Voltammetry of Copper Sulfide Particles and Nanoparticles

Suspensions of copper sulfide powders held at accumulation potentials of 0 to −0.2 V (vs ... However, after 3 or 4 aliquots, an additional HS- aliqu...
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Environ. Sci. Technol. 2005, 39, 7492-7498

Voltammetry of Copper Sulfide Particles and Nanoparticles: Investigation of the Cluster Hypothesis IRENA CIGLENEC ˇ KI,† D A M I R K R Z N A R I CÄ , † A N D G E O R G E R . H E L Z * ,‡ Center for Marine and Environmental Research, Rudjer Bosˇkovic´ Institute, Bijenicˇka 54, 10000 Zagreb, Croatia, and Department of Chemistry and Biochemistry, University of Maryland, College Park, Maryland, 20742

An association of Cu with sulfide in aerobic natural waters has been attributed to these components’ coexistence in clusters of sizes intermediate between mononuclear complexes and colloidal particles. This hypothesis is investigated here. Copper sulfide solid phases display sizerelated voltammetric behavior at Hg electrodes. Suspensions of copper sulfide powders held at accumulation potentials of 0 to -0.2 V (vs Ag/AgCl) produce voltammetric peaks near -0.15, -0.65, and -0.95 V during subsequent cathodic scans. The first two peaks arise from electrochemically generated Cu-oxyhydroxides and HgS; the -0.95 V peak arises from reduction of sorbed copper sulfide particles. Nanoparticles of radius ∼10-8 m produce the third peak even without stirring or accumulation. Still smaller analytes give only the first two peaks. Published evidence alleging production of subnanometer copper sulfide clusters during titrations of Cu2+ and HS- was not reproduced when sulfide oxidation was avoided. Instead, such titrations apparently generate nanoparticles. The titration stoichiometry is 1/1, consistent with previous descriptions of this process: Cu2+ + HS- f 1/2Cu2S‚S0 (brown sol) f CuS (green sol). Titrating Cu2+ into organic-rich (muscilaginous) Adriatic Sea water, which contains 10-7 M natural thiols and sulfide, produces solid products. In the future, voltammetry might prove useful for studying semiconductor sulfide nanoparticles in nature.

Introduction The surprising discovery of traces of sulfide in oxic marine waters (1-5) created particular interest in the stability of copper sulfide complexes, which originally were thought likely to account for much of this sulfide (6, 7; but see contrasting views 8, 9). Motivated partly by this interest, Luther et al. (10) undertook voltammetric titrations to explore the stability of sulfide complexes of several metals. For Cu at pH 8.3, each aliquot of added HS- initially caused a decrease in free Cu2+ in a molar ratio, ∆Cu/∆HS, of 1/1. However, after 3 or 4 aliquots, an additional HS- aliquot caused abrupt regeneration of Cu2+ in solution; subsequently, ∆Cu/∆HS was 2/3, or * Corresponding author phone: 301-405-1797; fax: 301-314-9121; e-mail: [email protected]. † Rudjer Bos ˇkovic´ Institute. ‡ University of Maryland. 7492

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sometimes 4/5. This odd behavior was first interpreted (10) as due to formation of CuS0 complexes in the initial stage and Cu2S32- complexes in the second stage. Onthe basis of the doubtful assumption of equilibrium control, stability constant values were derived. Later (11), laser ablationFourier transform mass spectra (FTMS), obtained from vacuum-dried, equimolar Cu2+/HS solutions, were used to infer that the titration products were actually clusters such as Cu3S30, Cu3S60, Cu3S90, etc.; EPR and NMR showed that titration products contained predominantly Cu(I), contrary to the authors’ earlier assumption. These observations raise a number of interesting questions. One concerns the actual size of the Cu-S products. Are the products small, reversibly formed complexes (e.g., Cu2S32-), are they oligomers of intermediate size (e.g., Cu3S90), or are they much larger products that might be more appropriately described as nanoparticles or colloids? Several lines of evidence favor nanoparticles (defined here as particles with radius 1-1000 nm), rather than small complexes or subnanometer clusters. First, titrations such as those in Luther et. al (10, 11) produce momentary supersaturations above 1010 (IAP/Ksp) with respect to Cu sulfide phases. Classical nucleation theory would predict no barrier to particle nucleation under such extreme conditions (12, 13). Second, many workers have in fact observed CuS nanoparticles when HS- is mixed with free or complexed Cu2+ (9, 14-16). Except when Cu2+ is bound by strong chelating agents, nanoparticles form quickly (17, 19); reported radii are 3-26 nm, implying particles containing 104-106 CuS units. Such nanoparticles are smaller than visible light wavelengths, and their existence is apparent by eye solely through coloration of the aqueous medium. After a period of hours, flocculation produces aggregates large enough to settle and become visible. In natural waters, Rozan et al. (21, 22) studied O2-resistant copper sulfide species and found that they possessed acidreactivity and FTMS properties like those of the alleged clusters formed in the laboratory. These species accounted for substantial fractions of filterable Cu and reduced sulfur in some natural water samples, implying that they might be important in biogeochemical cycles of both components. In some samples, more than half the sulfide-bound Cu was too large to pass an ultrafiltration membrane designed to retain globular proteins with a molecular mass above 3000 Da. Assuming that globular proteins are spheres with a density of 0.8 g/cm3, a 3000-Da membrane would have a nominal pore radius of 1.2 nm. This evidence supports also describing at least some of these products as nanoparticles. However, Rozan et al. suggest that organic coatings on much smaller Cu-S cores might explain the large sizes. Clarifying the nature of these species will influence profoundly how aquatic chemists and biologists model and interpret their role in natural ecosystems. Here, we explore what electrochemistry might contribute. We believe that ours is the first voltammetric study of Cu sulfide nanoparticles, which are semiconductors (23, 24). Heyrovsky´ and Jirkovsky´ (25) demonstrate that semiconducting oxide nanoparticles are electroactive, even with no stirring or accumulation period, and they discuss the principles involved. These principles also should apply to suspensions of semiconductor sulfide particles. In the interest of trying to clarify the nature of the Cu-S titration products, we first characterize voltammetric signatures of reference samples containing Cu-S analytes of different sizes and then investigate titrations similar to those of previous workers. 10.1021/es050586v CCC: $30.25

 2005 American Chemical Society Published on Web 08/19/2005

Experimental Section Distilled, demineralized water from a carbon adsorption/ ion exchange filtration system was used for reagent preparation and rinsing. Analytical reagent grade salts were used to prepare the electrolyte solutions, which in most cases consisted of 0.5 M NaCl + 0.03 M NaHCO3 adjusted to pH 8.2. Sulfide solutions were prepared from p.a. grade Na2S‚ 9H2O (Kemika, Croatia) or from anhydrous Na2S (Alfa Aesar). Chalcocite (Cu2S) was prepared by high-temperature reaction of Cu powder (Aldrich Gold Label, 99.995% pure) with elemental S (Aesar, 99.995% pure) in evacuated silica tubes. Nanoparticles, which previous workers have found to be semispherical particles consisting of poorly crystalline covellite (CuS, radius 6-12 nm), were synthesized as previously described (17). One vessel was prepared to contain 2 mM CuNO3 and 2.25 mM trans-1,2-diaminocyclohexyl-tetraacetic acid (DCTA, Sigma) plus N2-purged electrolyte (0.5 M NaCl and 0.03 M NaHCO3, pH 8.2). A second vessel was prepared to contain 2 mM NaHS plus the same N2-purged electrolyte. In a glovebox, equal volumes of the two solutions were rapidly mixed and sealed in septum bottles, which were then passed out of the glovebox. Subsequently, samples could be extracted from these bottles by syringe without exposing the samples, themselves, to air. After reaction times from 10 min to several hours, 5-mL aliquots were extracted, filtered with 200-nm Nuclepore filters, and characterized voltammetrically. Voltammetric titrations similar to those in refs 10 and 11 were performed by preparing 10 µM CuSO4 in 0.05 M NaCl+0.03 M NaHCO3. After deaeration with N2, variable amounts of Na2S were added. For each Na2S concentration, a fresh solution was prepared from Na2S‚9H2O to minimize cumulative loss of HS- by oxidation or volatilization during the titration. Measurements were done with negligible voltammetric accumulation (5 s rest time), starting from -0.05 V and scanning at 0.1 V/s to -1.1 V and then back. Voltammetric techniques (linear sweep voltammetry and cyclic voltammetry) using two electrochemical instrumentation systems were performed during the course of this work. One system was a BAS-100 B electrochemical analyzer (Bioanalytical Systems, West Lafayette, IN), connected to an automatic Controlling Growth Mercury Electrode (electrode surface 0.0259 cm2). The second was an Eco Chemie µ-Autolab connected to a 663 VA Metrohm stand (electrode surface 0.0054 cm2). The reference electrode with each system was an Ag/AgCl (3 M KCl) electrode connected to the solution by a supporting electrolyte bridge. Platinum served as the auxiliary electrode.

Results Cu2S Powders. Figure 1 shows a linear sweep voltammogram (LSV) obtained from a hanging mercury drop electrode (HMDE) that had been coated manually by touching it to a mound of finely ground Cu2S powder before it was immersed in deaerated electrolyte. Two prominent peaks occur. Peaks similar to the smaller one, labeled C3, have been observed using solid-state electrodes consisting of natural chalcocite (26) and using Cu2S film electrodes (27-29). In our case, the Hg drop serves primarily as an electrical contact to the Cu2S particles, which undergo reduction at C3.

Cu2S + H+ + 2e- T 2Cu + HS- EC3 ) EC3′ -

0.029 log([HS-]/[H+]) (1)

On the other hand, the very large, complex peak, C2, is not observed when simple Cu2S electrodes are used and occurs

FIGURE 1. Linear sweep voltammogram (LSV) of finely ground Cu2S powder physically applied to the hanging mercury drop electrode (HMDE) before it was immersed in sulfide-free electrolyte (0.5 M NaCl + 0.03 M NaHCO3, pH ) 8.2). Initial potential, Ei ) -0.20 V; scan rate, v )100 mV/s; accumulation period, td ) 0 s; HMDE area, A ) 0.0054 cm2. Potentials measured vs Ag/AgCl (3 M KCl). in the potential region associated with reduction of HgS to Hg metal.

HgS + H+ + 2e- T Hg(l) + HS- EC2 ) EC2′ -

0.029 log ([HS-]/[H+]) (2)

Complex peaks such as this, with maxima more negative than -0.70 V, are observed when multiple layers of HgS are deposited at Hg electrodes (30, 31). In this experiment, the solution contained no electroactive substance other than Cu2S particles, so the HgS must have been produced when Cu2S was oxidized in an anodic reaction prior to the scan (see below). After Figure 1 was recorded, several Hg drops were detached from the capillary. Particles adhering to the first drop became suspended in the stirred electrolyte. Figure 2A shows cyclic voltammograms obtained from this suspension. When a scan begins immediately after creation of a new Hg drop (i.e., no accumulation period), the cyclic scan is largely featureless. In contrast, when a new drop is held at the starting potential of -0.05 V for 120 s before the scan, then several peaks are apparent in the cathodic scan. These peaks can be due only to particles that have adhered to the Hg drop during the accumulation period. Attachment of Cu2S particles to Hg is probably analogous to this mineral’s adhesion to air bubbles in the froth flotation process and arises because chalcocite is hydrophobic at near-neutral pH (32). The voltammogram for the potential scan in the anodic direction is nearly featureless, as is a second cathodic scan immediately after the first. This evidence suggests that near the negative end of the cycle most of the Cu2S particles have been repelled from the electrode. Both the C2 and C3 peaks in Figure 2 are shifted by roughly +0.15 V relative to their positions in Figure 1. This was caused by the greater amount of analyte on the electrode in the case of Figure 1. At a fixed accumulation period (120 s), the relative intensities of C2 and C3 depend on the potential chosen for accumulation (Figure 2B). At the least negative accumulation potentials tested (-0.15 and -0.20 V), C2 is larger than C3. On the other hand, if accumulation occurs at -0.40 V, then no C2 is observed and C3 is correspondingly larger. From this, we infer that during accumulation at the less negative potentials, Cu2S particles that are sorbed to the Hg electrode become oxidized. Vela´squez et al. (26) observed such a VOL. 39, NO. 19, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 3. Scan 1 (dash-dot line): LSV of CuS nanoparticles grown in 1 mM CuDCTA2- and 1 mM HS-. Scans 2 (solid) and 3 (dashed): after progressive purging with N2. td ) 0 s; A ) 0.0259 cm2.

FIGURE 2. Cyclic voltammograms from Cu2S powder suspended in electrolyte. (A) td ) 120 s at Ei ) -0.05 V (heavy line); td ) 0 s (light line), v ) 0.1 V/s. (B) Three voltammograms recorded with td ) 120 s at different Ei (solid line, -0.15 V; dashed line, -0.20 V; and dash-dot line, -0.40 V); v ) 0.1 V/s, A ) 0.0054 cm2. Note absence of C2 in the scan beginning with accumulation at -0.40 V (dotted line). process at a simple Cu2S electrode and attributed it to a sequence of electro-oxidative processes producing various intermediate Cu sulfides, and finally “CuO” (where “CuO” represents a CuII oxyhydroxide of uncertain hydration state). At accumulation potentials less negative than -0.40 V, the net reaction is as follows: 1

/2Cu2S(s) + 1/2Hg(l) + H2O f “CuO”(s) + 1/2HgS(s) +

2H+ + 2 e- (3)

Accordingly, the HgS required by reaction 2 for the C2 peak is produced only if the potential during accumulation is sufficiently positive for reaction 3 to proceed to the right. If this condition is not met (e.g., accumulation at -0.40 V), then no C2 peak is observed and C3, due to reduction of unreacted Cu2S particles, is larger. (In contrast, when dissolved HS- is present in particle-free electrolyte, the C2 peak is observed at all accumulation potentials less negative than -0.6 V.) Both dissolved Cu2+ and “CuO” are products of the anodic reaction sequence, but at pH 8 the solubility of CuO is submicromolar, so a solid layer is likely to accumulate at the electrode. Reduction of this solid, and related oxidation products, can account for peaks labeled C1 and C4 during the cathodic scan (Figure 2A); e.g.:

“CuO”+ Hg + 2 H+ + 2e- f Cu(Hg) + H2O

(4)

In the Supporting Information, we present solubility experiments using the same chalcocite as above. The goal is to characterize dissolved Cu sulfide analytes produced by approaching equilibrium from undersaturation, avoiding high degrees of supersaturation attendant on mixing dissolved Cu2+ and HS-. The key experimental result is that 7494

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such analytes do not produce C3 peaks, although they do produce peaks very similar to C1 and C2. CuS Nanoparticles. A CuS sol was prepared by mixing equimolar HS- and Cu(II)-DCTA (17). Curve 1 in Figure 3 shows that two peaks at potentials of -0.65 and -0.99 V are observed with no stirring or accumulation period. These potentials are similar to those of the C2 and C3 peaks in Figure 2. Peaks C1 and C4 are absent because the anodic process described by reaction 3 proceeds to a negligible extent without an accumulation period. The substantial C2 peak in this figure is due to dissolved HS-, which exists at ∼10-5 M in equilibrium with amorphous CuS when free DCTA ≈ 10-3 M (see equilibrium constant data, 33). Voltammograms such as those in Figure 3 were obtained regardless of whether the sol had been freshly made or aged up to 1 day. Similar voltammograms were obtained using EDTA instead of DCTA. The C3 peak produced by nanoparticles (Figure 3) is broader and shifted relative to that produced by larger particles (Figure 2). On the basis of ref 29, we attribute these features to superimposition of peaks for sequential electroreductions, which progress from CuS through a series of sulfide phases (Cu1.5S, Cu1.8S, Cu2S) and arrive finally at Cu. Only the last of these reductions occurs in Figure 2 because the particles consist of Cu2S initially. Thus the shape and peak potential of C3 carries information about the composition of the electroactive Cu-S solid phases. When the solution represented by scan 1 of Figure 3 was purged with N2, the results represented by scans 2 and 3 in this figure were produced. Purging causes both C2 and C3 to decline while C5 grows. The latter is a CuDCTA2- reduction peak. Removing H2S causes the solution to become undersaturated with respect to CuS, allowing free HDCTA3- to attack and dissolve nanoparticles. Shea and Helz (17) found that CuS particles prepared by the method used for the experiment in Figure 3 are semispherical with radius 6-12 nm. The robust C3 peaks in Figure 3 show that such nanoparticles resemble larger particles voltammetrically, but can be distinguished by appearance of C3 under diffusion controlled conditions, without stirring or an accumulation period. Titrations of Cu2+ with HS-. With the foregoing background, titrations similar to those of Luther et al. (10, 11) were performed. Figure 4A presents voltammograms at several stages during titration of HS- into Cu2+. Increasing amounts of sulfide cause loss of broad anodic and cathodic peaks, near -0.1 V, which are related to dissolved Cu that has not reacted with sulfide. A sharp peak near -0.1 V, due to an unknown transient, is superimposed on the broad peaks in the middle two scans. A small C3 peak appears near -1.0

FIGURE 4. (A) Cyclic voltammograms at stages in a titration of sulfide into 10 µM Cu2+ contained in 0.05 M NaCl + 0.03 M NaHCO3 at pH 9. To minimize cumulative sulfide loss through oxidation and volatilization, a new solution of CuSO4 in supporting electrolyte was prepared and deaerated before each sulfide addition; Ei ) -0.05 V, td ) 0 s, scan rate 0.1 V/s. (B) Plot of C3 peak area normalized to its final value at equimolar Cu/HS (i.e., Q(C3)rel) vs the A1 peak area normalized to its initial value (i.e., Q(A1)rel). This shows that the C3 peak grows inversely to the A1 peak. V in the final three scans. Free sulfide, indicated by C2, does not appear until the equivalence point is passed (fourth scan). The small feature that we label C3 always appeared in our titrations as soon as sulfide was added. Figure 4B shows that it grew in proportion as the broad peaks associated with dissolved Cu declined during the titration. It was formed regardless of whether HS- was titrated into Cu2+ or Cu2+ was titrated into HS-; thus it forms and can be detected when either component is in excess. These C3 peaks are indicative of nanoparticles since they are observed without an accumulation period. The C3 peaks in Figure 4A differ in two ways from those obtained with CuS nanoparticles in Figure 3. First, they are conspicuously smaller, an effect caused by 100-fold less total analyte. Second, they are not accompanied by C2 peaks. As already stated, the C2 peaks in Figure 3 are due to free HSmaintained by CuS solubility in the presence of free HDCTA3-. The C3 signal can be enhanced by accumulation. Figure 5 shows that a much stronger C3 peak can be observed if a 30 s accumulation period is used. After accumulation at a potential of -0.05 V, a C2 peak also is prominent as a result of an anodic process like reaction 3. However, at more negative (less oxidative) accumulation potentials, the C2 peak no longer is produced during the 30 s accumulation period, and a larger C3 peak is observed instead. Figure 6 shows that the ∆Cu2+/∆HS- ratios found in our titrations are 1/1 within experimental uncertainty. In contrast, Luther et al. (10, 11) report ratios of 2/3 or sometimes 4/5 in later stages of titrations. In some preliminary experiments using Na2S‚9H2O from a long-opened bottle, we did observe lower stoichiometric ratios. The results in Figure 6 were obtained with Na2S‚9H2O from a freshly opened bottle. The implication of this finding is that ratios below 1/1 are controlled by oxidation artifacts (see Discussion). Note also that Figure 6 displays no discontinuity midway through the titration due to Cu2+regeneration, suggesting that this feature in Luther et al. (10, 11) also is an oxidation artifact. Because Figure 6 conflicts with Luther et al. (10, 11), we obtained an independent, more precise determination of the combining ratio of Cu2+ with HS- by classical titration methods (see Supporting Information). The combining ratio, ∆Cu2+/∆HS-, was very close to 1/1, ranging from 1.015 (

FIGURE 5. Cyclic voltammograms for solutions of 2 × 10-5 M CuSO4 and 2 × 10-5 M Na2S in supporting electrolyte after 30 s accumulation at three different potentials, Ei. 0.025 to 1.050 ( 0.028 in several experiments. A slight, but probably significant excess of Cu was found. Aged precipitates from this kind of reaction are often Cu-rich (33). Addition of Cu2+ to Adriatic Seawater. Figure 7 shows a titration of Cu2+ into a sample of muscilaginous, thiol-rich Adriatic Sea water that had been returned to the laboratory. The DOC concentration in this sample was 17.8 mg/L, mainly in the form of polysaccharides (34, 35). Adding Cu2+ to this sample quickly produces a strong C3 peak which grows at the expense of the C2 peak as Cu2+ is added. The C2 peak VOL. 39, NO. 19, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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Discussion Our principal findings are as follows. Copper sulfide solid phases are subject to reduction to Cu0 at potentials in the -0.9 to -1.0 V range (C3). The exact position and shape of this peak appears to carry particle-composition information, but we have not explored this systematically. Larger particles (e.g., from powders) sorb to the electrode slowly and thus require an accumulation period to be observable; nanoparticles (r e 10-8 m) sorb to the electrode rapidly enough to be observable without accumulation when total Cu concentrations are greater than 10-6 M. Although we have not explored detection limits for nanoparticles, the small sizes of C3 peaks in Figure 4A suggest that at submicromolar concentrations, accumulation periods will be needed even for nanoparticles. The titration experiments described in this paper are subject to a simple interpretation that is consistent with a considerable body of previous work (14-20). We propose that highly supersaturated mixtures of Cu2+ and HS- react in approximately an equimolar ratio as follows: FIGURE 6. Stoichiometry of a titration reaction (sulfide added to 10 µM Cu2+ in 0.05 M NaCl + 0.03 M NaHCO3 at pH 9). The integrated charge of the anodic, A1 peak (Figure 4A), normalized to the value at the start of the titration (i.e., Qrel) is plotted vs the concentration of sulfide added. The A1 peak is titrated to baseline by 10 µM HS-, demonstrating a 1:1 molar stoichiometric ratio. Contrast this result with Figure 4 in ref 10 and Figure 2 in ref 11.

FIGURE 7. Voltammetric scans in mucilaginous North Adriatic Sea water containing 1.4 × 10-7 M thiols (calibrated as sulfide). Scan 1 (dashed line): no Cu2+ addition. Scan 2 (solid line): 1.6 × 10-7 Cu2+ addition. Scan 3 (dash-dot line): 3.6 × 10-7 Cu2+ addition. Note that Cu2+ is added in stoichiometric excess to thiol sulfur without depleting it completely. This implies that nonthiol ligands are successfully competing for some of the added Cu2+ on the time scale of this experiment. Initial potential, Ei ) -0.20 V; scan rate, v ) 100 mV/s; accumulation period, td ) 60 s. here is due to dissolved natural sulfide and thiols in the sample. Figure 7 serves to make 3 points. First, solid products have formed in this natural water sample, as shown by appearance of C3 peaks. Second, it would be plausible that similar solid products would form in nature if metal-rich water mixes with water containing natural reduced S compounds. Rozan et al. (21, 22) apparently have described such products from natural waters. Finally, such products can be detected voltammetrically at least down to 10-7 M (as Cu). With improved detection limits, voltammetry could become a useful tool for studying nanoparticles in nature. 7496

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HS- + Cu2+ f 1/2Cu2S.S0(s) + H+ f CuS(s) + H+

(5)

The initial product is a brown sol, which transforms gradually to a green sol (18); the details of this process are not yet well understood. Samples that have been frozen in liquid N2 within minutes of mixing Cu2+ and HS- are found by X-ray spectroscopy to contain exclusively CuI and complementary amounts of S0 (36). Precipitates recovered from this process after flocculation initially are X-ray amorphous and have compositions in the range Cu0.9-1.4S (33). When stored dry, at room temperature, they transform to crystalline covellite over a period of months. Chalcocite (Cu2S) nanoparticles can be obtained by conducting reaction 5 in water-in-oil emulsions so that the S0 can escape to the organic phase around the aqueous micelles, thus blocking the second step (37). The novel ideas about copper sulfide cluster formation proposed by Luther et al. (10, 11) are based largely on voltammetric titration evidence that was not reproduced here (Figure 6). Those workers’ reliance on iodometric titrations to calibrate sulfide solutions probably led them to greatly overestimate their sulfide concentrations and consequently underestimate the ∆Cu2+/∆HS- stoichiometry. Iodine does not distinguish among HS- and its common oxidation products, S2O32- and SO32-. Additionally, if iodometric titrations were conducted in the mildly alkaline NaCl+NaHCO3 electrolyte that was used in most of their experiments, irregular titration stoichiometry, caused by partial oxidation of HS- all the way to SO42- (38), could cause further overestimation of HS-. Luther et al. (11) also used laser ablation mass spectrometry to support their view that clusters such as Cu3S30, Cu4S50, and Cu4S60 are significant species in aqueous solution. However, numerous workers who have studied gas-phase clusters using this approach emphasize that such clusters are products of the high-temperature plasmas created by the laser. These clusters retain little, if any, information about the molecular composition of the original sample (39-42). Using more appropriate mass spectroscopic methods, Sukola et al. (9) have been unable to find the clusters reported in refs 11 and 22. Where does this leave the question of the existence of aqueous copper sulfide clusters? Because small particles necessarily precede big particles in any precipitation process, there can be no doubt that subnanometer clusters form at least as transients in supersaturated solutions. However, in our opinion, no evidence yet proves that such transients have lifetimes great enough to make them significant

components of dissolved Cu on time scales of titrations in the laboratory. Whether such transients could persist in nature depends partly on consequences of slow particle growth rates at nanomolar Cu concentrations and on the possibility of cluster stabilization by natural surfactants. These effects deserve more study. Methods developed here might prove useful in this effort. Helz et al. (43) first explored whether subnanometer CuI (bi)sulfide clusters might exist, not as transients, but as stable species at much lower concentrations in equilibrium with Cu sulfide minerals. EXAFS spectra from 9 M NaHS solutions saturated with Cu sulfide minerals revealed that Cu occurred in the 2nd coordination shell to Cu. This proves that the predominant aqueous species were multinuclear. The measured Cu-S and Cu-Cu distances were similar to those in known Cu4(RS)62- clusters and different from distances in Cu-S solid phases or linear Cu2S(HS)22- complexes. However, owing to limits in spectroscopic sensitivity, they could determine nothing about Cu speciation in less concentrated HS- solutions, such as would be relevant in the environment. Thompson and Helz (44) explored the question of thermodynamically stable Cu-S clusters in less concentrated HSsolutions using solution chemistry methods. Their evidence favored existence of clusters, but their measurement precision was insufficient to exclude mononuclear and dinuclear complexes as predominant species. Existence of clusters as major dissolved Cu species in equilibrium with Cu sulfide minerals is suggested by both studies but remains to be proven. If Cu sulfide nanoparticles contribute significantly to the preservation of sulfide in oxic natural waters, as evidence in Rozan et al. (21, 22) might suggest, then mechanisms that retard their oxidation become an interesting problem for study. Oxidation rates observed in the laboratory differ immensely for reasons that require clarification (compare 9, 18, 19). Evidence in this paper raises a cautionary flag regarding several studies of stability constants for mononuclear Cu sulfide complexes (45, 46). These studies were based on titrations in which high degrees of supersaturation would have been attained. No evidence was presented that mononuclear complexes, rather than nanoparticles, were the products of the titration reactions.

Acknowledgments Support by Grants EAR 0229387 and EAR 9908532 from the U.S. National Science Foundation and Grant 098122 from the Ministry of Science and Technology of the Republic of Croatia is appreciated.

Supporting Information Available Experiments (a) to determine the voltammetric signature of Cu-S analytes obtained by dissolving Cu2S and (b) independently to confirm that the combining ratio of Cu2+ and HS- is approximately 1/1. This material is available free of charge via the Internet at http://pubs.acs.org.

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Received for review March 25, 2005. Revised manuscript received July 21, 2005. Accepted July 21, 2005. ES050586V