Lake Kinneret is a seasonally stratified freshwater lake located in northern Israel (32° 50′ N, 35° 35′E). The lake covers an area of 170 km² and is 40 m deep at its deepest point (Station A). Thermal stratification begins around April and ends during the winter (December-January). During stratification, oxygen depletion and sulfate reduction in the water column and sediments lead to the prevalence of euxinic conditions and an active reoxidative sulfur cycle (Knossow et al., 2015).

Anoxic water samples for aqueous incubations were taken in May 2016 from the water column at Station A below the chemocline (defined as the point at which O₂ became non-detectable, 17 m depth) at a water depth of 20 m. Oxygen concentrations were below detection (≤1 μM) and the sulfide concentration was 15 μM. Water was pumped from depth into covered glass containers that were sealed with a glass stopper, which prevented contamination of the samples by oxygen during transport to the laboratory, where the samples were processed the same day. Sediment samples for slurry incubations were taken from Station A in November 2016 from sediments underlying anoxic and sulfidic water. The cores were sealed, transported back to the laboratory, and processed the same day.

H₂S amendments were made from a stock solution of Na₂S•9H₂O, which was prepared in deoxygenated, deionised 18 MΩ (MilliQ®) water less than 1 h prior to addition. SO32-, S₂O32-, and S₄O62- amendments were also made from stock solutions of their sodium salts which were prepared in anoxic 18 MΩ water directly preceding addition to the experiment.

A polysulfide stock saturated with respect to elemental sulfur was made by preparing a mixture of 600 mM Na₂S•9H₂O and 6 M S⁰ in 50 mL deoxygenated deionized water. This solution was sealed, stirred and gently heated (40°C) for 4 h to (partially) dissolve elemental sulfur, then was allowed to stand overnight at room temperature to equilibrate. The pH was then adjusted to 7.4 (0.1–0.2 pH units lower than the pH of the experiments in order to prevent precipitation of S₈ upon addition) and stood for another 2 h to allow S₈ to precipitate and settle. The pH was confirmed, then the solution was filtered (0.2 μm) and used in experiments the same day.

Elemental sulfur colloids were prepared according to the method of Janek (1933) by the reaction of H₂S with SO32- under acidic conditions. 3.6 g Na₂SO₃ and 6.5 g Na₂S•9H₂O were dissolved separately in 50 mL deionized water. 1.5 mL of the Na₂SO₃ solution was added to the sulfide solution, followed by about 8 mL H₂SO₄ (25 %) added dropwise until the cloudiness imparted to the solution upon addition of the acid barely disappeared after stirring. At this point, 3 mL concentrated H₂SO₄ were added to the Na₂SO₃ solution, which was then poured into the sulfide solution, turning the solution a milky yellow color. This mixture stood for 1 h, then was filtered through a Whatman (Size 5) 12.5 cm paper filter. The filtrate was washed with deionized water to remove soluble polythionates, then was resuspended in 300 mL deoxygenated deionized water for use in experiments. The colloids created by this synthesis were previously characterized by Steudel et al. (1988), who suggested a micellular structure and determined a formula of x(NaHSO₄/Na₂SO₄)•yS_n•zNa₂S_mO₆ (n = 6–10, m = 4–16).

Sediment slurry experiments were prepared using anoxic, non-sulfidic sediment at a dilution of 1:1 (v/v) with anoxic, sterile water taken from the overlying water column. The added water was heat sterilized (autoclaved to 120°C) before addition in order to isolate the effects of the sediment microbial community from those in the water column. The sediment slurries (pH 7.5) were prepared in a glovebag under an anoxic atmosphere and allowed to rest for at least 12 h prior to the injection of the amendment solutions. Injections of anoxic amendment solutions were made following the conditions outlined in Table 1. All experiments were conducted in duplicate. Following injection, the slurry experiments were incubated at 25°C, the normal temperature of the chemocline of Lake Kinneret in the summer (Knossow et al., 2015), in the dark under gentle shaking and were returned to the glovebag for sub-sampling.

Aqueous incubation experiments were set up in 150 mL glass vials sealed with rubber stoppers and aluminum caps. One hundred twenty-five milliliter of sample water (pH = 7.6) was transferred to each vial under an anoxic atmosphere in a glove bag and allowed to rest at least 12 h. We note that although the water contained 15 μM sulfide at the time of sampling, at the time the experiments were begun, the sulfide concentrations were between 0 and 5 μM, which are typical of interface environments. Experiments were begun upon syringe injection of the amended species through the rubber stopper. Non-sterilized experiments were conducted in triplicates and an abiotic control for each experiment was constructed by heat sterilization of the water under anoxic conditions (autoclaved to 120°C) prior to amendment. The initial concentration of each amended species was chosen to be as close to an environmentally relevant concentration range as possible, but high enough to allow accurate measurement of its consumption over time (Table 2). Throughout the experiments, oxygen was monitored in sub-samples using a fiberoptic optode (detection limit 1 μM; Firesting Pyroscience) and no oxygen was detected during either the preparation or course of the experiments. All experiments were stored in the dark unless otherwise noted (i.e., Experiment 1).

Throughout the course of all incubation experiments, sub-samples were taken and sulfur speciation was quantified for each experiment (Table 1). All sub-samples were taken in a glovebag under an anoxic atmosphere (<0.1% O₂) to reduce the risk for oxygen contamination of the experiments and oxidation artifacts in the sub-samples. Sub-samples from the sediment slurries were taken using syringes and were filtered through a 0.45 μm prefabricated filter (Millipore) prior to analysis.

Due to the relatively long time scale of most experiments (days to week), it is probable that the microbial community changed in response to the substrate additions, and thus is not representative of in situ conditions. However, the lack of a lag phase in all incubations (with the exception of Experiment 6A) indicates that the capability to utilize these substrates is present and active, or is readily activated when they are provided.

Sulfide [operationally defined as S(-II) measured spectrophotometrically: ΣS(-II) = H₂S + HS⁻ + polysulfide S(-II)] was preserved in zinc acetate (20 % w/v) and measured using the spectrophotometric method of Cline (1969) with detection at 665 nm. The method detection limit is 1 μM.

Polysulfides were quantified via HPLC following derivatisation with methyl triflate (Kamyshny et al., 2004, 2006). Briefly, 0.1 mL filtered (0.2 μM) sample, 0.1 mL phosphate buffer (pH 7.6), and 6 μL methyl triflate were added simultaneously to 0.8 mL methanol. The derivatised samples were stored at −20°C until analysis. Concentrations of polysulfides of chain lengths 2–8 were determined in derivatised samples by reversed phase HPLC with UV-detection at 220 and 230 nm. The method detection limit is 3–10 μM depending upon chain length (Kamyshny et al., 2004).

Zero-valent sulfur (colloidal and dissolved S⁰, polysulfide S⁰, S₄O62-) was quantified as SCN⁻ by HPLC using a C-30 column modified with polyethylene glycol (5%) (Rong et al., 2005; Kamyshny, 2009). ZVS was converted to SCN⁻ via cyanolysis by injecting 5–10 mL of sample and 20 μL KCN (10% w/v) concurrently into 20 mL of boiling boric acid (1% w/v), after which the solution was returned to a boil and the volume was reduced to 5–10 mL. Corrections were made for contributions from S₄O62- and Sx2-, therefore ZVS concentrations presented here represent only dissolved and colloidal S⁰. Cyanolysis was chosen over extraction by organic solvents (e.g., chloroform or toluene) as determination of cyanide-reactive sulfur yields higher recovery of ZVS in Lake Kinneret waters, perhaps due to the presence of biologically produced hydrophilic sulfur (Knossow et al., 2015).

Tetrathionate was quantified immediately after sub-sampling in filtered samples using the same HPLC method described above for cyanide-reactive ZVS samples. The detection limit for this method is 0.5 μM. Tetrathionate concentrations in filtered subsamples were found to be stable over time scales of 1–3 h, which allowed sufficient time for accurate measurement by HPLC.

Thiosulfate and sulfite were quantified by HPLC following derivatisation by monobromobimane (Newton et al., 1981; Zopfi et al., 2004). The method detection limit for both S₂O32- and SO32- is 0.005 μM.

Sulfate concentrations were measured by ion chromatography (Metrohm) with a conductivity detector after filtration and dilution of sub-samples preserved in zinc acetate. Sodium carbonate/bicarbonate buffer was used as the eluent. The detection limit for this method is 10 μM.

All reactions were treated as pseudo-first order and consumption rates were calculated based upon the method of initial rates. Pseudo-first order behavior and kinetic constants were derived from plots of ln(concentration) over time.

This treatment of the data means that the kinetics may be described after the rate law given by Equation (3):

where C is the concentration of a particular species and k′ is the pseudo-first order rate constant. Equation 3 can be integrated to give Equation (4).

This assumes that the kinetics are dependent upon the concentration of the species of interest, and that all other factors are constant during the experiment (e.g., microbial biomass, chemical oxidants). As the electron donor or acceptor is unknown for the consumption reactions observed in these experiments, this yields the most realistic representation of the decomposition kinetics. It is important to consider, however, that as most of the reactions examined here are microbially mediated, it is therefore likely that the actual kinetics have a relationship to the substrate concentration more consistent with a Michaelis-Menten description.

During sediment incubations, the extremely rapid initial consumption of sulfide and polysulfide was likely due to abiotic reaction with iron minerals. The reactive iron content ranges from 50 to 100 μM Fe g⁻¹ (dry weight) at Station A (Eckert, 2000) and the formation of iron-sulfides is further suggested by the change in the sediment color from brown to gray/black during these incubations. The slower decrease in polysulfide concentrations after the initial consumption, however, may be due to either continued reaction with less reactive iron minerals or microbial consumption. Further work is required to constrain this, as polysulfides play an integral role in pyrite formation (Rickard and Luther, 2007), isotope exchange processes between reduced sulfur species (Fossing and Jørgensen, 1990; Kamyshny et al., 2014), and furthermore represent an important potential energy source for micro-organisms (Findlay, 2016).

It is important to note that the results of the sulfide and polysulfide incubation experiments are not directly comparable. Polysulfide concentrations are dependent upon sulfide and the rapid scavenging of sulfide by these sediments would result in a corresponding decrease in polysulfides. Therefore, the consumption rate of polysulfides should equal to or higher than that of sulfide under similar conditions. To avoid this being the primary process affecting polysulfide dynamics in these experiments, however, polysulfides were added at concentrations much higher than those of sulfide (Figure 1). The result of this is that polysulfide consumption appears slower relative to sulfide when corrected for concentration (Table 2), but this is an artifact of the experimental conditions and not an accurate description of the relative reactivity.

In water column samples, a clear biological influence on sulfide oxidation was observed in the light (Figure 3A; Table 2) over the dark rate in Experiment 2A, indicating that sulfide was oxidized phototrophically (either through phototrophic sulfur bacteria, or cyanobacteria, both of which are typically present at the chemocline, e.g., Rimmer et al., 2008). Despite this clear biological influence, the production of S₂O32-, the main oxidation product, was similar between the non-sterilized (0.2 μM h⁻¹) and control (0.3 μM h⁻¹) experiments (data not shown). This could be due to internal storage of oxidation products by phototrophic sulfur bacteria (van Gemerden, 1986), so that the solution chemistry reflects only the chemical oxidation in both instances. During sulfide oxidation under dark conditions in the sterilized control, S₂O32- concentrations increase with time in the control at a rate of 0.41 μM h⁻¹, whereas in the non-sterilized experiment they decrease at a rate of 0.051 μM h⁻¹ (data not shown), suggesting that although a biological impact on sulfide oxidation was not observed, microbial consumption of the products from the chemical oxidation occurred.

Polysulfide consumption in the aqueous experiments appears to be abiotic. Although the consumption rate in the non-sterilized experiments (39 μM h⁻¹) was higher than that in the control (21 μM h⁻¹), when variability between the triplicate non-sterilized experiments is accounted for, there is no statistical difference between the two rates (Table 2). The interpretation of these experiments is however complicated by the possibility of concurrent sulfate reduction (e.g., Figure 3B), which would strongly impact the concentration of sulfide and thus also polysulfide (Equation 1). If indeed there is biological oxidation, this could be partially masked by sulfate reduction in the non-sterilized experiments. Intriguingly, in the non-sterilized experiments S₂O32- concentrations consistently increased throughout the incubation period (from 6.7 ± 0.7 to 120 ± 60 μM), whereas in the control, concentrations were stable (2.3 ± 0.6 μM). This is in contrast to the sulfide oxidation experiments (Experiment 2), in which thiosulfate concentrations increased in the control and decreased in the non-sterilized experiments, and may therefore be a consequence of biological polysulfide oxidation, rather than sulfide oxidation.

This is the first report of directly measured rates of Scolloidal0 consumption under environmentally relevant conditions, although previous studies have indicated pelagic (Jørgensen et al., 1979) and sedimentary (Fossing and Jørgensen, 1990) S⁰ consumption. In the sediment incubations, rapid Scolloidal0 consumption similar to that observed for sulfide and polysulfide was likely abiotic, due to reaction with reactive iron minerals (pyritisation or absorption).

In the water column experiments, (Experiment 3A), Scolloidal0 consumption was the fastest of all amended species (normalized for concentration) and was biologically mediated (Figure 4A). One interesting feature of these experiments is that S₂O32- was also present in this amendment as a by-product of the synthesis, however, the rate of S₂O32- consumption in Experiment 3A is slower than expected based upon the kinetics determined from Experiment 5A (0.69 μM h⁻¹ ± 0.068). Moreover, a lag time was not observed in Experiment 5A (Figure 3E), as was observed in Experiment 3A. This, and a pause in Scolloidal0 consumption once S₂O32- consumption began (Figure 4A) indicate that a similar mechanism was responsible for both the use of both Scolloidal0 and S₂O32-, and that consumption of both species could not occur simultaneously.

One possible explanation for this observation may be the speciation of the sulfur colloids themselves. At 22 h, when S₂O32- consumption began, approximately 41 ± 4% of the Scolloidal0 had been consumed. This corresponds to the characterization of this synthesis by Steudel et al. (1988), which demonstrated that about 45% of the S⁰ present in the sols was in a form other than S₈, likely longchain polythionates. We thus hypothesize that this more reactive S⁰ was consumed first, followed by a switch to thiosulfate metabolism, and then again back to the remaining, less reactive ZVS (present as S₈) once thiosulfate concentrations were low. The second phase of Scolloidal0 consumption occurred three times slower than the initial consumption, further suggesting that two different species were involved. In contrast, S₂O32- concentrations were stable over the course of the experiment in the control (Figure 4B).

Consumption of S₂O32- and S₄O62- in the water column was largely a biological process (>90%; Figures 3E,F, Table 2). Sedimentary consumption was also likely predominantly biological, as both of these species are chemically stable under anoxic, non-sulfidic conditions. Notably, S₄O62- consumption was slower than S₂O32- consumption in the water column, whereas the opposite was true for sediment incubations. The rates of S₂O32- consumption observed in this study are notably lower than previous reports from both freshwater and marine sediments (Jørgensen, 1990a,b; Elsgaard and Jørgensen, 1992; Zopfi et al., 2004).

No oxidation products were observed in non-sterilized S₂O32- amendments in aqueous experiments. This could indicate oxidation to either S₄O62- or SO42- (which were not measured in these experiments), consistent with microbial S₂O32- oxidation in chemoheterotrophic bacteria. Disproportionation could also have occurred, followed by oxidation of sulfide (as discussed in the following section for SO32- amendments). In sediment incubation experiments, however, measurements of S₄O62- were made, and production was observed after about 35 h of incubation, indicating that S₂O32- oxidation to S₄O62- was responsible for at least part of the observed S₂O32- consumption, similar to experiments conducted in coastal sediments by Podgorsek and Imhoff (1999), in which S₄O62- accounted for up to 57% of oxidized S₂O32-. In our case, assuming that S₄O62- was only produced and not consumed, oxidation of S₂O32- to S₄O62- accounts for 10% of S₂O32- loss in the experiments. The remainder may be due to either reduction or disproportionation; previous studies using radio-labeled S₂O32- tracers demonstrated simultaneous oxidation, reduction and disproportionation throughout the sediment column (Fossing and Jørgensen, 1990).

This is the first report of S₄O62- consumption in aqueous environmental samples, and although the rate is slower than that observed in sediment incubations (this study, Zopfi et al., 2004), it is consistent with previous reports in its stoichiometry and microbial nature. S₄O62- can be reduced by a variety of different microorganisms, including but not limited to sulfate-reducing bacteria (c.f. Barrett and Clark, 1987). In the non-sterilized, aqueous incubations (Experiment 6A), S₄O62- consumption was accompanied by a nearly stoichiometric increase in S₂O32- (1.9; Figure 4C). This stoichiometry is consistent with S₄O62- reduction via tetrathionate reductase (Bak et al., 1993) after Equation (5).

where [H] represents the enzyme tetrathionate reductase. Based upon S₂O32- production, Equation (5) can account for 100% of S₄O62- consumption in these experiments.

S₄O62- also reacts with sulfide, forming S⁰ and which would also result in a 2:1 S₂O32-: S₄O62- stoichiometry (Equation 6; Sorokin et al., 1996; Podgorsek and Imhoff, 1999; Zopfi et al., 2004).

ZVS concentrations were highest at the start of the experiments (15 μM), before significant S₄O62- consumption was observed. In order to explain ZVS formation by Equation (6), 15 μM S₄O62- would need to have been consumed initially, which is not supported by the data. ZVS has been also found to form from the decomposition of S₂O32- due to acid generation by S₄O62- oxidation (Bak et al., 1993), however this is also not likely at the pH and buffering capacity of these experiments. This points to a source of ZVS not related to the microbial metabolism of S₄O62- (likely oxidation of sulfide in the initial sample prior to the commencement of the experiment) and indicates that S₄O62- consumption is due to microbial tetrathionate reduction via Equation (5).

Although S₂O32- was still the dominant product of tetrathionate consumption in the sediment slurry incubations (Figure 2B), the ratio between S₂O32- production and S₄O62- consumption was lower (0.5) than observed in the aqueous incubations. S₂O32- production after Equation (5) therefore accounts for only 25% of S₄O62- consumption. A lower ratio (1.5) results from S₄O62- disproportionation (Equation 7; Zopfi et al., 2004), however this process is also not consistent with the stoichiometry observed in the sediments.

There are several possible explanations for the lower production of S₂O32- than expected based upon the reaction stoichiometries in Equations (5, 7) that was observed in the sediment incubations. First, S₄O62- reacts with sulfide (Equation 6). Although sulfide was likely produced by microbial sulfate reduction, it was not observed in any of the sediment incubation experiments, as these sediments demonstrate very effective sulfide scavenging (Figure 1). Furthermore, apparent sulfate reduction rates were lower than S₄O62- consumption, similar to the observation made by Zopfi et al. (2004). It is therefore unlikely that Equation (6) was the dominant consumption pathway for S₄O62- in these experiments. It is more likely that S₂O32- formed during S₄O62- consumption was consumed biologically (e.g., via reduction or disproportionation).

Equations (2, 5) illustrate the interconnectivity between S₂O32- and S₄O62- production and consumption, as each is essentially the reverse of the corresponding process. Thermodynamic calculations indicate that consumption of S₄O62- is favorable to consumption of S₂O32- at all but very low ratios of S₄O62- to S₂O32- (≤1 × 10⁻⁹). Thus, in natural environments the relative concentrations of S₂O32- and S₄O62- should control the direction in which this reaction proceeds. It has been known from microbiological studies that the cycles of S₂O32- and S₄O62- are connected (Tuttle and Jannasch, 1972, 1973; Barrett and Clark, 1987; Bak et al., 1993), however, the role of S₄O62- in microbial metabolism and in the sedimentary sulfur cycle is still not clear. Nevertheless, the importance of S₂O32- metabolism has long been established in both freshwater (Jørgensen, 1990a) and marine (Jørgensen, 1990b) sediments. The observation of S₄O62- production during S₂O32- metabolism in this study and by Podgorsek and Imhoff (1999) thus points to a potential source of tetrathionate in sediments. Although the concentrations of S₂O32- used in these experiments are much higher than those found in most sedimentary environments (which likely is the cause for the production of detectable tetrathionate production), it is nevertheless likely that tetrathionate is produced under normal environmental conditions and Podgorsek and Imhoff (1999) observed its formation even in sulfidic sediments.

Based upon free energy yields, SO32- disproportionation (−233 kJ mol⁻¹; Equation 8) is expected to be the favored metabolism under anoxic conditions, followed by reduction via sulfite reductase (−171 kJ mol⁻¹; Equation 9, Krämer and Cypionka, 1989).

Sulfide was detected in neither sediment incubations nor non-sterilized and control water column experiments; however, the presence of S₂O32- in both sediment (Exp 13, Figure 2C) and non-sterilized water column (Experiment 7, Figure 4C) experiments may indicate its formation, as SO32- reacts with sulfide to form S₂O32- via equation 10 with a free energy yield of −167 kJ/mol (Krämer and Cypionka, 1989).

For water column experiments, based upon the rates of sulfide oxidation in the dark determined from Experiment 2, we calculate that anaerobic sulfide oxidation could account for the loss of all potential sulfide formed via disproportionation, thus masking its formation. The absence of major electron acceptors (O₂, Fe(III), MnO₂, NO3-) in these experiments makes disproportionation the most likely process responsible for SO32- consumption within the water column. The capacity for SO32- disproportionation is found among sulfate-reducing bacteria, which are known to be present in the chemocline of the lake (Eckert and Conrad, 2007). In sediment experiments, the extremely rapid removal of sulfide observed in Experiment 8 (Figure 1) can explain the very low sulfide concentrations measured in Experiment 13, making sulfite loss consistent with either disproportionation or reduction in both cases. The capacity for dissimilatory sulfite reduction appears to be predominantly found in sulfate-reducing bacteria (Barrett and Clark, 1987).

An alternative possibility is SO32- oxidation by sulfite oxioreductase via Equation (11), however, this reaction yields a much lower energy (−19.7 kJ mol⁻¹) than disproportionation (Krämer and Cypionka, 1989) and is thus a less likely explanation.

The results of these experiments provide several important insights into the cycling of intermediate sulfur species in natural systems and the biogeochemical controls on their consumption. First, consumption of all intermediate species, with the exception of polysulfide, appears to be predominantly biologically mediated in anoxic aqueous environments with low concentrations of trace metals (Fe ≤ 1 μM). Polysulfide concentrations will be controlled by their formation rate from sulfide and S⁰, balanced by geochemical and possibly biological consumption and are often not in equilibrium with S⁰ (Kamyshny and Ferdelman, 2010; Lichtschlag et al., 2012). As the water-columns of most stratified, non-polluted natural environments have low trace metal content, these results are widely applicable.

Second, in anoxic, non-sulfidic sediments, consumption of S₂O32- and S₄O62- will likely be microbially mediated, SO32- consumption will likely be both microbial and chemical (e.g., through reaction with organic matter) and the concentrations of S⁰ will be controlled by the presence and speciation of reactive iron (e.g., through pyrite formation). A comparison between the turnover rates measured here for S₂O32-, S₄O62-, and SO32- and those measured in previous studies from both freshwater and marine sediments (Table 2) shows that the rates measured for Lake Kinneret are generally lower, however the sequence of reactivity for these species, namely SO32- ≈ S₄O62- > S₂O32- is the same. Therefore, the patterns determined here likely also apply to the anoxic, non-sulfidic sediments that compose a large portion of coastal marine environments (Zopfi et al., 2004). In sulfidic sediments, S₄O62- and SO32- may react with sulfide, increasing the abiotic component of the consumption rate. In contrast, S₂O32- is produced even in sulfidic systems (Jørgensen, 1990b). Notably, the presence of sulfide will impact the thermodynamics of particular reactions, for example disproportionation of S⁰ or S₂O32- (Jørgensen and Bak, 1991).

Third, the speciation of zero-valent or elemental sulfur in a particular environment will expectedly have a strong impact on its reactivity. Therefore, bulk descriptions of S⁰ based upon extraction (e.g., in toluene or methanol) or even reactivity (cyanide-reactive sulfur) give only limited information regarding the biogeochemical reactivity or microbial availability of the measured sulfur in a system. The results of this study and previous experiments indicate that in sediments, elemental sulfur reacts quickly, either with iron to form pyrite (c.f. Rickard and Luther, 2007) or due to absorption onto mineral surfaces (Fossing et al., 1992). In contrast, in the water column, Scolloidal0 consumption is biological, with no significant chemical reactivity observed. Scolloidal0 is furthermore the most rapidly consumed species in aqueous experiments, indicating its potential importance as a microbial substrate. In contrast, S⁰ often accumulates in anoxic water columns and in marine sediments, suggesting that either its formation occurs rapidly, or that it is present in a less reactive form than the colloids synthesized here (e.g., as crystalline orthorhombic sulfur). This may also be impacted by how the S⁰ forms. Biologically produced S⁰ may be more hydrophilic and closer in reactivity to the colloids prepared here (Zöphel et al., 1988), and so would turn over rapidly, whereas inorganically produced S⁰ (e.g., from sulfide oxidation by oxide minerals) is hydrophobic, less reactive and therefore accumulates to larger crystals. This would indicate that it is the structure and bonding environment of the initial phases and not the size that control the reactivity.

Finally, we observe that the cycles of S₄O62- and S₂O32- are closely linked in both aqueous and sedimentary environments. In particular, the rapid consumption of S₄O62- in sediment experiments, which was also observed in marine environments, indicates a prevalent biological role for this intermediate. Moreover, the production of S₄O62- during thiosulfate consumption points to a potential source for low concentrations of S₄O62- in sediments, as S₂O32- consumption is significant throughout the sediment column in both freshwater and marine systems (Jørgensen, 1990a,b). Such loops within the oxidative sulfur cycle represent rate-determining processes that serve to retard the overall transformation of sulfide to sulfate.

The authors declare that the research was conducted in the absence of any commercial or financial relationships that could be construed as a potential conflict of interest.