre 0.21 Å above the mean plane of amine nitrogen atoms.16 The thiocyanate nitrogen atom occupies the apical position of the square pyramid, with a CuIIN distance (2.18 Å) larger than those of CuIIN(amine) bonds: primary 2.00 Å, tertiary 2.10 Å. Equilibrium studies were carried out in MeOH on the interaction of halide ions with complexes of formula [CuIIL2]2+, where L = en (ethylenediamine). Meen (N-methyl-ethylenediamine), diMeen (N,N0 -dimethylethylenediamine) and adiMeen.17 The diagram in Fig. 5 shows the log K values for the equilibria: [CuIIL2]2+ + X # [CuIIL2X]+, determined through spectrophotometric titrations. In particular, anion addition induces a more or less pronounced red-shift in the d–d band of the copper(II) bis-diamine complex. Two points have to be emphasised: (i) anion addition is favoured by the presence of methyl substituents on the amine nitrogen atoms (which is contrary to what expected in terms of steric repulsions); (ii) each [CuII(diamine)2]2+ receptor does not exert any selectivity in halide ion recognition, with the exception of [CuII(en)2]2+, for which log K values moderately decrease along the series Cl 4 Br 4 I. These equilibria were also investigated through calorimetry, from which DHo values were determined and TDSo values calculated. Fig. 6 displays the thermodynamic quantities DGo, DHo and TDSo (kcal mol1) associated to each [CuII(diamine)2]2+ + X # CuII(diamine)2X]+ equilibrium.

Fig. 6

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Thermodynamic quantities DGo, DHo and TDSo (kcal mol1) associated with the [CuII(diamine)2]2+ + X# [CuII(diamine)2X]+ equilibria in MeOH at 25 1C.17

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the much less solvated [CuII(diamine)2]2+ receptor. Thus, it appears that pre-coordination by a polyamine framework, leaving a vacant coordination site, does not substantially affect the affinity of metal ions for anions and opens the way to the design of more sophisticated receptors, hopefully capable of exerting selective recognition.


3. Anion binding by tetramine complexes moves to water: cyclam derivatives Cyclam (1,4,8,11-tetraazacyclotetradecane) forms the most stable tetramine complexes with divalent transition metals.20 The high thermodynamic stability results from the favourable preorganisation of the ligand, whose secondary amine nitrogen atoms are positioned at the corners of a square.21 On the other hand, transition metal ions fit well the cavity of the macrocycle, relaxed to its less energetic conformation. The great majority of isolated crystalline complexes show the so-called trans-III configuration (diastereoisomer RSSR, in Fig. 7), in which the metal is perfectly coplanar with the four amine nitrogen atoms, thus profiting from strong coordinative interactions.22 In principle, cyclam complexes can be considered as ideal subjects for the study of metal–anion interaction for two main reasons: (i) they offer plenty of space to the incoming anion along the z axis, and (ii) thanks to their inertness, they cannot be demetallated even on addition of a large excess of anion. In 1984, Hancock studied, through spectrophotometric titrations, the interaction of Cu(cyclam)2+ with anions in an aqueous solution (m = 0.1 M, 25 1C).23 Very weak interactions were observed and no reliable binding constants could be measured for F and Cl ions at the 102 M concentration scale (which indicates a binding constant log K o 1). Appreciable interaction was observed for polyatomic anions as N3 (log K = 2.1) and NCS (log K = 1.8). The intrinsically low magnitude of the binding constants has to be mainly ascribed to the endothermic anion dehydration, very high for monoatomic anions, less significant for N3 and NCS, in which the negative charge is spread over three atoms. Fig. 8 shows the molecular structure of the [CuII(cyclam)NCS]+,24 and [CuII(cyclam)N3]+,25 ternary complexes. Both complexes show a regular square pyramidal geometry, with the anion occupying the apical position. The CuII centre is only slightly displaced from the plane of four amine nitrogen atoms ([CuII(cyclam)NCS]+: 0.04 Å; [CuII(cyclam)N3]+: 0.16 Å). Cyclam can be permethylated at the amine nitrogen atoms to give 1,4,8,11-tetramethyl-1,4,8,11-tetraazacyclotetradecane

Fig. 7 The most frequently observed configurations of transition metal complexes of cyclam (trans-III) and Me4cyclam (trans-I).

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Fig. 8 Crystal structures of (a): [CuII(cyclam)NCS]+;24 displacement of CuII from the N4 plane: 0.04 Å, Cu–N(amine) distance 2.01 0.01 Å, Cu–N(NCS) distance 2.43 Å; (b) [CuII(cyclam)N3]+;25 displacement of CuII from the N4 plane: 0.16 Å, Cu–N(amine) distance 2.02 0.01 Å, Cu–N(N3) distance 2.31 Å. Hydrogens and counteranions omitted for clarity.

(Me4cyclam),26 which, on reaction with transition metal salts, gives complexes of trans-I configuration (R,S,R,S, see Fig. 7). The crystal and molecular structures of some divalent 3d metal complexes of Me4cyclam are shown in Fig. 9. All the complexes shows a square pyramidal geometry, with the four amine nitrogen atoms positioned at the corners of the square and an apically bound anion, as observed for CuII(cyclam)2+ complexes shown in Fig. 8. However, in the present cases, the metal lies well above the plane of nitrogen atoms, on the side of methyl substituents. In particular, the displacement from the N4

Fig. 9 The crystal and molecular structures of [MII(Me4cyclam)X]+ complexes (M = CuII, CoII, NiII), with relevant geometrical parameters (a): distance between the anion donor atom and the metal ion; (b) distance of the metal ion from the plane of the four amine nitrogen atoms; M–Nam, average distance between the metal ion and the nitrogen atoms of Me4cyclam. Not coordinated counter-anions and hydrogen atoms have been omitted for clarity. [CuII(Me4cyclam)NCS]+;27 [CoII(Me4cyclam)NCS]+;28 [NiII(Me4cyclam)N3]+;29 CoII(Me4cyclam)N3]+.30

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Tutorial Review plane is expected to favour the interaction with the anion. In this connection, one should observe that in the case of CoII ((ii) and (iv) in Fig. 9) and NiII complexes (iii) the metalanion distance is especially small, smaller than the average MIIN(amine) distance. The opposite occurs with the CuII complex (i), for which the length of the apical bond is larger than that of the CuIIN(amine) bond (see values reported in Fig. 9). On these bases, one should anticipate that CoII(Me4cyclam)2+ and NiII(Me4cyclam)2+ should be inclined to form with a given anion more stable ternary complexes than CuII(Me4cyclam)2+. The interaction of the rod-like triatomic anions N3, NCO and NCS with CoII, NiII and CuII complexes of Me4cyclam in water has been investigated by Kaden and Paoletti in 1982 through spectrophotometric titration experiments.31 Log K values for the equilibria of formation of ternary complexes are shown in Fig. 10a. As a general behaviour, CoII and NiII complexes form more stable ternary species than the CuII analogue, which may reflect the formation of more intense metal–anion interactions, as expected on the basis of structural data. Moreover, calorimetric studies showed that the stability of the ternary complexes of formula [MII(Me4cyclam)X]+ results from both favourable DHo and TDSo contributions.31 As an example, Fig. 10b shows the thermodynamic quantities for the equilibria involving the N3 anion. It appears that anion desolvation plays, also in the present systems, a determining role: (i) it makes the entropy change distinctly positive and (ii) it reduces the exothermicity associated the formation of the metal–anion coordinative bond. In particular, in the CuII(Me4cyclam)2+/N3 system, the exothermic effect of the CuIIN3 interaction and the endothermic effect of azide dehydration balance exactly, thus making complexation athermic and solely driven by the entropy term. Fig. 10b reports also, as green bars, the DGo values associated to the equilibria M2+ + NCO # [M(NCO)]+ in water.32 These values are of the same order of magnitude of those referring to Me4cyclam complexes. In particular, in the case of CoII and NiII, the interaction of the [MII(Me4cyclam)]2+ complex with NCO is favoured with respect to the corresponding aquaions, whereas for CuII the reverse behaviour is observed. Pre-coordination by Me4cyclam is especially

Fig. 10 (a) Log K values for the equilibria MII(Me4cyclam)2+ + X # [MII(Me4cyclam)X]+, in water at 25 1C (M = Co, Ni, Cu; X = N3, NCO, NCS); (b) thermodynamic quantities for the equilibria MII(Me4cyclam)2+ + NCO # [MII(Me4cyclam)NCO]+, in water at 25 1C.31

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Chem Soc Rev favourable, as it eliminates the endothermic dehydration of the uncomplexed aqueous ion and it prearranges the metal to the interaction with the anion. A comparison of the anion binding tendencies of metal complexes of cyclam and Me4cyclam can be made only in the case of copper(II): results are odd. In the case of NCS, it is CuII(Me4cyclam)2+ that forms the most stable complex: log K = 2.1 compared to 1.8 for CuII(cyclam)2+. On the other hand, for the N3 ion, the opposite trend is observed: CuII(Me4cyclam)2+ log K = 1.3, CuII(cyclam)2+ log K = 2.1. These subtle differences can be ascribed to a variety of factors, which include solvation effects, aniontetramine framework steric repulsions, binding mode of the anion (bent for N3, almost co-linear for NCS). Indeed, comparison of structural data available for [CuII(cyclam)NCS]+ and [CuII(Me4cyclam)NCS]+ complexes (crystal structures in Fig. 8 and 9, respectively) shows substantial differences in bonding parameters: for instance the axial CuIIN(NCS) bond is much shorter for the Me4cyclam complex than for the cyclam analogue (2.11 Å and 2.43 Å), in agreement with the higher value of the binding constant. In any case, Me4cyclam and cyclam complexes, as well as any other monometallic polyamine complex, do not form very stable ternary complexes with anions and do not provide any defined selectivity with respect to anion binding. In classical coordination chemistry, when a metal does not form a stable complex with a given unidentate ligand (e.g. ammonia), it is convenient to move to the corresponding bidentate ligand (e.g. ethylenediamine), in order to profit from the chelate effect.33 In the same way, in the case of anion binding, in order to have more stable ternary complexes, it seems convenient to profit from the chelate effect, thus designing polyamine complexes containing two (or even more) coordinatively unsaturated metal centres. As an additional benefit, steric constraints possibly present in the polyamine framework can generate a pronounced geometrical selectivity with respect to anion binding.

4. Two metals do it better: dicopper(II) complexes of TetraAminoEthylCyclam (TAEC) In 1986 Kida, Murase and co-workers reported the octadentate ligand 1,4,8,11-(2-aminoethyl)-1,4,8,11-tetraazacyclo-tetradecane (1, TAEC), in which four CH2CH2NH2 pendant arms have been appended to the four nitrogen atoms of cyclam.34 TAEC can coordinate two metal ions according to two different topological arrangements, as illustrated in Fig. 11. In the isomer I, each metal centre is coordinated by a tetramine subunit of type 2.3.2-tet, giving rise to a 5,6,5 sequence of chelate rings (the term p.q.r-tet refers to a linear tetramine of formula: NH2(CH2)pNH(CH2)q(CH2)rNH2). Among linear tetramines, 2.3.2-tet forms the most stable mononuclear complexes with transition metals (e.g. CuII), due to its ability to place its amine donor atoms at the corners of a square, according to a strain-free conformation.35 In particular, the metal ion lies on the plane of four nitrogen atoms. On the other hand, in the topological isomer II, each metal is bound by a tetramine subunit of the type of 2.2.2-tet (trien), with a 5,5,5

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Fig. 11 Topological isomerism of the dimetallic complexes of the octamine TAEC (1). The coordinating tendencies of the anion determine the formation of the isomer I (poorly coordinating anion X) or of the isomer II (strongly coordinating anion Y).

sequence of the chelate rings. As a general behaviour, in mononuclear trien complexes the tetramine fails to give a square planar coordination geometry and the metal centre (e.g. CuII) lies distinctly above the mean plane of four nitrogen atoms. As a consequence, CuII(trien)2+ is especially inclined to interact with anions to give five-coordinate ternary complexes of distorted square pyramidal geometry. Very interestingly, structural studies on crystalline dicopper(II) complexes of TAEC demonstrated that the formation of the topological isomers I and II is controlled by the nature of the anion: poorly coordinating anions X (e.g. ClO4) favour the formation of [CuII2(TAEC)X2]2+ complexes, according to the arrangement I; more coordinating anions Y (e.g. halides) promote the formation of the isomer II, in which Y bridges the two CuII centres: [CuII2(TAEC)Y]3+.34 The structural features of the two different ternary complexes are sketched in Fig. 11. Fig. 12a shows the molecular structure of the CuII2(TAEC)(ClO4)4 salt:36 one perchlorate oxygen atom is weakly coordinated to each metal centre (CuIIO distance: 2.56 Å). Each metal is moderately displaced from the mean N4 plane (0.19 Å) and the CuIICuII distance is 5.48 Å. Fig. 12b shows the molecular structure of the [CuII2(TAEC)Br]+ complex.36 The bromide ion bridges the two CuII centres, in the complex arranged in the topological mode II. Each metal is

Fig. 12 Crystal structure of (a) the [CuII2(TAEC)(ClO4)2]2+ complex (topological isomer I);36 (b) the [CuII2(TAEC)(Br]+ complex (topological isomer II).34 Hydrogen atoms have been omitted for clarity.

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displaced from the trien plane by 0.40 Å. The CuII–CuII distance is 4.66 Å. The binding tendencies of the dicopper(II) TAEC complex towards halides were investigated through spectrophotometric titration experiments in an aqueous solution made 0.03 M in NaClO4, at 25 1C. In particular, a solution of the CuII2(TAEC)(ClO4)4 salt was titrated with a standard solution of the chosen sodium salt. It is suggested that the dimetallic complex maintains in water the structural arrangement observed in the solid state and that the two weakly bound perchlorate ions are replaced by water molecules. Thus, the anion binding equilibrium can be represented as illustrated in Fig. 13. Notice that the progress of the equilibrium requires the occurrence of the configurational conversion of isomer I to isomer II, which probably involves a dissociative step. However, no time dependance of the spectra was observed and equilibrium was reached in the time scale of the titration experiment. Log K values for the investigated anions are reported in Table 1.37 Chloride and bromide form complexes of high and comparable stability, with K E 3 104, a value much higher than that estimated, for instance, for the interaction of Cl with CuII(cyclam)2+ (o102).23 Such an extra-stability should result (i) from the interaction with 2 metal centres (enthalpy contribution) and (ii) from the chelate effect (the anion binding by the second metal being strongly favoured, entropy contribution). Binding constant for I is 50-fold lower, which suggests that the cavity of the [CuII2(TAEC)]4+ receptor (isomer II) is rather rigid and suitable for the inclusion of Cl and Br. Iodide is probably too big to fit the cavity of the receptor relaxed to its minimum energy and its inclusion may induce an endothermic reorganisation of the polyamine framework. Such an effect is more evident in the case of the acetate ion: CH3COO is a truly ambidentate anion and should give a rather stable ternary complex. On the contrary, it shows the lowest binding constant (more than two orders of magnitude lower with respect to Cl and Br). The paradox is explained by the crystal structure of the [CuII2(TAEC)CH3COO]3+ complex, shown in Fig. 14a.39 The amplitude of receptor’s cavity has not been substantially expanded (the CuIICuII distance, 4.76 Å, is only slightly larger

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Fig. 13 Binding of an anion Y by the [CuII2(TAEC)]4+ complex in aqueous solution. It is hypothesised that a water molecule is weakly bound to each metal centre, to give a highly elongated square pyramidal coordination geometry. On anion coordination, a fast configurational conversion from I to II takes place, while each CuII centre assumes a more defined square pyramidal geometry (stronger apical interaction). Anion binding induces a red-shift of the absorption band of the CuII chromophore, from 560 nm to 620–630 nm (tentatively illustrated in the figure by a colour change from pink-violet to blue).

Table 1 Log K values of the equilibria: [CuII2(L)]4+ + Y # [CuII2(L)Y]3+, in aqueous solution (0.03 M NaClO4, 25 1C37

L

Cl

Br

I

CH3COO

N3

1, TAEC 2, TAEP 3, TAEH

4.47 3.22 2.57b

4.49 3.24 2.59b

3.87 2.76 2.19b

2.12 1.35

4.02a

a Measured in 0.5 M NaNO3;38 in the same medium log K for Cl = 3.57.38 b No added background electrolyte.

Fig. 14 The crystal structures of the ternary complexes: (a) [CuII2(TAEC)CH3COO]3+;39 (b) [CuII2(TAEC)N3]3+. Counter-anions and hydrogen atoms omitted for clarity.40

than that observed in the [CuII2(TAEC)Br]3+ complex, 4.66 Å). The CH3COO ion behaves as an ambidentate ligand, but the two CuII centres show a different coordinative arrangement: in one moiety, an acetate oxygen atom occupies a corner of the square of the pyramid, with a short CuO distance (1.99 Å), in the other, the acetate oxygen occupies the apical position of the pyramid, at a long CuO distance (2.22 Å). The asymmetry of the coordination modes seems to reflect the presence of serious steric constraints in the polyamine framework. The interaction of [CuII2(TAEC)]4+ with Cl, Br and N3 in water was also investigated by Hancock and Murase under different conditions (0.5 NaNO3, 25 1C).38 Log K values for halides were appreciably lower than those measured in the presence of a less concentrated background electrolyte, and, rather surprisingly, it was observed that N3 formed a more stable ternary complex than the most stable complex among halides, i.e. Cl (N3 4.0; Cl 3.6). Azide typically forms stable ternary complexes when it bridges the two metal centres as an ambidentate ligand (end-to-end). However, such a binding

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mode does not seem compatible with the rigid arrangement of TAEC complexes. Again, the crystal structure helped to solve the enigma (see Fig. 14b). The [CuII2(TAEC)N3]3+ complex shows configuration II, with a bridging N3 ion.40 However, azide is end-on bound, with one terminal nitrogen atom bridging the two metals. This type of coordination should not induce any deformation of the polyamine framework. In particular, both CuII centres show a nearly equivalent square pyramidal coordinative arrangement and are displaced by 0.45 Å from the plane of four amine nitrogen atoms.

Structural modifications were then made on the skeleton of TAEC, which consists of a 14-membered tetra-aza macrocycle (cyclam), functionalised with four aminoethyl side-arms. In particular, the atomicity of the macrocyclic ring was expanded to 15 (2, TAEP) and to 16 (3, TAEH).41 In the topological isomer II, TAEP hosts one metal centre in a 2.2.2-tet subunit and the other in a 2.3.2-tet subunit. With TAEH, both metals are coordinated by 2.3.2-tet moieties. Log K values in Table 1 indicate that the affinity towards halide ions decreases along the series [CuII2(TAEC)]4+ 4 [CuII2(TAEP)]4+ 4 [CuII2(TAEH)]4+. A satisfactory explanation of this trend is not straightforward, in view of the availability of a reduced set of crystallographic data. An interesting comparison, however, can be made for the structures of the two complex salts [CuII2(TAEH)ClO4](ClO4)3 and [CuII2(TAEH)N3](ClO4)3, shown in Fig. 15. In contrast to what observed with TAEC (complex in Fig. 12a, topological isomer I), the TAEH complex containing all ClO4 anions, Fig. 15a, shows an arrangement of type II.41 With TAEH, in fact, such a configuration guarantees both (i) the more thermodynamically favourable coordination of the two CuII ions by 2.3.2-tet subunits, and (ii) a larger separation of the

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Fig. 15 (a) Crystal structure of [CuII2(TAEH)ClO4](ClO4)3:41 only one ClO4 is shown, very weakly coordinated to the two CuII centres with one bridging oxygen atom (CuO distance: 3.11 Å); the CuCu distance is 5.16Å, each CuII ion is displaced from the N4 plane by 0.12 Å; (b) crystal structure of [CuII2(TAEH)N3](ClO4)3:41 N3 bridges the two CuII ions according to an end-toend mode, the CuCu distance is 5.44 Å, the metal ions are displaced from the N4 plane by 0.36 and by 0.32 Å.

two compartments (by trimethylenic spacers instead of ethylenic spacers). Notice that a ClO4 ion points an oxygen atom towards the cavity of the complex, but it should not give any important stabilising contribution, in view of the very long CuO distances (3.11 Å). Thus, that observed in the [CuII2(TAEH)ClO4]3+ complex can be assumed as the most stable conformation, showing, in particular, a CuCu distance of 5.16 Å. Fig. 15b displays the structure of the [CuII2(TAEH)N3]3+ ternary complex.41 In this case, the azide ion bridges the two metal centres according to an end-to-end bridging mode, a type of coordination allowed by a cavity larger than that available in the m-azido TAEC derivative. However, azide coordination seems to involve some conformational cost, as suggested by the CuCu distance (5.44 Å), distinctly larger than in the perchlorate complex in Fig. 15a. Disappointingly, no equilibrium studies have been carried out on the interaction of [CuII2(TAEH)]4+ with N3 and other polyatomic anions. In any case, the studies of Kida and co-workers on tetra-aminoethyl tetra-aza macrocycles have demonstrated that receptors containing two coordinatively unsaturated metal centres can effectively bind anions, even in a competing medium like water. Prepositioning of the two metals into a designed polyamine framework can open the way to anion selective recognition.

Tutorial Review Lehn predicted a variety of useful processes which could take place inside the cavity of dimetallic cryptates: fixation and transport, multicentre and multi-electronic catalysis, including N2 and O2 reduction, water splitting, mimicking of metalloproteins.42 More than three decades later, we can observe that investigated dimetallic cryptates have hardly performed the expected functions.43 With one nice exception: anion recognition. Examples in the following will illustrate how anion inclusion into dimetallic cryptates can be selectively modulated by varying the length and nature of the spacers connecting the two tren subunits. However, it seems convenient, at this stage, a brief reminder of the coordinating tendencies of the tripodal tetramine tren. If linear and cyclic tetramines favour the formation of ternary metal complexes of square pyramidal geometry, tren promotes the attainment of a trigonal bipyramidal coordination. In particular, the tertiary amine nitrogen atom occupies one of the axial positions of the trigonal bipyramid, the three primary amine groups stay in the equatorial sites and the other axial position is left available for an anion X or for a solvent molecule. Copper(II) is especially prone to the formation of five-coordinate complexes. The crystal structure of the [CuII(tren)NCS]+ complex is shown in Fig. 16a.44 A distinctive feature is that equatorial CuIIN bonds (2.08 0.06) are on average longer than axial CuIIN bonds (CuII–Ntert 2.04 Å), disclosing an axially compressed trigonal bipyramid. In particular, the CuIIN(NCS) distance is especially small (1.95 Å). The situation is opposite to what typically observed with linear tetramine and tetraaza macrocyclic analogues. As an example, the structure of the [CuII(trien)NCS]+ complex is shown in Fig. 16b.45 The complex shows an axially elongated square pyramidal geometry (CuII–N(amine) 2.03 0.05 Å), with an especially large CuII–N(NCS) distance (2.19 Å). These structural differences are quite general and suggest that CuII-tren2+ is a more appropriate subunit for anion binding than any other copper(II) complex with linear and cyclic tetramines. The coordinating tendencies of the bistren derivative 4 towards divalent 3d metal ions were investigated in water through pH titration experiments by Lehn and Martell.46 It was found that CuII formed a dinuclear complex very stable (log K = 27.3 for the equilibrium 2Cu2+ + 4 # [CuII2(4)]4+ in 0.1 M NaClO4), at least ten orders of magnitude more stable than dinuclear complexes of

5. The age of bistren cryptates In 1977 Lehn reported the synthesis of the cage-like octamine 4, in which two tren subunits have been linked by three –CH2CH2OCH2CH2 spacers.42 The polyamine was called ‘bistren cryptand’, in analogy to the structurally similar poly-ethereal systems, suitable for the selective inclusion of alkali and alkalineearth metal ions, to form mono-nuclear cryptates. It was observed that 4 was able to include two transition metals, e.g. CuII, in a stepwise mode, to give a di-nuclear cryptate. Lehn anticipated that, due to the coordinative unsaturation of the two metal centres, dimetallic bistren cryptates could encapsulate a variety of substrates, including anions. Indeed, preliminary spectroscopic studies (EPR and UV-vis) provided evidence for the inclusion of CN and N3 into the [CuII2(4)]4+ complex.42

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Fig. 16 Crystal structures of the ternary complexes of CuII and NCS with the branched tetramine tren (a),44 displaying a compressed trigonal bipyramidal geometry, and with the linear tetramine trien (b),45 showing an axially elongated square pyramidal geometry. Hydrogen atoms and uncoordinated counter-anions have been omitted for clarity.

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Fig. 17 The crystal structure of the complex [CuII2(4)OH]3+.47 The CuIICuII distance is 3.77 Å. The bridging OH ion points towards the ethereal oxygen atom of one of the three spacers and establishes a hydrogen bonding interaction.

CoII, NiII and ZnII. Moreover, it was observed that [CuII2(4)]4+ displays a huge affinity for the hydroxide ion, to give the [CuII2(4)OH]3+ complex. In particular, the constant for the [CuII2(4)]4+ + OH # [CuII2(4)OH]3+ equilibrium is more than five orders of magnitude larger than that observed for the reference equilibrium [CuII(tren)]2+ + OH # [CuII(tren)OH]+: log K = 9.7 and 4.8, respectively. An inspection of the crystal structure of the [CuII2(4)OH]3+ complex, in Fig. 17, provides a convincing account for such an extra-stability.47 In fact, the OH ion bridges the two CuII ions, but it is not co-linear with the two metal centres. It protrudes towards the ethereal oxygen atom of one of the three –CH2CH2OCH2CH2 spacers (CuIIOCuII angle 1551), in order to establish a H-bond interaction. It is probably this additional contribution, which can be referred as a second-sphere coordination, that

Scheme 1 Bistren derivatives with spacers providing ellipsoidal cavities with different major axes.

Scheme 2

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Synthesis of bistren derivatives. The amplitude of the cavity is determined by the nature of the dialdehyde.48

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Fig. 18 (a) Family of spectra taken over the course of the titration of an aqueous solution of [CuII2(5)]4+, buffered to pH = 8, with a standard solution of NaN3: anion addition promotes the development of an azide-to-metal charge transfer band; (b) symbols: molar absorbance at 400 nm; fitting line calculated on assuming the occurrence of the equilibrium [CuII2(5)]4+ + N3 # [CuII2(5)N3]3+, with log K = 4.75 0.05.49

Fig. 19 Crystal structures of: (a) [CuII2(5)N3]3+,50 the N3 ion is co-linearly coordinated to the two CuII centres (‘unnatural’ mode); (b) 5 alone;51 (c) half-cage complex [CuII(Me3tren)N3]3+:52 in the absence of steric constraints, N3 displays its ‘natural’ bent coordinating geometry.

5:51 its conformation is not exactly the same observed in the [CuII2(5)N3]3+ complex. For instance the distance between the two tertiary amine nitrogen atoms in the ligand alone (10.93 Å) is appreciably larger than that observed in the [CuII2(4)N3]3+ complex (10.09 Å): that means that stepwise inclusion of 2 CuII ions, then of N3, takes places with a contraction of the cavity, probably with some conformational cost. Another point of interest is provided by the mode of coordination of the N3 anion. In the absence of steric constraints, for instance in mononuclear copper(II) tetramine complexes, the coordinated azide is ‘bent’, i.e. it forms a CuNN angle ranging between 1201 and 1401. As an example, the structure of the [CuII(Me3tren)N3]3+ half-cage complex is shown in Fig. 19c (Me3tren = tris(2-(N-methylamino)ethyl)amine).52 The CuIINN angle of 1281 essentially reflects the sp3 hybridisation of the coordinated terminal nitrogen atom. On the other hand, encapsulation into the bistren dicopper((II) complex forces N3 to the co-linear coordination. This might represent, in principle, a further unfavourable contribution to the formation of the ternary complex. It will be shown in the following that, on the contrary, the two CuII ions inside the bistren framework appreciate such a type of ‘unnatural’ coordination. Analogous titration experiments were carried out with a variety of ambidentate polyatomic anions: in any case, on anion addition, the [CuII2(5)]4+ pale blue solution turned intense blue or green, and a new band or a distinct shoulder was observed to

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develop over the course of the titration.49 Absorbance vs. equiv. of X anion profiles indicated in all cases the formation of a 1 : 1 inclusion complex, [CuII2(5)X]3+, and the corresponding equilibrium constant (at pH = 8) could be determined. The trend of log K values could not be rationalised in terms of anion shape, electrical charge or position in the spectrochemical series. A convincing key of interpretation was provided by the anion bite length, i.e. the distance between two consecutive donor atoms of the anion acting as a bidentate ligand (Fig. 20). In particular, when plotting log K values vs. anion bite length, a sharp peak selectivity was observed, in favour of the azide ion. Thus, it appears that the N3 ion has the right length to place its donor atoms (terminal nitrogens) in the two axial positions left available by the two CuII centres, without inducing any endergonic conformational rearrangement of the receptor’s framework. Anions of either larger (i.e. NCS) or smaller bite length (e.g. NO3) force the cryptate to leave its relaxed conformation and to adjust the distance between the two CuII centres at the required value: the more or less significant rearrangement of the cryptand framework (involving either expansion or contraction) is reflected in a concurrent decrease of the log K value. Thus, the [CuII2(5)]4+ receptor neither recognises the shape of the anion nor its electrical charge: it recognises its bite length.

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Fig. 21 The crystal structure of the cryptate [CuII2(6)(H2O)2]4+. The two labile axially coordinated water molecules can be replaced in solution by the oxygen atoms of dicarboxylates, providing linear recognition.53

Fig. 20 Bite length of selected anions: distance between two consecutive donor atoms of an anion X acting as an ambidentate ligand. Peak selectivity for the stability of the [CuII2(5)]3+ complex with respect to the bite distance of the anion X (log K of equilibria: [CuII2(5)]4+ + X # [CuII2(5)X]3+, at 25 1C, in a solution buffered to pH = 8.49

The sharp and definite selectivity displayed by the [CuII2(5)]4+ system as an anion receptor seems to be ascribed to the scarce deformability of the bistren framework, which, in turn, may be associated with the rigidity of the 1,3-xylyl spacers joining the two tren subunits of cryptand 5. As a consequence, the [CuII2(5)]4+ cryptate, at any pH, is not able to include monoatomic anions like halides, which can act as ambidentate ligands, but are too small to bridge the two metal centres. As a matter of fact, addition of even a large excess of Cl, or of other halide ions, to a solution of [CuII2(5)]4+, at any pH, neither induces a colour change nor a modification of the absorption spectrum.

6. A bistren cryptate exerting linear recognition of dicarboxylates Bistren cryptates offer an ellipsoidal cavity suitable for the inclusion of polyatomic anions. The major axis of the ellipsoid (more precisely, a prolate spheroid, i.e. a rotation ellipsoid with two equal minor axes and one major axis) can be varied at will through a thoughtful choice of the dialdehyde to be used in the Schiff base condensation with tren. As an example, the major axis of the ellipsoid can be extended by inserting 4,4 0 -ditolyl spacers (6). The structure of the dicopper(II) complex of 6 has been determined (see Fig. 21).53 The solid complex [CuII2(6)(H2O)2](NO3)4 contains two water molecules, each one occupying the vacant axial position of each CuII centre. It is suggested that the [CuII2(6)(H2O)2]4+ species is present also in an aqueous solution, where the two labile water molecules can be replaced by the donor atoms of an ambidentate anion. The length of the ellipsoidal cavity seemed appropriate for the inclusion of dicarboxylates, both aromatic and aliphatic. Quite disappointingly, the replacement of the water oxygen atom by the carboxylate oxygen atom did not cause any significant spectral modification, also in view of the diluted conditions imposed by the poor solubility of the cryptate salt. Thus, anion inclusion equilibria

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had to be investigated through spectrofluorimetry, by using the method of indicator displacement.54,55 According to this paradigm, to a solution of the anion receptor R, a substoichiometric amount of the fluorescent indicator In is added. An RIn complex forms, in which an electron/energy transfer process takes place, so that In fluorescence is quenched. Then, the envisaged anion X is added and displaces In from R, due to either its greater affinity or to a mass effect or both. When released to the solution, In displays its natural fluorescence, thus signalling the formation of the stable complex RX and anion recognition. The process is pictorially illustrated in Fig. 22. Rhodamine (8) is a dye, which, when excited at 496 nm (isosbestic point), emits at 571 nm (orange fluorescence). It contains a 1,4-benzene-dicarboxylate subunit, suitable for bridging the two CuII centres of the [CuII2(6)]4+ receptor.

Preliminary experiments showed that, on interaction with [CuII2(6)]4+, the fluorescence of rhodamine is quenched, which suggests (i) inclusion of the dicarboxylate subunit into the cryptate and (ii) occurrence of an intra-complex energy or electron transfer process involving the excited fluorophore and the CuII centre(s). A log K = 7.0 0.2 was determined for the equilibrium: [CuII2(6)]4+ + In # [CuII2(6)In]3+. Then, in a typical displacement experiment, a non-fluorescent solution

Fig. 22 The fluorescent indicator displacement paradigm. The receptor quenches the emission of the indicator through an intracomplex mechanism. When displaced by the anion, the indicator releases its fluorescence, thus signalling anion recognition.55

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Fig. 23 (a) Variation of the fluorescence emission of a solution 2.5 107 M in rhodamine and 2.5 106 M in [CuII2(6)]4+, buffered to pH = 6, when titrated with a solution of 1,n-benzenedicarboxylate (n = 2, phthalate; n = 3, isophthalate; n = 4, terephthalate; (b) bars, left vertical axis: log K values for the equilibria [CuII2(6)]4+ + O-O2 # [CuII2(6) O-O]2+ (O-O2: aromatic dicarboxylate; horizontal dashed line: log K value for the equilibrium: [CuII2(6)]4+ + In # [CuII2(6)In]3+ (In: rhodamine), left vertical axis; diamonds: D, difference (absolute value in Å) between the distance of the oxygen atoms of the coordinated water molecules in the [CuII2(6)(H2O)2]4+ complex and the two oxygen atoms of the dicarboxylate, right vertical axis; the dashed line is the least-squares straight-line.53

2.5 106 M in [CuII2(6)]4+ and 2.5 107 M in rhodamine, buffered to pH = 7 was titrated with a solution of terephthalate (1,4-benzene-dicarboxylate). On dicarboxylate addition, the fluorescence of rhodamine was progressively restored, due to indicator displacement and anion encapsulation. Fig. 23a shows the corresponding titration profile. It is observed that rhodamine fluorescence is restored on first additions of terephthalate, signalling quick replacement and formation of a stable receptor–anion complex. In particular, total extrusion of the indicator is achieved on addition of 1 equiv. of terephthalate (with respect to the receptor [CuII2(6)]4+). On the other hand, on titration with isophthalate, a competitive behaviour is observed, with incomplete restoring of rhodamine fluorescence even after the addition of a ten-fold excess of anion. Finally, phthalate is totally unable to displace the indicator. On treatment of titration data, the equilibrium constants for the interaction of [CuII2(6)]4+ with positional isomers of phthalate were determined and their values are reported in the bar diagram in Fig. 23b. The diagram illustrates well the paradigm of indicator displacement. The envisaged analyte, terephthalate, exhibits a distinctly larger constant than the indicator, which, in turn, shows a constant larger than the interferring analytes, isophthalate and phthalate. Selectivity in favour of terephthalate has a geometric nature, as inferred form available structural data. In the crystallographically investigated [CuII2(6)(H2O)2]4+ complex, the distance between the coordinated water oxygen atoms is 7.36 Å. This can be considered the ‘ideal’ distance which a dicarboxylate ion should span with two oxygen atoms, in order to form a stable inclusion complex. Terephthalate possesses two potentially donating oxygen atoms at a distance of 7.39 Å, as judged from the crystal structure of its alkaline metal salts, and can therefore replace the two coordinated water molecules without inducing any conformational rearrangement of the bistren framework. Rhodamine contains the same 1,4benzene-dicarboxylate fragment as terephthalate, but it forms with [CuII2(6)]4+ a complex one order of magnitude less stable than plain terephthalate. This may result from the steric repulsive effects

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exerted by the fluorogenic substituent. Isophthalate and, in particular, phthalate form much less stable complex than terephthalate, because their most favourable OO distances (4.98 and 3.16 Å, respectively) strongly deviate from the ideal value of 7.36 Å. Quite interestingly, an inverse linear relationship correlates log K and the absolute value of the difference between the ideal OO distance and the most favourable one observed in any isolated aromatic dicarboxylate (diamonds in Fig. 23b): the smaller the difference, the more stable the receptor–dicarboxylate complex. This study was then extended to linear aliphatic dicarboxylates of formula OOC(CH2)nCOO (n = 2, succinate; n = 3, glutarate; n = 4, adipate; n = 5, pimelate). Fig. 24a displays the change of the fluorescent intensity of rhodamine, obtained on titration of a solution 2.5 106 M in [CuII2(6)]4+ and 2.5 107 M in rhodamine, buffered to pH = 7, with a standard solution of dicarboxylate. Glutarate (n = 3) and adipate (n = 4) displace the indicator and revive fluorescence, whereas succinate (n = 2) and pimelate (n = 5) do not, confirming that the [CuII2(6)]4+ receptor is able to exert anion length recognition. Pertinent log K values for association equilibria are reported in the bar diagram in Fig. 24b. Glutarate and adipate have similar binding constants, in spite of the fact that their most favourable OO distances, as observed in their alkaline metal salts, differ significantly from each other (glutarate: 7.20 Å; adipate 7.83) and are appreciably different from the ‘ideal’ one. It is suggested that both anions can span the vacant positions inside the receptor at a low conformational cost, in view of the flexible nature of their spacers. They show a log K value similar to that of rhodamine and, in titration experiments, they are able to displace the indicator due to a mass effect. On the other hand, succinate and pimelate are one too short and the other too long (favourable OO distance: 6.04 and 9.67), and their inclusion in the receptor must involve a drastic rearrangement of the bistren framework. In any case, glutarate and adipate show a binding constant one order of magnitude lower than terephthalate, which may reflect the loss of entropy they can experience in complexation with respect to the rigid aromatic dicarboxylate.

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Fig. 24 (a) Variation of the fluorescence emission of a solution 2.5 107 M in rhodamine and 2.5 106 M in [CuII2(6)]4+, buffered to pH = 7, when titrated with linear aliphatic dicarboxylates of formula OOC–(CH2)n-COO (n = 2, succinate; n = 3, glutarate; n = 4, adipate; n = 5, pimelate); (b) bars, left vertical axis: log K values for the equilibria [CuII2(6)]4+ + O-O2 # [CuII2(6) O-O]2+ (O-O2: aliphatic dicarboxylate); horizontal dashed line: log K value for the equilibrium: [CuII2(6)]4+ + In # [CuII2(6)In]3+ (In: rhodamine), left vertical axis; diamonds: D, difference (absolute value in Å) between the distance of the oxygen atoms of the coordinated water molecules in the [CuII2(6)(H2O)2]4+ complex and the two oxygen atoms of the dicarboxylate, right vertical axis.53

Quite interestingly, [CuII2(6)]4+ forms a stable complex, with L-glutamate, which presents the same carbon skeleton of glutarate.56 L-Glutamate is an important neurotransmitter and [CuII2(6)]4+ recognises it in the presence of any other neurotransmitter, including aspartate and GABA.

7. A metallocyclam trifurcate receptor for tricarboxylates Recognition of tri-carboxylates, e.g. citrate, requires a more complex receptor’s design. A successful example refers to the trinuclear metal complexes of the tris-cyclam ligand 7.57 In the absence of any coordinating anion, it is expected that a trinuclear complex [MII3(7)]6+ displays an unfolded structure, in which the three metallocyclam subunits stay as far as possible each other, in order to minimise electrostatic repulsions between metal centres. Indeed, such an arrangement has been observed in the crystalline complex salt [NiII3(7a)](ClO4)6 H2O, shown in Fig. 25.58 The six ClO4 ions (not shown in the figure) are not coordinated to the diamagnetic NiII centres, which show a square geometry. It has been suggested that the complex [MII3(7)]6+ can rearrange to assume a bowl conformation at a moderate energy cost, in order to bind a terdentate anion, e.g. a tri-carboxylate, as roughly sketched in Scheme 3. In particular, the anion should be able to establish a binding interaction with each metal centre. Such a possibility has been verified with the complex [CuII3(7)]6+. It has been shown in previous sections that the CuII(cyclam)2+ fragment is coordinatively unsaturated and is especially inclined to bind an anion to give a five-coordinate square pyramidal ternary species. In order to verify the binding tendencies of [CuII3(7)]6+ with multidentate anions, a non-fluorescent solution 2 105 M in [CuII3(7)]6+ and 5 107 M in 5-carboxyfluorescein (9), buffered to pH = 7, was titrated with standard

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Fig. 25 The crystal structure of the complex salt [NiII3(7a)](ClO4)6H2O.58 Non-coordinating ClO4 ions and water molecule have been omitted for clarity, as well as hydrogen atoms. Each low-spin NiII centre shows a square coordination geometry.

solutions of 1,3,5-benzene-tricarboxylate, 1,3-benzene-dicarboxylate (isophthalate) and benzoate. Only the terdentate anion 1,3,5-benzene-tricarboxylate was able to displace the indicator, restoring fluorescein emission (see the titration profile in Fig. 26), whereas the dicarboxylate and the monocarboxylate derivative failed, even if added in a large excess. This strongly suggests that each –COO group of 1,3,5-benzene-tricarboxylate establishes a coordinative interaction with a CuII centre and that this energy contribution compensates the energy to be spent when moving from the unfolded conformation of [CuII3(7)]6+ (reasonably similar to that observed in the crystal structure of Fig. 25) to the bowl-shaped arrangement, capable of hosting the tri-carboxylate anion and providing full coordination. The tri-carboxylate anion citrate, 10, an analyte of multipurpose interest in food science, stimulated the design of selective receptors and sensors. The tripodal trication 11, in which a 1,3,5-triethylbenzene subunit has been armed with

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Scheme 3

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The interaction of a tris-metallocyclam receptor with a multidentate anion (e.g. a tri-carboxylate) capable of coordinating each metal centre.

Fig. 26 Fluorescence profiles recorded over the course of the titration with an aromatic carboxylate of an aqueous solution 2 105 M in [CuII3(7)]6+ and 5 107 M in 5-carboxy-fluorescein (9), and buffered to pH = 7.57

three methylguanidinium side chains, forms with citrate a 1 : 1 complex, with an association constant log K = 3.8 and has been used to determine citrate in a variety of beverages.59 Affinity toward citrate has been later improved by linking to the aromatic platform side chains containing the guanidiumcarbonyl pyrrole functionality (12, log K = 4.9 at pH = 7).60 The trimetallic receptor [CuII3(7)]6+ works much better, as it forms with citrate a 1 : 1 complex with an association constant log K = 5.6 at pH = 7, which is only slightly lower than that measured for 1,3,5-benzenetricarboxylate (log K = 5.8). Moreover, [CuII3(7)]6+ is capable of

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discriminating citrate from competing analytes present in beverages, as shown by the titration profiles displayed in Fig. 27.

8. Nitro–nitrito linkage isomerism in [NiII(diamine)2X2] complexes in solid state and in solution There exists a particular type of isomerism that is specifically related to the metal–anion interaction: linkage isomerism.

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Fig. 27 Fluorescence profiles recorded over the course of the titration of an aqueous solution 2 105 M in [CuII3(7)]6+ and 5 107 M in 5-carboxyfluorescein (9), and buffered to pH = 7, with a variety of interferents in beverages.

¨rgensen isolated the red and the yellow forms of a In 1894 Jo cobalt(III) ammonia compound, to which the same elemental analysis corresponded: CoCl2(NO2)6NH3.61 In 1907, Werner,62 within the frame of the theory of coordination compounds, assigned to the two complex salts the formula: [CoIII(NH3)5(NO2)]Cl2 and demonstrated that in the red form the acidic residues NO2 were bound to the metal through the oxygen atom (today we say: nitrito complex) while in the yellow form they were bound through the nitrogen atom (nitro complex). Since then, the nitro–nitrito interconversion in solution for complexes of this type has been investigated in detail, a process made complicated by the inertness of the CoIII centre.63 In the 1960s Goodgame and Hitchman discovered that highspin nickel(II) bis-diamine complexes could give rise to nitro–nitrito isomerism.64,65 These complexes exhibit a transoctahedral geometry, as sketched in Fig. 28. The nitro group is a stronger ligand than the nitrito one: it is higher in the spectrochemical series and it is able to accept electrons from the metal in a p mode. Thus, these complexes

Fig. 28 Nitro- and nitrito- coordination modes in high-spin trans-octahedral [NiII(diamine)2X2] complexes. In the absence of steric crowding, nitro coordination is favoured. In the presence of hindering substituents on the amine nitrogen atoms, the less space demanding nitrito coordination takes place.

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Chem Soc Rev tend naturally to assume nitro coordination: [NiIIL2(NO2)2], with L = plain ethylenediamine, N-alkyl derivatives). However, nitro coordination is more space demanding and, in the presence of steric crowding due to substituents on the nitrogen or carbon atoms of the diamine (e.g. L = N,N 0 - or N,N-diethylethylenediamine), nitrito coordination is favoured: [NiIIL2(ONO)2]. Formation of either isomer can be visually perceived, as the nitro complexes show a pink-red colour (which is associated to a d–d band centred at around 500 nm, as measured in reflectance in solid complexes or in absorbance in solution of a noncoordinating solvent, e.g. CHCl3). On the other hand, nitrito complexes exhibit a blue colour (d–d band centred at 580–600 nm). An interesting case is provided by the complexes with the two isomeric diamines diMeeN and adiMeen. The [NiII(diMeen)2(NO2)2] complex shows nitro coordination (see the crystal structure in Fig. 29a).66 On the other hand, the complex with the asymmetrically substituted dimethyl-ethylenediamine shows nitrito coordination, [NiII(adiMeen)2(ONO)2] (Fig. 29b).66 Such a behaviour would indicate that the two methyl groups exert a stronger repulsive effect (or leave less space for the anion) when they are both on the same nitrogen atom than when they are allocated each one on one of the two nitrogen atoms of the diamine. Fig. 30 shows the spectra taken on the solid complexes, the red [NiII(diMeen)2(NO2)2] and the blue [NiII(adiMeen)2(ONO)2]. Noticeably, the two solid isomeric complexes display a very different thermal behavior. In fact, on heating the solid [NiII(adiMeen)2(ONO)2] blue complex, neither color change nor spectral modifications are observed. On the contrary, on heating the solid [NiII(diMeen)2(NO2)2] complex, the colour gradually changes from pink-red to blue and definite spectral changes are observed.67 In particular, Fig. 31a shows the spectra of the solid NiII(diMeen)2(NO2)2, spread on filter paper, taken in absorbance, over the 25140 1C temperature interval. On heating, the band of the nitro complex, centred 500 nm, decreases in intensity, while a new band develops at 580 nm (nitrito isomer), indicating that the nitro form smoothly converts to the nitrito derivative. However, at the highest investigated temperature, 140 1C (above which supporting paper begins to char), the conversion is not complete. On making the rough assumption that the limiting spectrum of the [NiII(diMeen)2(ONO)2] complex is similar to that observed for the [NiII(adiMeen)2(ONO)2] complex

Fig. 29 The crystal structures of: (a) [NiII(diMeen)2(NO2)2, nitro coordination;66 (b) [NiII(adiMeen)2(ONO)2, nitrito coordination.66 Hydrogen atoms have been omitted for clarity.

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spectral features are not altered on heating the solution. On the other hand, a solution in CHCl3 of a nitrito complex, e.g. [NiII(N,N 0 -diethylethylenediamine)2(ONO)2], shows the absorption bands of both nitrito (590 nm) and nitro form (500 nm). Moreover, on increasing temperature, the intensity of the band of the nitrito form increases, while that of the band of the nitro complex decreases. The presence of a sharp isosbestic point at 570 nm indicated the occurrence of the endothermic nitro-tonitrito interconversion equilibrium: [NiIIL2(NO2)2] # [NiIIL2(ONO)2]

Fig. 30 Spectra, taken in absorbance, of the solid complexes [NiII(adiMeen)2(ONO)2], blue colour, and [NiII(diMeen)2(NO2)2], red colour, spread on filter paper, at 25 1C.

(Fig. 30), conversion should have taken place to ca. 50%. The presence of defined isosbestic points at 450 and 540 nm suggests that only two species coexist in the solid state equilibrium, thus indicating that the two NO2 ions, on heating, make a simultaneous pirouette, moving from the nitro to the nitrito manner of coordination and no mixed species (e.g. [NiII(diMeen)2(NO2)(ONO)]) form as an intermediate, during the conversion process. On cooling, the band of the nitrito species decreases, while the band of the nitro form is restored. The reversible thermochromic change shows a very moderate hysteresis, as shown by the diagram in Fig. 31b, which confirms the kinetic simplicity of the process. Studies on the nitro–nitrito interconversion of nickel(II) Nalkyl substituted ethylenediamine complexes in solution had been carried out by Goodgame and Hitchman well before the investigations in the solid state.65 In particular, it was observed (i) that solid nitro complexes, e.g. [NiII(N-ethylethylenediamine)2(NO2)2], when dissolved in a non-dissociating solvent, e.g. CHCl3, maintain nitro coordination and (ii) that red colour and

(4)

for which the following thermodynamic quantities were determined: DGo = 0.1 kcal mol1, DHo = 2.3 kcal mol1, TDSo = 2.2 kcal mol1 ( 0.5 kcal mol1). It appears that in the investigated complex nitro and nitrito forms exhibit a similar stability, which results from the balance of the enthalpy term, favouring nitro, and of the entropy term, favouring nitrito. The endothermicity of the nitro-to-nitrito conversion (4) is to be associated with the lower strength of the NiIIONO interactions with respect to the NiIINO2 interactions. On the other hand, the favourable DSo term reflects the higher mobility of the NO2 anion when coordinated through the oxygen atom, which involves: (i) the NO2 rotation along the NiIION axis (easier and less sterically hindered than the rotation of the N-bound anion), (ii) swapping of the coordinated and uncoordinated oxygen atoms of each NO2 ion, according to an oscillatory motion (a movement not possible in the nitro analogue). Nitro- and nitrito- bis-ethylenediamine complexes of nickel(II) disclosed a unique example of reversible thermochromic behaviour both in the solid state and in solution. Quite interestingly, such an appealing phenomenon can be designed and controlled through simple modifications of the diamine framework.

9. Epilogue Chemistry, as a young discipline, is compulsorily determined to open new routes and make new inventions (stimulated in that by the needs of an evolving society). Thus, it has no much time

Fig. 31 (a) Spectra of the solid NiII(N,N0 -dimethylethylenediamine)2(NO2)2 complex, spread on filter paper, taken in absorbance over the temperature range 25132 1C. On heating, the band at 500 nm (nitro) decreases, whereas the band at 580 nm (nitrito) increases; (b) heating (red symbols) and cooling (blue symbols) profiles: the vertical axis reports the difference between current absorbance at 610 nm (pertinent to the band of the nitrito form) and the absorbance at 610 nm, measured at 251(see the black double-head arrow set at 610 nm, in Fig. 31a).67

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Tutorial Review to dedicate to a revisitation of the past. However, a reconsideration of the studies carried out a century ago or more can help to develop and strengthen modern lines of research or to solve problems encountered in today’s chemistry. An example is provided by anion recognition, as we tried to illustrate in the present review. Anion recognition (and sensing) is an actively practised topic of supramolecular chemistry, with important fall-outs in medicine, food and environmental sciences.68 Supramolecular chemistry is the discipline studying non-covalent interactions, and anion recognition typically refers to the establishing of weak and reversible interactions between the anion and a given receptor. The first deliberately designed anion receptors, reported in the late 1970s and the early 1980s, were cage-shaped polyammonium,69 and polyalkylammonium compounds,70 capable of including anionic substrates in force of electrostatic and/or hydrogen bonding interactions. A high preorganisation of receptor’s framework (a cage) was required in order to impart stability to the inclusion complex and to compensate the strongly endergonic term due to anion desolvation. Later, a large variety of neutral receptors containing polarised NH fragments (from amides, ureas, thioureas, pyrroles)71 were introduced, capable of interacting with the anions only through hydrogen bonding. However, in most cases the anion–receptor interaction was too weak to compensate dehydration energy and recognition studies had to be carried out in non-protic media (CHCl3, MeCN, DMSO, in order of increasing polarity). On the other hand, transition metal ions establish with anions strong and reversible interactions and form stable complexes in water. In order to control and to address the interaction, it is opportune to pre-coordinate the metal ion with a polyamine, in such a way that one binding site is left vacant. Very conveniently, pre-coordination by the polyamine maintains metal affinity for anions intact. Then, selectivity towards anions can be improved through the design of the polyamine framework, in particular imposing appropriate geometrical constraints (thus playing the same game played with the synthesis of purely organic receptors). However, metals bring an additional benefit: according to their electronic configuration, they manifest different geometrical preferences, which adds a further element of selectivity. All these advantages (possibility to carry out recognition studies in water, improved selectivity) should stimulate anion chemists to make a more extended use of metals in the design of effective receptors. A final, less ‘scientific’ point could be mentioned. Anions are colourless as well as their organic receptors. Thus, the process of anion recognition in most cases has to be followed only indirectly (e.g. through the shift of an 1H NMR signal). On the other hand, transition metal containing receptors show a variety of colours which typically undergo drastic changes on interaction with anions, generating in the operator surprise, emotion and pleasure. Similar feelings have been probably experienced by Alfred Werner and his students when synthesising and manipulating their metal amine complexes. This may have been a further, perhaps not secondary, stimulus for the development of coordination chemistry at the University of

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Chem Soc Rev Zurich, more than one century ago (with the additional gratification of being able to rationalise this coloured chemistry within the frame of the locally and freshly developed ‘Theory of constitution and configuration of higher order compounds’).72

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