Alkylsulfides of Ag(I) and Au(I) as Metallosurfactants - Langmuir (ACS

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Alkylsulfides of Ag(I) and Au(I) as Metallosurfactants Coenraad R. van den Brom, Manfred Wagner, Volker Enkelmann, Katharina Landfester, and Clemens K. Weiss* Max-Planck-Institut f€ ur Polymerforschung, Ackermannweg 10, 55128 Mainz, Germany Received July 9, 2010. Revised Manuscript Received August 13, 2010 Several representative, interfacially active silver(I) nitrate alkylsulfide complexes were synthesized and characterized in detail. The complexes form extended structures in the solid state and in solution. Interestingly, a two-phase approach, in which aqueous silver nitrate is combined with organically dissolved sulfides, leads to the in situ formation of the complexes at the water-organic interface and a strong reduction of the surface tension. Despite their low solubility in water or organic solvent, these complexes are capable of stabilizing eicosane emulsions and dispersions in water. Thus, these silver nitrate alkyl sulfides represent a new class of metallosurfactants in which the metal ion is crucial for the interfacial activity. Gold(I) chloride alkyl sulfides show the same effect to a lesser extent.

1. Introduction Metallosurfactants are amphiphilic (d- or f-block) metal complexes that show surface or interface activity1 or can form micellar structures.2 Such complexes are not only of fundamental interest, but also give rise to many applications:2,3 The metal introduces functionality and increased reactivity into the surfactant, such that they may be used, e.g., for the formation of catalytic metallomicelles4 or as enzyme mimics.5 Conversely, they provide a means to direct metal ions toward polar/apolar solvent interfaces, enabling, for instance, catalysis at the interface,6 templating the formation of metal-bearing mesostructures,7 or the buildup of optoelectronic thin films.8 Most metallosurfactants reported in the literature apply metal ions that are relatively well shielded by chelating ligands9 or use ligands that are surfactants themselves. Thus, the origin of their interfacial activity is essentially the same as for conventional surfactants. In principle, direct solvent-metal interactions via coordination should also lead to interfacial activity: hydration of a metal complex with one or more hydrophobic ligands can lead to amphiphilicity (especially in combination with electrostatic effects). To our knowledge, this type of interfacial activity and thus this class of metallosurfactants have received only a little attention in the literature. Rather, it seems that this possibility is often overlooked when metal ion extraction is investigated: the *E-mail: [email protected]. (1) Fallis, I. A.; Griffiths, P. C.; Griffiths, P. M.; Hibbs, D. E.; Hursthouse, M. B.; Winnington, A. L. Chem. Commun. 1998, 665. (2) Griffiths, P. C.; Fallis, I. A.; Tatchell, T.; Blishby, L.; Beeby, A. Adv. Colloid Interface Sci. 2008, 144, 13. (3) Griffiths, P. C.; Fallis, I. A.; Chuenpratoom, T.; Watanesk, R. Adv. Colloid Interface Sci. 2006, 122, 107. (4) Zhang, Y. X.; Zeng, H. C. Langmuir 2008, 24, 3740. (5) Kriste, A. G.; Vizitiu, D.; Thatcher, G. R. J. Chem. Commun. 1996, 913. (6) (a) Fallis, I. A.; Griffiths, P. C.; Cosgrove, T.; Dreiss, C. A.; Govan, N.; Heenan, R. K.; Holden, I.; Jenkins, R. L.; Mitchell, S. J.; Notman, S.; Platts, J. A.; Riches, J.; Tatchell, T. J. Am. Chem. Soc. 2009, 131, 9746. (b) Hay, R. W.; Govan, N.; Parchment, K. E. Inorg. Chem. Commun. 1998, 1, 228. (c) Polyzos, A.; Hughes, A. B.; Christie, J. R. Langmuir 2007, 23, 1872. (7) (a) Pileni, M. P. Langmuir 2001, 17, 7476. (b) Amos, K. E.; Brooks, N. J.; King, N. C.; Xie, S. H.; Canales-Vazquez, J.; Danks, M. J.; Jervis, H. B.; Zhou, W. Z.; Seddon, J. M.; Bruce, D. W. J. Mater. Chem. 2008, 18, 5282. (8) Chu, B. W. K.; Yam, V. W. W. Inorg. Chem. 2001, 40, 3324. (9) (a) Bowers, J.; Danks, M. J.; Bruce, D. W.; Webster, J. R. P. Langmuir 2003, 19, 299. (b) Domínguez-Gutierrez, D.; Surtchev, M.; Eiser, E.; Elsevier, C. J. Nano Lett. 2006, 6, 145.

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tendency to accumulate at the interface is overcome by a large excess of extractant and by the addition of soft counterions (e.g., ClO4-)10 or adduct-forming adjuvants.11 On the other hand, supramolecular structure-forming capabilities of such complexes have been reported, notably for metallomesogens. For example, the complexation of imidazoles to AgNO3 led to the formation of smectic liquid crystals,12 and dendritic isocyanides even formed micelles upon complexation of coinage metals.13 Dynamic, in situ formation of amphiphiles from two nonamphiphilic components at the aqueous-organic interface provides an interesting means to control the locality of the surface-active species. Only a few methods to realize such a system have been reported. The best-known example is the deprotonation of organically dissolved (fatty) acids by base, which has been applied in systems ranging from oil-recovery14 to droplet stabilization in miniemulsions15 or microfluidics.16 Although interfacial complexation reactions are known,17 to our knowledge, no targeted attempts to prepare metallosurfactants from two-component systems have been reported in the literature. Here, we present silver(I) and gold(I) sulfide complexes that present a class of metallosurfactants with metal ion head groups. Interestingly, these complexes can both be synthesized beforehand and formed in situ at a water-organic interface and thus provide the first identified example of a two-component metallosurfactant. Silver and gold are interesting candidates for the preparation of metallosurfactants for a number of reasons. Both Ag(I) and Au(I) have a tendency toward linear twofold coordination. When two different ligands are present, this leads to the formation of linear asymmetric compounds, which provides a strong geometric advantage for interfacial activity. Such complexes are comparable in shape to conventional low molecular (10) Ohki, A.; Takagi, M.; Ueno, K. Anal. Chim. Acta 1984, 159, 245. (11) Imura, H.; Namai, T.; Ishimori, K. I.; Hayashi, S.; Ohashi, A.; Ohashi, K. Bull. Chem. Soc. Jpn. 2005, 78, 2146. (12) Lee, C. K.; Hsu, K. M.; Tsai, C. H.; Lai, C. K.; Lin, I. J. B. Dalton Trans. 2004, 1120. (13) Coco, S.; Cordovilla, C.; Donnio, B.; Espinet, P.; Garcı´ a-Casas, M. J.; Guillon, D. Chem.;Eur. J. 2008, 14, 3544. (14) Amaya, J.; Rana, D.; Hornof, V. J. Solution Chem. 2002, 31, 139. (15) Aschenbrenner, E. M.; Weiss, C. K.; Landfester, K. Chem.;Eur. J. 2009, 15, 2434. (16) Ward, T.; Faivre, M.; Stone, H. A. Langmuir 2010. (17) Michel, T.; Nitsch, W. Chem. Ing. Technol. 1990, 62, 738.

Published on Web 09/17/2010

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Article Chart 1

Figure 1. Detail of the crystal structure of 2 showing its tetrameric cyclic structure. Colors: Ag = purple, S = yellow, O = red, C = gray. Hydrogens are omitted for clarity.

weight surfactants. Ag(I) and Au(I) both form reasonably stable complexes with strong soft S- and P-based Lewis bases.18 Hydration of Ag(I) and Au(I) is weaker than that of, e.g., 3d metal cations;19 therefore, sulfide ligands are not displaced by water. Silver- and gold-based surfactants may be interesting for catalytic applications or as precursors for metallic nanostructures. First, we describe the characterization of the prepared complexes with some emphasis on their crystal structure. Then, we assess their behavior in solution by means of diffusion-ordered NMR (DOSY) experiments. The interfacial activity is shown by tensiometry, for both a priori and in situ prepared complexes. Finally, we show how these complexes were used to stabilize aqueous dispersions of eicosane.

2. Results and Discussion 2.1. Complex Preparation and Characterization. The isolation of silver nitrate sulfides was first reported by R^ay in 1931.20 The same convenient concept of combining AgNO3 and a sulfide SR2 in a suitable solvent followed by precipitation and recrystallization of the complex was employed here to prepare the novel compounds silver nitrate didodecylsulfide 1, silver nitrate dodecylmethylsulfide 2, and silver nitrate dihexylsulfide 3. Remarkably, by following the original procedure by R^ay with dibenzyl sulfide 4a,20 the reported complex [AgNO3(SBz2)] was not obtained, but the bis-coordinated complex [AgNO3(SBz2)2] 4 was isolated instead (see Chart 1). All compounds were obtained in good yields and are slightly light-sensitive. In particular, 2 (with only one alkyl chain) bears a structural resemblance to conventional surfactants like, for example, sodium dodecyl sulfate (SDS) and N-hexadecyl-N,N,N-trimethylammonium bromide (CTAB). Thermal analysis by differential scanning calorimetry (DSC) showed that the complexes 2-4 have well-defined melting points (at 77.7, 63.5, and 114.5
C, respectively). Interestingly, 1 features two endothermal peaks on heating, at 37.7 and 64.6
C, which implies the formation of a different nonliquid phase between these temperatures. Detailed structural investigation of this phase is outside the scope of this paper. Sulfides of gold(I) chloride are known even longer21 than those of silver nitrate. Following R^ay’s method,21c we prepared gold chloride didodecyl sulfide 5, gold chloride dodecylmethylsulfide 6, and gold chloride dibenzylsulfide 8 as crystalline compounds. These reactions require reduction of Au(III) by the sulfide. For 5 (18) (a) Lancashire, R. J. In Comprehensive Coordination Chemistry; Wilkinson, G., Ed.; Pergamon: Oxford, 1987; Vol. 5. (b) Puddephatt, R. J. In Comprehensive Coordination Chemistry; Wilkinson, G., Ed.; Pergamon: Oxford, 1987; Vol. 5. (19) (a) Holland, P. M.; Castleman, A. W. J. Chem. Phys. 1982, 76, 4195. (b) Ni, B.; Kramer, J. R.; Werstiuk, N. H. J. Phys. Chem. A 2005, 109, 1548. (c) Asthagiri, D.; Pratt, L. R.; Paulaitis, M. E.; Rempe, S. B. J. Am. Chem. Soc. 2004, 126, 1285. (20) R^ay, P. C.; Adhikari, N.; R^ay, H. J. Indian Chem. Soc. 1931, 8, 689. (21) (a) Phillips, F. C. J. Am. Chem. Soc. 1901, 23, 250. (b) Herrmann, F. Ber. Dtsch. Chem. Ges. 1905, 38, 2813. (c) R^ay, P. C.; Sen, D. C. J. Indian Chem. Soc. 1930, 7, 67.

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and 6, the resulting sulfoxide byproducts were removed by column chromatography. The hexylsulfide complex 7 was a labile and corrosive oil at temperatures above -20
C, impeding its full characterization. Both 7 and the sulfides of gold(I) nitrate (analogous to 5-8) were too light- and temperature-sensitive to be of use for further testing as metallosurfactants. 2.2. Crystal Structures. The solid-state structure as revealed by single crystal X-ray diffraction of complexes 1-4 is fundamentally interesting, since Ag(I) is known for its coordinative flexibility: the exact number of ligands and binding angles is to a great extent determined by the ligands.22 Moreover, extended structure formation similar to that in the solid might play a role in the solution behavior of these compounds. The structures of 1, 2, and 3 are based on bridging Ag-S bonds. The packing of 2 is completely different from that of 1 and 3, due to the asymmetric and symmetric shape of 2a and 1a/3a, respectively. Crystals of [AgNO3(SMeDodec)] 2 are monoclinic, with space group P21/c (Supporting Information Table S1). A lamellar structure is observed, consisting of “inorganic” layers (AgNO3S) interspersed by organic layers consisting of interdigitated alkyl chains in a herringbone arrangement. Figure 1 shows that four asymmetric units form eight-membered Ag-S-Ag-S 3 3 3 rings, in which all Ag-S distances are about 2.54-2.59 A˚. The two inequivalent Ag centers both have S-Ag-S angles of around 120
. Every Ag connects to two nitrates via one oxygen and via two oxygen atoms, respectively. Thus, the nitrates connect the Ag-S rings. Crystals of [AgNO3(SHex2)] 3 are monoclinic with space group C2/c. This crystal structure is also lamellar, but the hexyl chains of adjacent lamellae do not interdigitate. Instead, the hexyl chains sticking out of one inorganic layer form densely packed slabs on either side. The formation of an extended structure based on Ag-S bridging is observed, but unlike 2, this leads to the formation of a helical zigzag chain structure, as can be seen in Figure 2a and b. Crystals of 1 have a similar structure, due to the correspondence of 1a and 3a (Figure 2c and d), though the plane-plane distance is of course larger for 1. The differences between 1 and 3 illustrate how the crystal structure is determined by a balance between two types of interactions driving the induction of crystalline order: (a) Ag-S and Ag-NO3 coordination in the inorganic layer and (b) van der Waals interactions between the alkyl chains. The structures favored on the basis of these two factors are not necessarily commensurate. In the case of 3, the strength of the coordinative interactions dominates, leading to a (suboptimal) nonparallel arrangement for the alkyl chains. The longer chains in 1 lead to stronger interchain interactions and hence to an all-parallel packing for the alkyl chains (Figure S1b, Supporting Information). Here, the incommensurability between the coordinative ordering and alkyl stacking is overcome by the torqued bonds between the first few carbon atoms of every chain and tilting of the entire alkyl slab (Figure 2c,d). (22) Fielden, J.; Long, D. L.; Slawin, A. M. Z.; Kogerler, P.; Cronin, L. Inorg. Chem. 2007, 46, 9090.

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Figure 2. Details of the crystal structure of 3, (a) projected along the a-axis and (b) projected along the b-axis, revealing a helical chain structure based on Ag-S bridging; and details of the crystal structure of 1 projected (c) along the a-axis and (d) along the b-axis, illustrating the similarity to 3. (Colors as in Figure 1).

Figure 3. Details of the crystals structure of 4 showing how the benzyl groups of two coordinating sulfides 4a prevent the formation of extended structures. (Colors as in Figure 1.)

The crystal structure of 4 is remarkably different. In these orthorhombic crystals with space group P21nb, there are two sulfide ligands per Ag(I) cation, and thus, there are no extended Ag-S structures (Figure 3). The Ag(I) cation still binds to two sulfur atoms, but these are not bridging, nor are the nitrates. Thus, the Ag(I) cations are shielded from their surroundings due to the steric requirements of the benzyl groups. (23) (a) Johnson, C. S. Prog. Nucl. Magn. Reson. Spectrosc. 1999, 34, 203. (b) Cohen, Y.; Avram, L.; Frish, L. Angew. Chem., Int. Ed. 2005, 44, 520.

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2.3. Diffusion-Ordered NMR. Diffusion-ordered NMR spectroscopy (DOSY-NMR)23 can reveal whether extended structures as observed in the solid state of 2 and 3 are also present in solution, since complexes of different size or aggregates give rise to different diffusion constants. Here, we show a detailed investigation of the solution behavior of 2 in THF by combining 1H NMR and DOSY-NMR. The exact dimensions of a solvated molecule are hard to determine by DOSY-NMR, since for such small entities, the solvent cannot be considered a continuum and their shape usually deviates substantially from a sphere.24 Comparing sizes of similarly shaped molecules based on the same approximations is justified, though. Indeed, DOSY-NMR has recently been used to investigate the coordination of Ag(I)22,25 and the formation of aggregates of Au(I) complexes by aurophilic interactions,26 so the method seems suitable for the current investigation. Straightforward 1H NMR can also be used to obtain information about chemical or complexation equilibria, provided one or more nuclei have different chemical shifts in different complexation states:27 If the dynamics of an equilibrium are fast compared to the NMR time scale, a weighted average of the chemical shifts of the individual species is observed.27 2.3.1. Effect of Concentration. Variation of the concentration of 2 has a clear effect on both the 1H NMR spectrum and on the diffusion constants observed by DOSY. In the 1H NMR spectrum (see Supporting Information, Figure S5), the chemical shifts of the protons of both the sulfur-bound methyl (SCH3) and R-methylene (SCH2CH2 3 3 3 ) groups show a concentrationdependent shift when the concentration is varied between 2.2 μM and saturation at about 13 mM as can be seen in Figure 4a. The concentration dependence clearly shows that there is a dynamic equilibrium. At low concentrations, the shifts of 2 converge toward the corresponding shifts for free 2a, which are represented by the dotted lines in Figure 4a (recorded separately). This means that the complex is mostly dissociated at low concentrations. By fitting to a simple 1:1 complexation model (see Supporting Information), log(K) for this equilibrium was calculated as 3.5 ( 0.2 at 298 K, which agrees well with the values of 3.4-3.8 reported in the literature for similar compounds with thiodiglycol28 or tetrahydrothiofurane.29 A justification for this approach (namely, that the exchange is fast on NMR time scales) is provided by the addition of 0.5 equiv of free 2a to a solution of 2 (at ca 11 mM), since this leads to only one coalesced R-methylene triplet and, similarly, only one methyl singlet. Figure S6 (Supporting Information) shows that in a DOSYNMR spectrum of 2 at ca 13 mM in THF-d8 recorded at 298.4 K, the complex is well-separated from THF and water. The approximate complex size of 2 in THF-d8 at varying concentrations (0.71 mM until saturation at ca 13 mM) was calculated using the Stokes-Einstein relation (see Supporting Information) and the diffusion constants obtained from DOSY-NMR spectra (Figure 4b). Figure 4b and Supporting Information Table S2 clearly show that the average size increases as a function of concentration. At all accessible concentrations, it is substantially higher than that of the free ligand, as indicated by the dotted lines in Figure 4b. (24) Macchioni, A.; Ciancaleoni, G.; Zuccaccia, C.; Zuccaccia, D. Chem. Soc. Rev. 2008, 37, 479. (25) Scheele, U. J.; Georgiou, M.; John, M.; Dechert, S.; Meyer, F. Organometallics 2008, 27, 5146. (26) Balzano, F.; Cuzzola, A.; Diversi, P.; Ghiotto, F.; Uccello-Barretta, G. Eur. J. Inorg. Chem. 2007, 5556. (27) Connors, K. A. Binding Constants; John Wiley & Sons: New York, 1987. (28) Sigel, H.; Rheinberger, V. M.; Fischer, B. E. Inorg. Chem. 1979, 18, 3334. (29) Tilley, R. I. Aust. J. Chem. 1990, 43, 1573.

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Figure 4. (a) Proton chemical shifts of the sulfur-bound methyl and R-methylene groups in 2 as a function of concentration at 298.3 K and fits of these shifts to a 1:1 binding isotherm (solid lines). The dotted lines marked * and † represent the corresponding shifts for the R-methylene and methyl groups of the free sulfide 2a, respectively. (b) Diffusion coefficients D (circles) and corresponding approximate hydrodynamic radii Rh (squares) for 2 at various concentrations in THF-d8, determined from DOSY. The dashed lines connecting the data points are guides for the eyes only. The dotted lines marked † and § denote D and Rh, for the free sulfide 2a, respectively.

Lower concentrations could not be assessed, due to the signal-tonoise ratio, but it is likely that, at low concentrations, D would converge to that of the free sulfide, analogous to the concentration effect on the chemical shifts of the methyl and R-methylene protons. At the other end of the concentration range, it is very difficult, though, to estimate what the calculated Rh of the “pure” 1:1 complex would be.24 Therefore, the formation of larger aggregates (dimers, trimers, etc.) is not immediately obvious from Figure 4b. However, a comparison of the 1H NMR and DOSY results above shows that at concentrations of 10-∼13 mM there is little shift in δ, whereas Rh increases from 4.0 to 4.5 A˚. The former implies that the fraction of free ligands does not change, and the latter that an increase in average size still occurs. This is an indication that there indeed is formation of higher-order aggregates [Agx(NO3)y(SR2)z] (with at least z>1) not only in solid state, but also in solution, in agreement with conclusions drawn from solubility data.30 (30) Mikhailova, M. P.; Mikhailov, V. A. Zhur. Neorg. Khim. (Russ. J. Inorg. Chem.) 1983, 28, 499.

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Figure 5. (a) Chemical shifts of selected proton resonances of 2 in THF-d8, upon addition of H2O at 298.3 K. The dashed lines connecting the data points are guides to the eye. The dotted lines marked * and † represent the corresponding shifts for the R-methylene and methyl groups of the free sulfide 2a, respectively. (b) Diffusion coefficients D (circles) and corresponding approximate hydrodynamic radii Rh (squares) for the same samples determined from DOSY. The dashed lines connecting the data points are guides for the eyes only. The dotted lines marked † and § denote D and Rh, for the free sulfide 2a, respectively.

2.3.2. Effect of Adding Water. Water might affect the equilibrium or equilibria that are observed, since it can solvate the complexes 2 (either in aggregated form or the individual species) and the small fraction of free AgNO3 (AgNO3 has a very low solubility in pure THF). Remarkably, the addition of water (up to a molar fraction xwater = 0.10) to a 10 mM solution of 2 does not lead to a change in the observed complex hydrodynamic radius; it stays around 4.0 A˚ (Figure 5b). At the same time, the 1H-signals of both methyl and R-methylene protons do shift as a function of water content (ca 0.058 and 0.068 ppm upfield, respectively), whereas the other peaks (e.g., the ω-CH3) do not shift significantly (Figure 5a): The minute upfield shift of the ω-CH3 resonance is attributed to a change of the general solvent environment. This shift may arise since the resonances of residual protons of the THF (that were used as a reference) slightly shift in the presence of water. In a control experiment, the addition of water to free 2a in THF-d8 (up to xwater = 0.10) only led to minute shifts ( ∨ 1a 2a 3a 4a

AgNO3 59

50 36.4 Dodec2S DodecMeS 50 28.4 50 32.0 Hex2S 50 29.0 Bz2S dodecane 47.9 a The error in γ is about 0.5 mN 3 m-1.

NaNO3 59

water -

47.0 46.5 45.7 44.5 48.4

49.1 47.6 46.7 47.7 49.4

The combined result of changing chemical shifts and constant complex size indicates that the observed shift in δ is not caused by a shift of the complexation equilibria, but by a local change in the surrounding of the complex which is strongest near the sulfide’s headgroup. This is consistent with the Ag(I) ions showing a preference for solvation by water over solvation by THF. Moreover, it implies that the interfacial activity (see below) is indeed related to hydration of Ag(I) at the organic/aqueous interface. 2.4. Interfacial Activity of [Ag(I)NO3(SR2)]. Spinning drop tensiometry (SDT) is a versatile technique to measure interfacial tensions,31 in which the interfacial tension γ is determined from the shape of a droplet of one phase (e.g., decane) in a capillary filled with another immiscible phase (e.g., water).32 When the capillary is spun, the shape of the droplet is governed by the balance between centrifugal and interfacial forces. The use of a closed system and the possibility to vary the centrifugal forces by changing the spinning rate make SDT a more reliable method to determine interfacial tensions than, for example, the use of the Du No€uy ring method with a two-layer system. Here, we present an investigation of the interfacial activity of the complexes 1-4, both prepared a priori and in situ, which revealed that the in situ prepared complexes dramatically reduce the interfacial tension. We first determined the interfacial tensions of saturated solutions of 1-3 in dodecane, cyclohexane, and toluene against water. In all these cases, no significant depression of the interfacial tension was observed in comparison to the pure solvents (see Supporting Information Table S3). This is primarily due to the low solubility of these complexes in either water or apolar solvents at ambient temperature, which is estimated to be considerably lower than 1 wt % for all these complexes. In contrast, AgNO3 is highly water-soluble and all sulfides used as ligands are well-soluble in apolar solvents. Therefore, the interfacial tension γ was determined for two-component systems of sulfide-containing n-dodecane droplets versus aqueous AgNO3. Table 1 shows that γ lies in the range 28-36 mN 3 m-1 for all sulfides 1a-4a (50 mM in n-dodecane) versus AgNO3 (59 mM in water). This is substantially lower than γ = 49.4 mN 3 m-1 obtained for neat dodecane (of the purity used in our experiments) versus water. Reference experiments with either only AgNO3 or only one of the sulfides yielded reductions of γ of no more than ca 3 mN 3 m-1 at most. In order to exclude any effects of the ionic strength, another series of reference measurements was performed using the noncoordinating NaNO3 at the same concentration (59 mM in water), which led to depressions of γ of less than ca 5 mN 3 m-1. These results strongly imply that the Ag(I) sulfide complexation at the interface is responsible for the observed interfacial activity. The trend of increasing γ for the series 2a-1a-3a can probably be explained by the increasing bulkiness (31) Seifert, A. M. In Drops and Bubbles in Interfacial Research; M€obius, D., Miller, R., Eds.; Elsevier Science: Amsterdam, 1998. (32) Viades-Trejo, J.; Gracia-Fadrique, J. Colloids Surf., A 2007, 302, 549.

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Figure 6. The interfacial tensions γ between the sulfide 2a in dodecane and aqueous solutions of AgNO3, at 297 K and at varying concentrations of both components. The same SDT data (points) are represented twice in order to show fits to the Szyszkowski equations (the lines) at constant concentrations (a) of AgNO3 and (b) of 2a, respectively.

of the alkyl chains and the corresponding decreasing hydrophile to lipophile balance (HLB). The interfacial activity of these two-component systems was investigated in more detail for 2a. The graphs in Figure 6 clearly show that increasing the concentration of either sulfide 2a or AgNO3, or of both, led to a reduction of γ, down to 15 mN 3 m-1 at concentrations of 100 mM for 2a and 118 mM for AgNO3, respectively. Using higher concentrations of 2a or AgNO3 (0.37 M 2a vs 118 mM AgNO3 or 97 mM 2a vs 0.47 M AgNO3) led to the spontaneous formation of emulsions upon attempting to insert a droplet of the dodecane solution of 2a into the aqueous AgNO3 solution. The emulsification was immediately followed by the formation of a white precipitate which was identified as 2 by IR and 1H NMR. Therefore, the in situ formation of 2 at the interface is responsible for the observed reduction in interfacial tension. The observed emulsification and precipitation indicate that γ was reduced even further than 15 mN 3 m-1, but they hinder further quantitative measurements. In comparison with the well-known conventional cationic surfactant hexadecyltrimethylammonium bromide (CTAB), which has a critical micelle concentration (cmc) of about 0.92 mM,33 the two-component approach allows much higher concentrations (over 100 mM) without leveling out of the interfacial tension, presumably because the low solubility of the complex in either phase prevents the formation of micelles. Instead, crystallization of 2 occurs upon saturation of the interface. CTAB yields an interfacial tension between dodecane and water of ca 6 mN 3 m-1 at (33) (a) Kosaka, O.; Sehgal, P.; Doe, H. J. Surfactants Deterg. 2005, 8, 347. (b) Medrzycka, K.; Zwierzykowski, W. J. Colloid Interface Sci. 2000, 230, 67.

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298 K at its cmc.33b Though complex 2 is much less active at the same concentration, the much higher concentrations of the components in both phases enable a similar surfactant density at the interface as illustrated below. Generally, assessment of γ as a function of surfactant concentration may yield information about the interface excess Γ and the interfacial adsorption constant Kads.34 Development of an elaborate model for our particular case, in which the surfactant forms in situ from two components in two different bulk phases, in which the AgNO3 is dissociated, and either component is only drawn to the interface in the presence of the other, is outside the scope of this paper. Therefore, we took a simplified approach using the empirical Szyszkowski equation:35 γ ¼ γ0 - A lnð1 þ BcÞ where γ0 is the interfacial tension in the absence of surfactant, c is the surfactant concentration, and A and B are fit parameters. In Figure 6, least-squares fits are displayed using this equation for either the variation of γ as a function of AgNO3 at constant sulfide concentration (Figure 6a) or vice versa (Figure 6b). The χ2 and R2 values of these fits (Table S4, Supporting Information) show that these are good fits. Mathematically, the Szyszkowski relationship may be derived from a combination of the Gibbs and Langmuir isotherms for surfactant adsorption.35 Thus, the saturation surface excess Γ¥ (or the effective surface area at saturation ω = 1/Γ¥) and Kads and ΔGads may be determined from A and B, since A = RTΓ¥ and B = Kads, ΔGads =-RT ln(Kads). These values for the fits in Figure 6 are tabulated in Supporting Information Table S4 but need to be interpreted with great caution, since (1) the model does not take the two-component system into account, (2) the Langmuir isotherm ignores lateral motional freedom at the interface, leading to an underestimation of Kads,35a and (3) the model ignores interactions between surfactants at the interface. Nevertheless, a few remarkable observations may be made: (1) The values for ω (0.35-0.56 nm2 per molecule) and of ΔGads are roughly the same for both the fits with constant c(AgNO3) and those with constant c(2a). (2) The values of ω are of a physically reasonable size both in comparison with conventional surfactants (cf. ω = 0.28 nm2 for CTAB)33b and with the minimum possible area estimated from dense packing in the crystal structure of 2 (0.15 nm2). The value of ω decreases at increasing concentrations of the component kept at constant concentration. This decrease indicates a denser packing at saturation of the interface (i.e., a larger maximum surface excess Γ¥), in agreement with a naive picture in which more stabilizing Agþ induces a higher sulfide surface excess and vice versa. (3) The value of ΔGads is roughly the same in all cases, implying that the same distribution/binding process determines the interfacial tension in all cases. The latter implication makes it tempting to compare this ΔGads derived from the apparent adsorption constant Kads (ΔGads ≈ -11.5 kJ 3 mol-1) with ΔGassoc corresponding to the bulk association constant Kassoc for the silver-sulfide complexation at 298 K. As mentioned above, log(Kassoc) is about 3.4-3.8, corresponding to ΔGassoc ≈ -20 kJ 3 mol-1 (at 298 K). Therefore, the concentrations of both Agþ and the sulfide at the interface are probably (34) (a) Butt, H.-J.; Graf, K.; Kappl, M. Physics and Chemistry of Interfaces; 2nd ed.; Wiley-VCH: Weinheim, 2006. (b) Everett, D. H. Basic principles of colloid science; RSC: London, 1988. (35) (a) Lyklema, J. Fundamentals of Interface and Colloid Science; Academic Press: San Diego, 2000; Vol. III. (b) Prochaska, K. Adv. Colloid Interface Sci. 2002, 95, 51.

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orders of magnitude higher than in the respective bulk phases. In conclusion, the two-component SDT measurements provide strong evidence that the [AgNO3(SR2)] complexes 1-4 are indeed interfacially active, provided a sufficient concentration can be reached at the interface. Such a concentration cannot be reached by dissolving the complex in either water or the organic phase, but this problem is effectively overcome by using a two-component approach. Indeed, the strong Ag-S affinity combined with the low solubility of the complexes in either bulk phase essentially traps them at the interface. The gold(I) complexes 5 and 6 at 1 or 5 wt % in dodecane lead to a modest reduction of the interfacial tension against water from 50 to ca 43 mN 3 m-1. The softer character of the chloride ions compared to nitrates, on one hand, enhances the solubility of these compounds in the organic phase, but this clearly is at the expense of their interfacial activity. In principle, in situ preparation of these or similar complexes should be possible by using a water-stable and -soluble source of aqueous Au(I) cations. 2.5. Eicosane Emulsification. A simple emulsification procedure based on sonication was used to demonstrate that the Ag(I) and Au(I)-based metallosurfactants 1-7 are not only interfacially active, but can also be used to prepare and stabilize n-eicosane dispersions in water. Such dispersions were made using both a priori and in situ prepared complexes, the latter via the same two-component approach as in the previous section. n-Eicosane was used as the dispersed phase, since its melting point (37
C) allowed emulsification at moderate temperatures and subsequent cooling instantaneously yielded a dispersion of stable solid particles, that eases characterization by light-scattering, solids content determination, and electron microscopy. Before drawing a comparison between the different complexes, a detailed investigation of one model system, customarily based on 2, is presented. 2.5.1. Particle Size and Solids Content for a Model System. In Table 2 and Figure 7, the solids contents and average particle sizes are compared for eicosane dispersions stabilized with 2, AgNO3/2a (the two-component approach), or CTAB as surfactants and with only AgNO3, only 2a, or no surfactant as control experiments. In all cases, the same amount of eicosane (3 wt % relative to water) and the same molar amounts of sulfide or surfactant were used. The solids contents show that very good stabilizing efficiencies were obtained for all three surfactant systems, whereas only little eicosane is dispersed in the control experiments. The aliquots used for solids content determination were taken within an hour after emulsification. The control dispersions were so unstable (due to creaming) that their solids contents would have been even lower after a longer time of standing. Since formation of the complex is an equilibrium process, the effect of the silver-to-sulfur ratio was investigated too (samples B, C, D). Varying the ratio form 1:1 to 2:1 did not lead to an appreciable difference in the formed dispersions (vide infra). At 10:1, no stable dispersions could be formed, however, presumably because the resulting high ionic strength of the aqueous phase increases the interfacial tension and destabilizes the dispersion. The above observations were quantified by the average particle sizes as determined by photon cross-correlation spectroscopy (PCCS). PCCS is a dynamic scattering technique that is particularly suitable for quickly analyzing relatively dense dispersions, since the cross-correlation configuration allows correction for multiple-scattering phenomena. Moreover, the measurements presented below were corroborated by more detailed angledependent dynamic light scattering (DLS) analyses for DOI: 10.1021/la1027543

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Table 2. Preparative Details and Analysis of Aqueous Eicosane Dispersions Stabilized by 2, by the Combination of 2a and AgNO3, and Several Controls

surfactant system

2, 2a, CTAB

AgNO3

conc. w. r.t. eicosane

in water

-1

(wt %)

(mmol 3 kg )

14.2

368.8

-1

(mmol 3 kg )

ratio Ag:S ICP analysis

abs.

eff.

(mol:mol)

(mol:mol)

wt %

(%)

1:1

before dialysis: 0.91 ( 0.12 after dialysis: 0.05 ( 0.01

3.28

96

3.26 2.96

95 82

n.a. 0.21 0.39 3.40 0.41

n.a. 7 12 101 14

A

2

B C

2a þ AgNO3 2a þ AgNO3

8.0 8.0

371.9 371.9

11.5 23.0

1:1 2:1

D E F G H

2a þ AgNO3 2a only AgNO3 only CTAB Blank

8.0 8.0 ; 13.5 ;

371.9 371.9 ; 369.2 ;

114.6 ; 23.0 ; ;

10:1 ; ; ; ;

Figure 7. Comparison of the approximate particle sizes for 2-stabilized eicosane dispersions in water measured by PCCS. The lower, middle, and upper boundaries of the boxes denote the sizes x10, x50, and x90 at which the cumulative size distribution reaches 10%, 50%, and 90%, respectively. The sample labels correspond to those in Table 2.

representative samples as described in the Supporting Information. Figure 7 provides a graphical representation of the particle size distributions obtained by PCCS. In this graph, the x50 value represents the center of the cumulative size distribution that was calculated using an NNLS algorithm. The x10 and x90 values (the sizes at 10% and 90% of the cumulative size distribution, respectively) are also presented to give an impression of the approximate width of the size distributions. Both 2 and AgNO3/2a stabilize particles of about 300 nm in diameter. CTAB-stabilization yields smaller particles (ca 170 nm in diameter), but the sizes for all three surfactant systems are of the same order of magnitude. The control experiments yield dispersions with macroscopic flocs, as well as some particles or aggregates that are much larger and hence are outside the range that is reliably accessible with our PCCS setup (i.e., >1 μm). This observation convincingly shows that only complex 2 (prepared either a priori or in situ) and CTAB are capable of stabilizing eicosane dispersions. 2.5.2. Silver Content and Location. The silver and sulfur contents of several dispersions were determined by inductively coupled plasma optical emission spectroscopy (ICP-OES), both before and after dialysis against water. The measured silver content was expressed as a percentage of the amount of silver originally added during formulation. Both for the 2-stabilized and 15800 DOI: 10.1021/la1027543

solids content

prep.

before dialysis: 1.60 ( 0.25 after dialysis: 0.06 ( 0.01

for the AgNO3/2a-stabilized dispersions, the Ag and S contents and their ratios are consistent with the original formulation, when taking into account the accuracy of the ICP method (see Table 2 and Table S6, Supporting Information).36 Upon dialysis, however, the amount of silver is drastically reduced, which can only happen if the silver is in contact with the continuous water phase. At the same time, no sulfur is lost from the samples, which indicates that the sulfides are indeed tightly adsorbed on (or located inside) the eicosane particles. The loss of Ag only is exactly what one would expect if the complexes are located at the particle interface, since the Ag(I)-sulfide bonding is a dynamic equilibrium, as reported in the literature28,29 and demonstrated by our NMR experiments. The alternative explanation, namely, that the Ag(I) is predominantly present as “free” solvated ions in the continuous phase, could be ruled out by conductometry. The AgNO3/2a-stabilized eicosane dispersion (B) had a conductivity of 1.027 mS 3 cm-1. This is only 2/3 of the conductivity obtained for the applied AgNO3 stock solution (1.498 mS 3 cm-1) at the same temperature (4
C). This lowering of the conductivity indicates that one of the conducting species (Agþ and NO3-) has a limited mobility, consistent with the silver cations being located predominantly on the surface of the eicosane particles. Finally, it should be noted that the dialyzed samples (i.e., where most of the silver cations were removed) became unstable, leading to creaming of the eicosane and thus providing evidence that silver(I) plays an essential role in stabilizing the particles. This observation is consistent with our hypothesis that the [Ag(I)SR2]þ complexes stabilize the interface and protect against particle coalescence by a combination of hydration and electrostatic interaction. 2.5.3. Solidity, Stability, and Shelf Life. The eicosane particles in dispersions stabilized by AgNO3/2a or CTAB were solid at room temperature as determined by differential scanning calorimetry (see Supporting Information Figure S10). This solidity obviously enhances the stability of these dispersions. Indeed, the dispersions stabilized by 2 or AgNO3/2a are similar in stability to those stabilized by CTAB: they could be stored (at 5
C in darkened containers) for at least two weeks without a significant change in particle size as determined by PCCS. After one month, only a moderate increase in average particle size was observed (see Supporting Information Table S8). The stability of the complexes trapped at the interfaces is indicated by the fact that all dispersions stay white (i.e., do not show blackening indicative of decomposition) and emphasized by the fact that the silver can be washed out by dialysis, which would not be possible for surface-attached or precipitated silver or silver-oxide decomposition products. (36) Vogel, N.; Hauser, C. P.; Schuller, K.; Landfester, K.; Weiss, C. K. Macromol. Chem. Phys. 2010, 211, 1355.

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2.5.4. Surface Coverage. For AgNO3/2a stabilized eicosane emulsions, it was found that doubling the AgNO3 concentration while keeping the 2a content of the eicosane phase constant (i.e., Ag:S = 2:1) led to particles of the same size (vide supra). This observation suggests that the sulfide content (and hence the total amount of complex 2 formed at the interface) determines the particle size. The surface area of a single eicosane particle (of ca. 300 nm diameter) can be compared with the estimated area that would be occupied by a single complex 2. Here, one might take the estimated area from either the crystal structure (0.15 nm2) or ω obtained above by tensiometry (0.35-0.56 nm2). The former is favored since it reflects the densest possible packing and the sulfide concentration in the eicosane is much higher than in the dodecane-based tensiometric analyses. This comparison reveals that full coverage of the particle surface requires 1.88  106 sulfides per eicosane particle (see Supporting Information for details). On the basis of the eicosane/2a ratio used, an estimated number of 2.5  106 sulfides were present per eicosane particle. The good agreement between these two numbers is remarkable and is consistent with the hypothesis that most of the sulfides are located at the surface of the eicosane particles, in the form of the complex 2. In summary, all the analyses presented in this section provide a very consistent description of 2-stabilized eicosane dispersions. The dispersions are essentially the same, regardless of whether they are made using the a priori prepared complex or they rely on in situ formation, sporting a similar size, solids content, and stability. Moreover, compelling evidence is presented that the formation and stabilization of the particles is indeed due to the presence of hydrated (and charged) Agþ-alkylsulfide species at the eicosane/water interface. 2.5.5. Eicosane Emulsions Stabilized by Further [Ag(I)SR2] Complexes. Table S7 and Figure S9 (Supporting Information) summarize the results that were obtained upon preparation of eicosane dispersions stabilized with 1, 3, and 4 using the same molar concentrations of all complexes or sulfides compared to the eicosane. Both a priori and in situ prepared complexes were used. The size distributions as represented in Supporting Information Figure S9 show that the complexes of all sulfides give rise to stable emulsions with sizes of the same order of magnitude. Three important observations can be made: First, the sizes obtained for dispersions made with a priori prepared complexes are similar to those using the two-component approach. This observation supports the idea that in the twocomponent systems using 1a-4a indeed the corresponding complexes 1-4 are formed and, conversely, that the ultrasonication provides sufficient energy and agitation to overcome the low solubility of 1-4. Second, the dispersions prepared with complexes based on the asymmetric ligand 2a are slightly smaller and have a narrower size distribution than those with the bulkier ligands 1a and 3a, which would be consistent with the view that surfactants with a single alkyl chain are more suitable to stabilize direct emulsions for geometric reasons. Interestingly, the aromatic ligand 4a leads to the same particle size as 2a, which correlates with the roughly equal water/dodecane interfacial tension that was reported above for AgNO3 with 2a and 4a, respectively. Third, control experiments with only the sulfides 1a-4a gave rise to flocculent dispersions with large aggregate sizes (>1 μm), low solids contents (7-12 wt %), and a high tendency to fully cream out within the course of one day. In summary, we provided evidence that the class of silver nitrate sulfides with apolar (aliphatic or aromatic) ligands is effective in stabilizing aqueous eicosane emulsions and can hence be classified as metallosurfactants. Langmuir 2010, 26(20), 15794–15801

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2.5.6. Eicosane Emulsions Stabilized by Gold (I) Complexes. The gold (I) complexes 5, 6, and 8 were also tested for their ability to prepare and stabilize aqueous eicosane emulsions, following the same procedure as for the Ag(I) complexes. For 5 and 8, this resulted in eicosane particles in the same size regime as for the corresponding silver complexes (ca. 300-600 nm, Supporting Information Figure S9), albeit with a much lower efficiency in terms of the solids content (Supporting Information Table S7). Complex 6 decomposed during ultrasonication, resulting in a gray/black flocculated sample. Altogether, these observations show that sulfide complexes of Au(I)Cl can in principle be used as metallosurfactants, though they are less efficient than their Ag(I)NO3 counterparts.

3. Conclusion We have shown that alkylsulfide complexes of Ag(I)NO3 and Au(I)Cl with alkyl chains of six or more carbon atoms could successfully be synthesized. Single crystals of the Ag(I) complexes 2 and 3 provided an example of the coordinational flexibility of silver and sulfur, which gave rise to extended cyclic or helical structures in the solid state. In THF solution, on the other hand, NMR investigations revealed that a dynamic equilibrium was established in which the monomeric complex was the dominant species for 2, though both dissociated species and (especially at higher concentrations) extended structures were also present. It was also shown by DOSY-NMR that water preferably hydrated the Ag-S headgroup of these complexes, without inducing dissociation of the Ag-S bond. This observation provided an important indication as to why these complexes are amphiphilic. The complexes 1-4 are not very soluble in either water or apolar solvents. However, they do assemble at the interface between such phases, where they considerably lower the interfacial tension. Such interfacial assembly of the complexes could elegantly and efficiently be achieved by a two-component approach in which AgNO3 is introduced in the aqueous and a sulfide in the organic phase, respectively. Alternatively, the preformed complexes 1-4 could also be driven toward the interface by mechanical agitation (sonication). Thus, stable dispersions of eicosane in water could be formed in the presence of 1-4 by ultrasonication, regardless of whether the preformed complex was present or whether it was allowed to form in situ from AgNO3 and the sulfide. A series of tests provided conclusive evidence that in both cases the same species stabilized the interface (in the sense of lowering the interfacial tension) and the dispersion (in the sense of preventing aggregation or coalescence of the eicosane particles). To a lesser extent, Au(I)Cl sulfides displayed the same behavior as their AgNO3 counterparts. In conclusion, we have clearly shown that alkyl sulfide complexes of Ag(I) and Au(I) are promising amphiphiles that can be used as metallosurfactants. In particular, the two-component approach provides a promising way to direct metal ions precisely to the interface. Further development of this concept (especially for gold) is possible, for instance, by introducing more strongly coordinating ligands and improving the stability of the aqueous species. Acknowledgment. The authors gratefully acknowledge N. Vogel and C. Hauser for performing the ICP-OES measurements and B. M€uller and Dr. A Kr€oger-Brinkmann for the DLS measurements. Supporting Information Available: Experimental Section, additional graphs, tables and explanations. This material is available free of charge via the Internet at http://pubs. acs.org. DOI: 10.1021/la1027543

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