concentration and form of dissolved sulfide in the oxic water column of the ocean

Marine Chemistry, 27 (1989) 165-177 165 Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands

Concentration and Form of Dissolved Sulfide in the Oxic Water Column of the Ocean

GEORGE W. LUTHER, III and ELIZABETH TSAMAKIS

College of Marine Studies, University of Delaware, Lewes, DE, 19958 (U.S.A.)

(Received June 27, 1988; revision accepted April 12, 1989)

ABSTRACT

Luther, G.W., III and Tsamakis, E., 1989. Concentration and form of dissolved sulfide in the oxic water column of the ocean. Mar. Chem., 27: 165-177.

During May through July 1987, we analyzed for hydrogen sulfide in the oxic water column of the eastern Mediterranean Sea (May-June) and the northwest Atlantic Ocean (July). Cathodic stripping square wave voltammetry was used to detect and quantify sulfide species. This is a direct method requiring no sample preparation. Our measurements show a consistent profile of 2 nM in the Mediterranean Sea, except for increases that occur at the oxygen maximum in the photic zone and above the hypersaline anoxic brine of the Bannock Basin (2.9 mM sulfide maximum). The measurements in the Atlantic Ocean are lower and more variable as a result of different source waters. Sulfide increases are noted near the surface and in cold pool waters along the East Coast of the United States.

Experimental manipulation of seawater samples suggests that the sulfide is bound covalently to a metal. The metal(s) is likely to have an electron configuration consistent with inert complexes, which do not dissociate readily to form free sulfide. Free sulfide reacts with iodate and oxygen in seawater. Sulfide in an inert metal complex should not dissociate and react with iodate and oxygen readily. We propose that the kinetic stability of sulfide-metal complexes is the primary control on the distribution and reactivity of sulfide in the oxic water column of the ocean, rather than thermodynamic considerations alone. This kinetic effect results from high ligand field stabilization energies associated with the electron configuration of the metal (e.g. d 3 and d 6 low spin octahedral metal complexes).

INTRODUCTION

From May through July 1987, we participated in two cruises to study sulfur and iodine speciation in the ocean. The first was to study the chemistry of the anoxic hypersaline basins (Tyro and Bannock ) located in the Mediterranean Sea, and the second was to study the chemistry of Atlantic Ocean waters at the shelf slope interface. One of our objectives was to determine the sulfide concentration in the oxic water column of the sea and to gain insight into the nature or species of sulfide present using polarographic techniques. Several

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DISSOLVED SULFIDE IN OXIC WATER COLUMN OF OCEANS 167

workers have hypothesized that sulfide should be present because of COS hydrolysis (Elliott et al., 1987 ) or because of metal complexation (Dyrssen, 1988; Elliott, 1988). Cutter and Oatts ( 1987 ) have measured sulfide (0.51 + 0.07 nM ) in the surface waters of the Atlantic Ocean via an indirect gas chromatographic method. Recently, Cutter and Krahforst (1988) have reported on sulfide dis- tributions in waters from the photic zone of the Atlantic Ocean off the south. eastern coast of the United States. To our knowledge, no one has at tempted to determine sulfide throughout the oxic water column of the ocean. This study provides a data set for sulfide throughout the oxic water column of the Medi- terranean Sea above the Bannock Basin {4000 m depth, Fig. 1A), and at several locations in the northwest Atlantic Ocean (Fig. 1B ) to a depth of 1500 m. Sulfide data were obtained by cathodic stripping square wave voltammetry (CSSWV). Because voltammetric methods offer the possibility of direct mea- surement without chemical pretreatment of the solution, it is also possible to gain information regarding the species of sulfide and to determine its concentration.

EXPERIMENTAL

Water samples were collected with Go-Flo bottles that had been acid washed to ensure against sulfide contamination. This was very important in the Med- iterranean Sea research because of previous sampling performed in the hypersaline anoxic brines. Subsamples were withdrawn from the Go-Flo bottles into 50-ml polypropylene syringes (Aldrich) via three-way plastic stopcocks. The Atlantic Ocean samples were also filtered through 0.4-/~m filters. A 10-ml aliquot of the subsample was pipetted directly into the polarographic cell. The aliquot was purged for 2 min with high purity Ar gas to remove oxygen. Sulfide was analyzed by CSSWV using an EG + G Princeton Applied Research model 384B-4 polarographic analyzer system, with a model 303A static mercury drop electrode in the hanging mercury drop electrode mode. For analytical purposes, sulfide was deposited for 60 s at a potential of -0 .40 V (versus the saturated calomel electrode (SCE)). A 2 mV s -1 scan increment with a 100-Hz square wave pulse and 50-mV square wave pulse height was used to strip the sulfide from the Hg drop (scan range -0 .40 to -0 .90 V). Analytical precision on five replicates of each subsample is generally + 10% (relative standard deviation). Precision on replicates of seawater obtained from four different Go-Flo bottles at the same depth (50 m) and location near the Bannock Basin was _+ 12% relative standard deviation {4.00 _+ 0.45 nM). Typically 0.5 nM of sulfide yields a current of 10 nA. We estimate a minimum detection limit (MDL) for 60-s deposition times on 10-ml samples of 0.1 and 0.05 nM at - 0 . 4 V and - 0 . 1 V deposition potentials, respectively. This compares well with the MDL for iodide using CSSWV (Luther et al., 1988). We chose to measure our samples in a polarographic cell that had been purged for 2 min with Ar. This should re-

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move any 'free' sulfide (H2 S, S H - ). In fact, no sulfide peak is normally observed until after purging to remove the oxygen that gives a background cur-- rent over the voltage range scanned. Thus, oxygen and iodate coexist with the sulfide present. The type of sulfide will be discussed in light of this and other experimental observations.

Sulfide standards were prepared from anhydrous sodium sulfide (Alia Prod- ucts). The sodium sulfide was dissolved in nitrogen-purged water. One pellet ( ~ 0.1 g) of' sodium hydroxide (Fisher) was added to each preparation to pre- vent loss of sulfide as H2S. Gaseous H2S and COS were purchased from Mathe- son Gas and used without further purification.

RESULTS AND DISCUSSION

Polarographic experiments

Figure 2 shows voltammograms of a seawater sample at different initial deposition potentials. Sulfide gives a wave near - 0.59 V at seawater pH, whereas iodide gives a wave at - 0 . 2 2 V (if Triton X-100 is added to the solution iodide gives a wave at - 0 . 3 1 V (Luther et al., 1988) ). Argon was used to purge the sample, typically for 2 min, and no other chemicals were added to the sample,

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Continued purging has little effect on the current of the wave. The sulfide peak is sharper at a - 0.1 V deposition potential than at - 0.4 V potential. This may be related to Cu 2+ also depositing at - 0 . 4 V. Because this is a cathodic scan (negative direction), sulfide is stripped off the Hg electrode, whereas Cu may remain in the Hg drop. A Cu/sulfide interaction on the Hg drop should result in a broadening of the sulfide peak, particularly if the Cu 2 + is bound to the sulfide in solution before the deposition step.

The following experiments were performed on seawater samples to verify that the peak is due to sulfide, which enhances the electrochemical oxidation of the Hg electrode (eqn. (1); Turner et al., 1975 ).

S H - + H g ~ H g S + H + + 2 e - (1)

Addition of I ml of 0.1 M sodium hydroxide to seawater samples causes a shift of the peak to near -0 .63 V, but no significant increase or decrease is observed in the current of the wave. Addition of excess acid causes a shift of the peak towards the iodide peak (near - 0 . 2 V), or the peak is not observed as a result of sulfide loss because of acidification and purging (the potential of the iodide peak is unaffected by acidification). The current of the peak decreases sub- stantially when the deposition time is set at 0 s rather than 60 s, verifying that the preconcentration step is an oxidation process at the Hg electrode. A mil- lion-fold excess of Zn 2+ ( >_ 200 ttl of 50 mM zinc acetate solution) added to seawater samples causes the peak to disappear, presumably because of ZnS precipitation or because of shifting the peak to more positive potentials (in many cases the peak shifted to near - 0.35 V on zinc addition). Addition of a surface-active agent, such as Triton X-100 (50 ttl of 0.2% solution), to 10 ml of seawater also causes a decrease in the current of the peak. Sulfide forms an HgS film on the mercury electrode (Turner et al., 1975 ) rather than a precipitate, which the halides form (Colovos et al., 1974). When Triton X-100 is added, it apparently prevents efficient deposition of sulfide on the Hg electrode at these low levels because of the Triton X-100's surface-active properties at the Hg electrode. The peak potential of H2S and COS, which are blown as pure gases through the purge port of the model 303A electrode and into 0.565 M sodium chloride solutions containing 4 mM bicarbonate, is identical with that in seawater. The similar potentials are likely fortuitous because sodium chloride solutions do not mimic seawater exactly and were not purified to remove trace metals. Sodium chloride solutions are used because seawater contains iodate which reacts with free sulfide (Zhang and Whitfield, 1986).

All of the above tests (chemical and electrochemical) demonstrate that the peak in seawater samples is due to a sulfide species. In the case of COS, it hydrolyzes to hydrogen sulfide in aqueous media (Greenwood and Earnshaw, 1984; Elliott et al., 1987). However, the sulfide peak measured in oxic seawater appears to be complexed (bound in a covalent chemical bond) rather than free because (1) the argon purging does not remove the sulfide peak readily, whereas

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solutions of H2S and COS are removed by purging and (2) available oxidants in seawater (02, IOa- ) do not react quantitatively with the sulfide. The reaction of iodate with sulfide is discussed below.

In a series of experiments we added sodium sulfide standards to ( 1 ) seawater and (2) solutions of 0.565 M sodium chloride containing 4 mM bicarbonate for the purpose of understanding the physico-chemical and electrochemical behavior of sulfide in seawater. When free sulfide is added to seawater for the purpose of standard additions (20-30 nM increments), no increase in current is observed for the sulfide peak until most of the iodate present reacts with the free sulfide. Iodate ranges from 300 to 460 nM in seawater, with the highest values observed at depths greater than 100 m (Elderfield and Truesdale, 1980; Luther et al., 1988). Figure 3 shows voltammograms of seawater beh~re and after excess free sulfide is added. The iodide peak increases after free sulfide addition. (In studies with excess S H - to IO:~ , >_ 90% of the IO:~- in seawater is converted to I- . ) This observation, that free sulfide reacts with iodate, gives evidence to support the complexation of sulfide to metals through a chemical bond (e.g. M-S) . In laboratory studies, organic thiols do not appear to react appreciably with IO:~ . However, thiols are not the likely cause of the sulfide

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peak because these compounds give peaks at more positive potentials than inorganic sulfide (Wang, 1985; Luther et al., 1986).

In bicarbonate-buffered sodium chloride solutions (pH 8), the addition of free gaseous H2S yields a peak near - 0.59 V. However, on continued purging with Ar, the sulfide peak decreases in current quickly because of the open system that is used for purging. With care in sample handling, standard curves using sodium sulfide standards can be made ranging from 5-10 nM to higher sulfide concentrations. A 10-nM solution gives a current of about 200 nA. Stan- dard additions of free sulfide in a closed cell, rather than in an open one, should yield very reproducible data for free sulfide (e.g. Opekar and Bruckenstein, 1984). However, our data from oceanic samples are obtained in an open cell system. These data obtained from ocean samples are more reproducible than those obtained with sodium sulfide standard in sodium chloride solutions (pH 8). In addition, the sulfide peak is quite stable because seawater samples stored in trace-metal-cleaned polypropylene bottles for 10 months at room tempera- ture still give the peak. The chemical and electrochemical experiments described above can be repeated on samples stored in the dark at room temper- ature. All of these experiments support the possible c0mplexation of sulfide by metals, as does the experiment described below.

Finally, we synthesized the metal sulfide complex, [ Cr (H20) ~SH ] 2 +, as per Ramasami and Sykes (1976), in order to compare its electrochemical behavior with that of the sulfide observed in the ocean samples. The complex was sep- arated by ion exchange chromatography at pH 1 and gave an identical visible absorption spectrum with that reported by Ramasami and Sykes (1976). The chromatography extract was used without further purification. The chemical and electrochemical behavior of the complex in sodium chloride/sodium bicarbonate solutions (pH 8) is almost identical to that of sulfide in this matrix. The peak for the complex is similar to the sulfide peak in seawater (Fig. 2). However, the peak is sharp at both -0 .1 and - 0 . 4 V deposition potentials, which suggests that the metal sulfide complex in seawater is principally due to a metal other than Cr. The addition of Zn (II) solutions (104-108 excess) shifts the peak to near -0 .35 V. Thus, Zn(II) addition to both [Cr(H20)sSH] 2+ and seawater solutions does not precipitate ZnS readily, as suggested by Dyrs- sen (1988). Methods to measure H2S via a ZnS solid preconcentration step should not be very efficient. Cutter and Krahforst (1988) noted poor H2S recovery with such a method.

Field observations

Figure 4 shows the data obtained in the oxic water column of the Mediter- ranean Sea above the Bannock Basin location (see Fig. 1A). All data were obtained during the day. The salinity ranges from 38.481%o at the surface to 38.665%o at depth. The sulfide levels are consistently near 2 nM except at two

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depths. Firstly, at the oxygen and fluorescence maxima (near 50 m) there is a significant increase in sulfide. This could be related to the production of sulfide where primary production is high. The processes responsible for the production and/or release of sulfide over the entire water column are unknown, although anaerobic processes may be occurring in marine snow or fecal pellets (Karl and Knauer, 1984; Alldredge and Cohen, 1987). The second increase in sulfide in the water column occurs at 3125 and 3150 m. This is about 100 m above the oxic-anoxic interface between the Mediterranean Sea and the hypersaline anoxic basins, which contain as much as 2.9 mM H2S. This increase could be due to flux of H2S from the bottom waters of the anoxic brines. It is also possible that H2S and metal sulfides released at mid-ocean rise vents can contribute sulfidic material for dispersal to the ocean in a similar manner.

Table 1 shows data obtained in the Atlantic Ocean at 10 locations. These locations (Fig. 1B) were chosen to understand the extent and the chemistry of the nutrient-rich cold pool waters. Their source is the Gulf of Maine (Church et al., 1984). Only stations GF 32 and PG 11B are not affected by these cold pool waters. At GF 32 (a gulf stream station), sulfide is highest in concentration in the surface, and there is a significant increase at the oxygen and fluorescence maxima, as in the Mediterranean Sea data set. However, there is a marked decrease in sulfide concentration with depth for stations PG-11B and GF-32. This decrease differs from the Mediterranean Sea profile. Although salinity varies more in the Atlantic Ocean than the Mediterranean Sea, salinity alone is not likely to be the determining factor for these observations.

The salinity data suggest different source waters for the Atlantic Ocean, i.e. shelf waters, slope waters, cold pool waters, Gulf Stream waters and North Atlantic deep waters. All of these source waters have different biochemical/ chemical characteristics including nutrient andpzoductivity variations (Church


TABLE 1

Sulfide profile data ( _+ 1 a) for 10 locations in the nor thwes t Atlant ic Ocean during July 1987

Sample location Depth Sulfide Salini ty 02 Temp. (M) (nM) (%o) (%) ( °C)

PG 10 5 0.949_+0.115 31.562 107.0 23.48 37.16.07 N 20 ~ 1.95 +_0.193 32.941 121.4 10.65 74.41.52 W 35 b 1.26 -+0.133 33.173 101.0 8.05

50 b 1.23 -+0.098 33.338 92.4 8.20

PG l l B 5 0.24 _+0.096 33.909 105.8 24.84 37.16.39 N 40 ~ 0.854-+0.097 35.953 107.4 19.15 74.25.14 W 225 0.33 _+0.10 n.m. n.m. 13.51

700 0.42 -+0.077 35.887 80.8 4.73 1170 0.17 _+0.063 34.963 85.0 4.16

PG 12 4 0.871 _+ 0.104 31.504 107.9 23.96 37.19.07 N 24 b 1.62 _+0.108 32.875 107.9 8.59 75.00.34 W 37 b 1.01 _+0.121 32.883 106.7 8.48

PG 25C 5 0.935_+0.158 31.370 107.7 22.87 37.49.59 N 34 b 1.54 _+0.162 32.896 93.7 7.00 74.22.35 W 56 b 1.11 _+0.119 32.923 91.7 7.02

PG 50 3 1.70 +_0.165 31.264 112.6 22.69 38.25.02 N 35 b 1.52 -+0.073 33.016 97.5 7.61 73.45.21 W 50 b 0.99 -+0.10 33.428 97.7 9.05

69 0.74 -+0.089 34.018 89.3 10.39

PG 54 5 1.69 -+0.71 31.212 110.6 23.04 38.35.02 N 27 a'b 2.26 _+0.140 32.806 97.1 7.51 74.01.38 W 50 b 1.34 +_0.034 32.797 82.0 6.99

GF 32 5 0.595 -+ 0.070 33.637 105.9 25.32 36.24.06 N 25 0.388 -+ 0.071 35.656 105.9 26.73 73.50.10 W 40 a 1.43 -+0.087 34.574 111.7 18.88

100 0.519 _+ 0.092 35.950 67.2 16.72 200 0.353 -+ 0.059 35.146 50.7 9.07 300 0.505 -+ 0.056 35.073 61.5 7.07 600 0.374 -+ 0.042 34.996 79.6 4.85

1000 0.590 -+ 0.042 34.958 85.8 4.14 1500 0.357 _+ 0.095 34.966 85.7 3.82

GF 42 5 0.670 _+ 0.047 31.849 106.1 23.90 36.56.10 N 25 0.793 _+ 0.020 35.383 117.0 19.49 74.39.21 W 40 b 1.00 _+ 0.070 33.687 95.0 9.04

88 0.774 _+ 0.074 34.820 88.4 12.15

GF 43 5 0.812 +_ 0.087 31.649 106.2 23.92 36.57.06 N 20 0.854_+0.054 35.243 114.5 21.22 74.43.53 W 50 b 1.02 -+0.137 34.034 96.7 10.06

80 b 0.963 -+ 0.092 34.200 91.2 10.62

GF 44 5' 0.953 _+ 0.133 31.699 106.6 23.80 36.58.15 N 15 0.864 -+ 0.119 34.414 118.0 20.85 74.48.39 W 355 1.51 -+0.162 33.309 110.3 8.90

535 1.16 _+0.247 33.627 92.6 8.90

aOxygen and fluorescence maxima. bCold pool waters. n.m. = not measured.

] 74 G.W. LUTHER Ill AND E TSAMAKI. ~,

et al., 1984 ). As an example of these characteristics, GF stations 42, 43 and 44 are along an east to west transect. Sulfide levels increase westward suggesting possible nearshore or estuarine processes affecting sulfide distribution. These observations are similar to those of Cutter and Krahforst (1988), who analyzed samples from off the southeastern coast of the United States. PG stations 10, 12 and 25C are generally similar in sulfide distribution to GF 42, 43 and 44. PG 50 and 54 are the stations that are most affected by the nutrient-rich cold pool waters. These stations also have the highest sulfide levels in the Atlantic Ocean samples that we analyzed. The high sulfide levels correspond to the lower water temperatures ( < 9 ° C ), and higher trace metal and nutrient levels (T. Church, personal communication, 1988) associated with the cold pool waters. These cold pool data suggest that sulfide results from decomposition processes.

Speciation of sulfide

From our voltammetric and chemical experiments, it is possible to comment on the form and reactivity of sulfide in the oxic water column of the ocean. Thermodynamic work by Dyrssen (1988), which neglects organic chelation, shows that sulfide should be bound by trace metals. Specifically, 95% of the sulfide should be bound by Cu(II) as dissolved [Cu(S)]o, and 5% by Hg(II) as [ Hg (S) ] o. Although thermodynamics explains the formation of a complex, thermodynamic concepts alone do not explain the reactivity of a bound sulfide complex once it is formed in the oxic water column. We address reactivity and kinetic stability of bound sulfide through knowledge of the electron configuration of a given metal and ligand field theory.

Our results show that the sulfide in seawater persists for months, even though 02, H202 and IO3- are present for oxidation. Also, our results show that iodate will react with free sulfide to form iodide and, perhaps, sulfate (we did not detect thiosulfate ). The rate of this reaction in our seawater samples is consistent with the laboratory studies of Zhang and Whitfield (1986).

From this and other observations we conclude that sulfide is bound in a chemical bond to metals. However, metal-ligand bonds are frequently labile (eqn. 2 ). If the ligand dissociating from the metal is sulfide, iodate in seawater would react with the sulfide {eqn. 3)

[Cu (H20)~(S)1°+ H+-, [Cu(H20)~] 2+ + SH- (2)

S H - + IO~ - - , I - + oxidized sulfur products (3)

Certain metal complexes are known to be kinetically inert as a result of the electron configuration of the metal. This inertness stems from a stronger-than- expected metal-ligand bond that does not dissociate readily. Example s of kinetically inert metal cations in octahedral symmetry are d 3 (Cr 3 + ) a n d d 6 (low spin) electron configurations. In square planar symmetry the d s configuration


(Pt 2+, Pd 2+ ) is kinetically inert (Huheey, 1983 ). As an example of slow ligand exchange, the rate of water exchange for Cr 2+ (d4; t2g3eg 1 ) is 1015 times faster than Cr ~+ (d3; t2g 3 ). This relatively slow rate of exchange is attributed to ligand field stabilization energies resulting from these electron configurations (Hu- heey, 1983).

Cu 2 + (d 9) may also give kinetically stable complexes as a result of the Jahn- Teller effect (Huheey, 1983). This effect is responsible for much of the behavior of Cu(II) complexes (e.g. Cu 2+ tends to form square planar, rather than octahedral, complexes and has very low formation constants on the addition of a fifth and sixth ligand). Thus, an 'idealized' Cu (S) o or Cu (SH) + complex ion should have less reactivity (more kinetic stability).

Stable sulfide complexes with the metal cations Cr 3+ (Ardon and Taube, 1967; Ramasami and Sykes, 1976) and Ru 2+ (Kuehn and Taube, 1976) are known in aqueous solution. The chemical behavior of these complexes should serve as an example for other metal sulfide complexes. Both of these complex ions are unstable in acid solution (pH 0-2) to oxygen in air. At pH 2, tl/2 for the Cr complex is 55 h (Ardon and Taube, 1967). Their reactivity increases with increasing acidity. The instability of [ Cr (H20) 5SH ] 2 +, or any other metal sulfide complex, results from the loss of H2S as the ligand on protonation of the SH- ligand (eqn. 4)

[M(H20)sSH] n+ + H30+-. [M(H20)6] (n+1)+ +H2S (4)

The increased acidity favors an Snl or dissociative mechanism. This mechanism explains the success of the method for H2S determination by Cutter and Oatts (1987). However, the final H + concentration of their samples is 0.5 M, which may not enable the recovery of all the H2S from strong metal sulfide complexes.

The kinetic inertness (stability) of such complexes to oxidation should increase with pH because a dissociative process is less favorable. To dissociate the sulfide ligand from the metal complex at higher pH, other mechanisms are required. We discuss the chelate effect and photochemical bond cleavage using Cu 2 + as the metal example because of the interest in copper complexation to sulfide in seawater (Dyrssen, 1988), and because our sulfide values in the Med- iterranean Sea parallel the copper values of H. de Baar (personal communication, 1988).

The S 2- ligand could be displaced from Cu 2+ (or any other metal ion) by multidentate ligands, such as humic acids and polypeptides. The complete complexation of Cu 2+ by humic acid (the chelate effect) would yield several water molecules and one SH- ligand (eqn. 5). The net result is an increase in entropy,

[Cu(H20)x(S) ]°+humic acid (HA) +H +-~Cu(HA) +x H20+SH- (5)

which makes zig more negative. The chelate effect should be important in the surface microlayer and in the upper I00 m of the water column where primary

176 (;,W. I ,UTHER I11 AND I~;. I ' S A M A K I N

productivity is. highest. The diel data of Cutter and Krahforst (1988) show that sulfide concentrations are lower in periods of daylight. This could be related to increased organic chelate production to displace sulfide, or to sulfur metal bond-breaking via photolysis. Both processes would result in dissocia- tion of sulfide from metal complexes and its release to the atmosphere as H~S at the surface microlayer, or the reaction of' free SH- with IO:~ to produce more reduced iodine, as is observed in the surface waters of the ocean.

CONCLUSION

Using a direct voltammetric method, we determined sulfide in the oxic water column of the ocean. Sulfide is present throughout the water column, but shows maxima at the oxygen maximum in the photic zone offshore, and in nutrient rich nearshore waters. Reduced sulfur as sulfide is in a bound state because it does not react with iodate in seawater, and it is not purged from the polarographic cell. It is probable that sulfide is bound to metals as predicted by Dyrs- sen (1988) and Elliott (1988). We cannot state to which exact metal (s) the sulfide is bound at this time, although copper is likely based on our experiments at different deposition potentials. However, sulfide should be bound to metals that are inert kinetically as a result of the electron configuration of the metal and ligand field effects. Metals complexed with sulfide as a ligand are also likely to be complexed with organic chelates, which should serve to stabilize further metal-sulfide bonds. More work is needed to establish the exact form of the sulfide, how sulfide is produced, and how long it remains in the oxic water column of the ocean. More data is needed that can couple trace metal, sulfide and iodine forms throughout the oxic water column of the ocean. Metal- sulfide complexes, which exhibit kinetic inertness, are probably even more important in sulfidic waters and may account for the enhanced solubility of some metals (compared with metal sulfide precipitation ) in these waters.

ACKNOWLEDGMENTS

The Mediterranean data set was collected as part of the Anoxic Basins Cruise (ABC). The ABC was financed by the Netherlands Council of Oceanic Re- search (NRZ) and was part of the sea-going programme of the Utrecht Work- ing Group in Marine Geosciences. Cornelis H. van der Weijden was a gracious chief scientist. G. de Lange, H. de Baar and J. Middelburg provided valuable assistance and encouragement throughout. NRZ and Netherlands Institute for Sea Research (NIOZ) technicians are thanked for technical assistance and Captain Blok and his crew of the R/V "Tyro" for their gracious cooperation.

The Atlantic Ocean data sets were collected with the enthusiastic assistance of T. Ferdelman and T. Church, the chief scientist. C. Culberson provided the oxygen measurements. C. Branson-Swartz provided valuable assistance on the


free sulfide methodology. Captain McCann and his crew of the R/V "Cape Henlopen" are thanked for their helpful cooperation. We thank D. Nuzzio of EG&G for helpful discussions and the loan of a model 384B-4 system.

Ship time was provided under NSF grant OCE-8541747 to T. Church. G.W.L. acknowledges support from NSF grant OCE-8696121 for all phases of this work and from the NOAA, Office of Sea Grant, Department of Commerce, under grant NA 86AA-D-SG-040 for the Atlantic Ocean phase of this work.

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