Sulfur in various valence states between sulfate (+6) andsulfide (–2) plays a key role in microbially mediated subsurfaceredox reactions, including those involved in the oxidation oforganic carbon (Kasten and Jørgensen 2000). Bacterial sulfatereduction (BSR) is estimated to account for ~50% of sedimen-tary organic carbon remineralization in modern marine sedi-ments (Jorgensen 1982), as well as consuming a significantproportion of subsurface methane via anaerobic oxidation(Niewöhner et al. 1998; Reeburgh 2007). Sulfide oxidation, inturn, is a key anaerobic microbial metabolism, linked to iron,manganese, and nitrate reduction (Burdige and Nealson 1986;Schippers and Jørgensen 2002; Gevertz et al. 2000). Inbetween sulfate and sulfide are intermediate valence state sul-fur species that are also involved in a myriad of subsurfaceredox reactions, linking the subsurface biogeochemical sulfurcycle to other key redox cycles, such as iron and manganese(e.g., Thamdrup et al. 1993; Böttcher and Thamdrup 2001).Ultimately, the oxidation of organic carbon and/or the burial
of reduced sulfur species (principally the mineral pyrite)impact the electron balance of Earth’s surface environment(Berner 1987; Canfield 2005). As a result of the wide impor-tance of the sulfur cycle, the subsurface redox transformationsof sulfur are of critical interest.
One of the primary tools for exploring the biogeochemicalsedimentary sulfur cycle is through the measurement of thesulfur and oxygen isotopic compositions of sulfate (δ34SSO4 andδ18OSO4, respectively). Of these two isotope systems, sulfur iso-topes have been far more widely applied (e.g., Rees 1973;Habicht and Canfield 1997; Brüchert et al. 2001; Habicht et al.2002; Canfield et al. 2006). Each change in the valence stateof sulfur—from the oxidized form of sulfate through to thereduced form of sulfide—partitions sulfur isotopes, such thatthe light 32S is continually concentrated in the reduced prod-uct (Canfield 2001).
In contrast, δ18OSO4 is a relatively new geochemical toolapplied to the subsurface sulfur cycle. Whereas δ34SSO4 yieldsinteresting insight on the overall throughput of sulfurthrough microbial communities, the δ18OSO4 better illumi-nates the internal cell dynamics of sulfur cycling (Mizutaniand Rafter 1973; Fritz et al. 1989; Brunner et al. 2005, 2012;Mangalo et al. 2007, 2008; Wortmann et al. 2007; Antler et al.2013). Oxygen atoms in sulfate molecules equilibrate intra-cellularly when sulfur is in the intermediate valence state ofsulfite (Kohl et al., 2012; Wortmann et al. 2007; Brunner et al.2012; Kohl et al. 2012). δ18OSO4 is useful for tracking the sul-fur atom as it transitions among its various valence states dur-ing the subsurface sulfur cycle. Additionally, oxygen isotopes
Controls on the abiotic exchange between aqueous sulfate andwater under laboratory conditionsVictoria C. F. Rennie* and Alexandra V. TurchynDepartment of Earth Sciences, University of Cambridge, Downing Street, Cambridge, CB2 3EQ
AbstractThe oxygen atoms in sulfate are known to exchange with water at low pH and at high temperature; howev-
er, it is unclear what the timescale for exchange is for the pH and temperature conditions commonly experi-enced in the laboratory. We present a time series of sulfate-oxygen isotope data for solutions with two differentsulfate concentrations (28 mM and 11 mM), at a range of low to intermediate pH values (1 to 5), using bothhydrochloric and acetic acid. Using water enriched in 18O, we show that there is negligible exchange of oxygenatoms between sulfate and water over the course of 390 days. We use the external uncertainty in these resultsto calculate a lower bound estimate on the timescale for oxygen isotope exchange under these conditions. Thelower bound of the timescale for oxygen isotope exchange between sulfate and water at laboratory pH is ~2 ×105 hours (~25 y), which is broadly in agreement with previous estimates. This result validates the use of δ18OSO4
as geochemical tool for a variety of solutions that are subjected to low pH at room temperature.
AcknowledgmentsWe are grateful to Adina Paytan and two anonymous reviewers for
improvements to the manuscript, and to David Hodell for assistancewith the oxygen isotope analysis of the water, and James Rolfe for tech-nical assistance with the running of the TC/EA. We are also grateful toOlivier Rouxel for helpful discussions, and Jenny Mills, Xiaole Sun, andJoseph Nicholl for their feedback on early drafts of the manuscript. Thiswork was supported by a NERC studentship NERC grant NE/H011595/1and ERC starting investigator grant 307582 to AVT.
in sulfate provide valuable insight into different pathways ofsulfide oxidation, both in the modern environment (Calmelset al. 2007), and possibly over geological time (Newton et al.2004).
When oxygen atoms are added to a sulfate moleculethrough sedimentary microbial redox processes, the resultingδ18OSO4 should remain “fixed” for the lifetime of that sulfateatom in the natural environment. This is because under openmarine pH and temperature conditions, abiotic sulfate-oxygenexchange is understood to be very slow relative to the resi-dence time of sulfate in the ocean (aqueous sulfate is esti-mated to fully exchange oxygen isotopes with water over 109
years under marine pH and temperature conditions [Zak et al.1980]; this is longer than the residence time of sulfate in theocean, which is ~107 years). Using δ18OSO4 rests on the assump-tion that there is also minimal oxygen isotope exchangebetween sulfate and water during the variety of conditionsimposed by sample storage and laboratory processing. Aftersampling, aqueous sulfate samples can be exposed to the addi-tion of acid, a common practice with pore fluids, seawatersamples, and culture media before analysis. Additionally, theextraction of sulfate from evaporites and carbonates by min-eral dissolution in acid at low temperature, exposes sulfate toacidic conditions. The isotope fractionation factors for oxygenisotope exchange between water and sulfate (SO4
2–), bisulfate(HSO4
–), or sulfuric acid (H2SO4) are not well constrained, but
are expected to be in the range of 18-40‰ (Zeebe 2010; Lloyd1968; Mizutani and Rafter 1969). Thus, if even partial oxygenisotope exchange between sulfate-oxygen and water-oxygendoes occur during sample storage or handling, it will overprintthe original δ18OSO4.
Oxygen isotope exchange between sulfate-oxygen atomsand water-oxygen atoms has been shown to occur in excep-tionally low pH (<0) solutions at 25°C and 100°C (with a half-life [time for half of sulfate-oxygen atoms to exchange—τ1/2] of~105–102 h, Hoering and Kennedy 1957— see Fig. 1). A linearextrapolation of these experiments to a pH value of 1, yieldshalf-life exchange times on the order of 107 h (~1000 y). Athydrothermal temperatures (100-300°C), oxygen isotopeexchange between sulfate and water has also been shown tooccur between pH 4-8 (τ1/2 ~ 104–100.5 h; Chiba and Sakai1985). Extrapolating these results to laboratory temperaturesand a pH of 1 yields τ1/2 estimates of ~ 103 h (using activitycoefficients from Wirth 1971). The discrepancy in these twoextrapolations for oxygen isotope exchange between sulfateand water at laboratory temperature and pH is likely to be dueto the very different ionic strength of the two experimentalsetups, and because the Chiba and Sakai (1985) equationsmake assumptions about the chemical species undergoing iso-tope exchange that may not be valid at different pHs.
Another oxygen isotope exchange study was performedwith samples across a large temperature range (25–448°C),
Rennie and Tuchyn Isotope exchange between sulfate & water
167
Fig. 1. Summary plot of previous data on the half-life of oxygen isotope exchange between sulfate and water using experimental data from Hoeringand Kennedy (1957). Lloyd, (1968), and Chiba and Sakai (1985), modified from Seal et al. (2000). Calculations of pH for Hoering and Kennedy’s exper-iments are from Seal et al, 2000. Theoretical relationships between pH and half-life for Chiba and Sakai’s solutions are also shown. Using their derivedequations, and assuming that H2SO4 is the rate determining step, we apply that equation to our laboratory temperature and pH range, assuming thesample solution composition as Chiba and Sakai (1985): ΣS = 0.01 mol/kg, ΣNa = 0.1mol/kg, taking activities from Wirth (1971), and assuming a reac-tion order of 1.
which suggested that sulfate-oxygen exchange with water wasstill appreciably rapid at low temperature and circumneutralpH (τ1/2 ≈ 10
8 h at pH 7 and 25°C – Lloyd 1968, see Fig. 1).Unfortunately, the chemistry of the solutions in this experi-ment (and therefore the pH) were not well constrained, andmay have been much lower than reported (Chiba and Sakai1985). Thus far, the above studies have not made it possible torule out oxygen isotope exchange between sulfate and water inthe laboratory. It is, however, critical to establish whether sul-fate-oxygen can be re-set by isotope exchange during labora-tory processing, because this often involves increasing the tem-perature and decreasing the pH of the original solution (Chibaand Sakai 1985; Hall and Alexander 1940; Garus et al. 1967).
In this study, we have tested for the likelihood of isotopicexchange between sulfate-oxygen and water-oxygen in aque-ous samples at room temperature, over a range of pH valuesand over a timescale common to laboratory procedures. Weuse water enriched in 18O, so that, even with the smallest esti-mated fractionation factor from the literature, oxygen isotopeexchange should be apparent. We estimate a lower bound onthe timescale for exchange under these conditions, based onour results.
Materials and proceduresExperimental setup
Two parallel experiments, using solutions of sulfate at twodifferent concentrations and sulfate-oxygen isotopic composi-tions, were monitored for 390 days. Each experiment con-tained duplicate samples at ~ pH 1, 3, and 5 in hydrochloricacid and ~ pH 3 and 5 in acetic acid, as well as a nonacidifiedsample. The samples were maintained at laboratory (i.e., room)temperature (the temperature in the laboratory varies between20 and 25°C), and sampled regularly for the oxygen isotopecomposition of sulfate. The oxygen isotope composition of thewater used was enriched to 20‰ by evaporation (water was leftto evaporate, and repeatedly sampled until it reached 20‰,
before being used in this experiment) and thus any exchangeover the course of the experiment should produce sulfate-oxy-gen isotope compositions increasing toward 40‰ or higher.The concentrations and initial isotope compositions of theexperimental solutions are displayed in Table 1.
The oxygen isotope composition of the initial aqueous sul-fate was measured by dissolving the starting sulfate mineralsand precipitating insoluble BaSO4 via the addition of saturatedbarium chloride solution. The sulfate in the experimentalsolutions was similarly extracted; 0.5 mL experimental solu-tion was removed and the sulfate precipitated as BaSO4.Oxygen isotope analyses
The resultant barite precipitates were cleaned of possibleBaCO3 contamination by addition of cold 6M HCl (removedrapidly to prevent possible exchange), followed by two rinsesin 18.2 MQ water, before being dried. Approximately 180 μgdry precipitate was then weighed into silver capsules andbaked overnight at 50°C to remove any adsorbed water. Cap-sules were crushed and introduced into a Thermo continuousflow Delta V Plus as CO, via a ThermoFinnigan thermal con-version elemental analyzer (TC/EA) at 1450°C. Samples weremeasured in duplicate, except for the solutions at pH 1 (HCl),which were measured in triplicate. Samples were bracketed insets of twelve using triplicate measurements of two standards,NBS-127 (δ18O = 8.6‰) and an internal standard (δ18O =12‰). Reproducibility, as measured using the standard devia-tion of the 6 bracketing NBS-127 samples was better than0.4‰.
Results are presented in delta notation relative to ViennaStandard Mean Ocean Water (VSMOW):
(1)
and reported in units of permil (‰).The oxygen isotope composition of the water was analyzed
by cavity ringdown spectroscopy on a Picarro, calibrated with
∂=⎛
⎝⎜
⎞
⎠⎟×
R -� R
R1000sample VSMOW
VSMOW
Rennie and Tuchyn Isotope exchange between sulfate & water
168
Table 1. Summary of the experimental setup, the starting solution concentration and the starting oxygen isotope composition of thesulfate and the water.
Experiment Sulfate mineral [SO4]2– (mM) δ18OSO4 (‰) δ18OH2O (‰) Vial number Acid pH
3 international standards, GISP, SLAP and SMOW, andreported in delta notation versus standard mean ocean water(VSMOW). Each sample was measured a total of 9 times, withthe first three measurements discarded and the remaining sixaveraged.
AssessmentThe oxygen isotope composition of sulfate in both experi-
ments was monitored over the course of 390 days. No sulfatemineral precipitation was observed in the experimental vialsover the course of the experiment. Results are presented inFig. 2, grouped by experiment and acid composition. Theδ18OSO4 of all samples analyzed are within 2σ of the originalsulfate-oxygen isotope composition (denoted by the graybars). The sulfate-oxygen isotope composition of the lowestpH samples in both experiments (Fig. 2b&e) does show anapparent ~ 1‰ increase in δ18OSO4 compared with the samplesin solutions at higher pH over the course of the experiment,
however these results are still within 2σ analytical error of theoriginal δ18OSO4. This potential increase in the oxygen isotopiccomposition of sulfate in low pH samples is in the correctdirection for oxygen isotope exchange, given the δ18O of thewater at 20‰; sulfate that had equilibrated its oxygen atomswith this water should have a δ18OSO4 approaching 44‰. How-ever, our signal is not distinguishable from the noise withinthe data.Analytical noise
The δ18OSO4 measurements for the duration of the experi-ments span the full range of ±2σ analytical error, which is cal-culated from repeated measurements of standards. Highlyvariable δ18OSO4 may be the result of contamination of theexperimental solutions over time, contamination during pro-cessing, water bound within the barite lattice, or mass spec-trometer-related instability. If the experiment vials havebecome contaminated, then it would be expected that theexperiment with lower sulfate concentrations (experiment 2)
Rennie and Tuchyn Isotope exchange between sulfate & water
169
Fig. 2. δ18OSO4 over time for 28mM (a-c) and 11mM (d-f) solutions of sulfate at pH values 1-5. The grey bars indicate the original isotopic compositionof the sulfate ±2σ. (0.8‰, n=10). Sample error bars are 2σ of the variation in the bracketing standards on the mass spectrometer (always taking the 2σfrom the most variable of the two groups of bracketing standards), except for pH1 and control solutions, where the error is 2σ from 3 sample replicates.Solutions were made up from water with a δ18O composition of 20‰ and sulfate with starting δ18OSO4 values of 14.6‰ (a-c) and 11.6‰ (d-f), so thatexchange would become immediately apparent (fractionation factors range from 18‰ – 40‰; Zeebe, 2010; Lloyd, 1968; Mizutani and Rafter, 1969).The solutions were sampled regularly for 390 days.
would show higher variability. Moreover, it is possible thatthis variability would increase over the timescale of the exper-iment, especially if we were introducing the contaminant dur-ing sampling. In our experiments, δ18OSO4 results are not morevariable in experiment 2 than experiment 1, and do not getsignificantly more variable with time, but instead remainwithin ±2σ of the original δ18OSO4. This suggests that theobserved variability is not the result of increasing contamina-tion of the sample vial over time, but is rather due to variableperformance in the TC/EA (although it is not possible to ruleout the contribution from lattice bound water).
TC/EA measurements have reported relatively poor exter-nal precision for nitrates, sulfates, and phosphates, and canvary from 0.1‰ (1σ) to 0.6‰ (1σ) within the same laboratory(Vennemann et al. 2002; Boschetti and Iacumin 2005). The 1σanalytical error reported here (0.4‰) is consistent with thisrange, especially as we report the standard deviation of δ18OSO4
on duplicate or triplicate (rather than quadruple) mea-surements. There are several hypotheses for the variable per-formance of the TC/EA-CFIRMS (continuous flow isotope ratiomass spectrometry), including memory effects resulting fromoxygen isotope exchange between newly generated CO andresidual BaO from previous samples during subsequent sampleanalyses (Boschetti and Iacumin 2005). The buildup of mem-ory effects was mitigated in our analytical setup by running ablank sample every ~ 12 samples. However, this may not besufficient to eliminate all memory effects, which may havecontributed to some of the variability we observe.
Another likely cause of the poor reproducibility of repeatedδ18OSO4 analyses is asymmetrical peaks, which are oftenobserved in the determination of δ18O via TC/EA-CFIRMS.Asymmetrical peaks result in the poor calculation of the sam-ple’s isotope composition, and can occur due to sluggishpyrolysis or poor flushing of the sample CO through theTC/EA and to the mass spectrometer (LaPorte et al. 2009). Inthis experiment, sample peaks did occasionally show moretailing near the end of a run—samples were analyzed as they
were generated over the course of the experiment, and so inany given reactor, samples could fall at the beginning, middle,or end of the life of the crucible. Due to the increased heliumflow restrictions that develop over the life of a crucible, this islikely to contribute to the variation in the analytical uncer-tainty over the course of the experiment.
Sluggish or incomplete pyrolysis can also be the result ofpoor placement of the graphite crucible, such that it does notlie in the hottest part of the furnace. Variations in the heightof the crucible result in a greatly decreased CO yield (to a ~75% yield—Boschetti and Iacumin 2005). This is unlikely tobe problematic in the TC/EA setup used in this study, wherethe packing depth has been optimized over years of analyses.
DiscussionTimescales for exchange over low temperature
Using the lack of change in the δ18OSO4 over the course ofthe experiment, we can calculate an upper limit for the rate ofoxygen isotope exchange between water and sulfate-oxygensuggested by our data. We perform this calculation using the“analytical error envelope” of our isotopic measurement as thebounding conditions for maximum possible isotope exchangebetween sulfate and water that could have occurred over the390 days. We calculate the error envelope using the startingvalue for δ18OSO4 (14.6‰ and 11.6‰ for experiments 1 and 2,respectively), and assuming that the final δ18OSO4 is within0.8‰ of the starting composition after 390 days as shown inFig. 3.
We calculate the timescale for oxygen isotope exchange (orlack thereof) between sulfate and water using the followingequation, derived initially by Chiba and Sakai (1985):
(2)
where αe αf , and αi are the oxygen isotope fractionation fac-tors at equilibrium, the end and the start of the experiment,respectively. R is the overall rate of isotope exchange, and X
α α
α α
⎛
⎝⎜
⎞
⎠⎟ = −
⎛
⎝⎜
⎞
⎠⎟ tln
-
-R
4X+Y
4XYye f
e i1
Rennie and Tuchyn Isotope exchange between sulfate & water
170
Fig. 3. An estimate of the maximum possible oxygen isotope exchange between sulfate-oxygen and water over the duration of the experiment. Noisotopic exchange that was distinguishable from the starting isotope composition by 2σ was observed, thus the maximum possible isotopic excursionmust be ±2σ of the original oxygen isotope composition. We therefore take the maximum to be δ18Ooriginal SO4 + 2σ.
and Y are the total number of sulfate and water molecules inthe solution, and y1 is the mole fraction of those water atomsthat have a 16O atom (It can be assumed that y1 ≈ 1). Whenhalf of the sulfate atoms have exchanged oxygen atoms withwater, Eq. 2 becomes
(3)
Our experimental set-up is clearly far from isotopic equilib-rium for the sulfate-water system. However, the half-life of theoxygen isotope exchange can still be calculated, and a plot ofln (αe – αf)/(αe – αi) against run time, t, should produce astraight line with a slope of – R(4X + Y)/4XY). This derivationis valid irrespective of the order of the reaction (Chiba andSakai 1985) and further assumes that only one sulfate speciesis undergoing exchange, neglecting exchange from co-occur-ring sulfate species (e.g., both sulfate and bisulfate undergoingexchange simultaneously). At a pH 1 and 25°C, the majorityof the sulfate in solution is in the form of HSO4
– (see Fig. 4),and so this derivation is likely to hold.
We assume that the final experimental δ18OSO4 is elevatedby 0.8‰ (+2σ analytical error) after 390 days. We furtherassume a range of possible oxygen isotope equilibrium frac-tionation factors for sulfate and bisulfate, respectively, of18–23‰ and 28‰ (Zeebe 2010; Lloyd 1967; Mizutani andRafter 1969) to estimate the fastest possible half-life. Usingthis we estimate the half-life for isotopic exchange for the two
experiments at pH 1, as illustrated graphically in Fig. 5. Ourcalculation suggests a half-life for oxygen isotope exchangebetween sulfate and water of 2.2 × 105 – 2.9 × 105 h, for a rangeof oxygen isotope fractionation values between sulfate andwater at room temperature. These calculations suggest that itwould take at least 25 years for aqueous sulfate at low-to-inter-mediate pH at room temperature to isotopically exchange halfits oxygen atoms with water. Our lowest calculated timescalefor oxygen isotopic exchange between sulfate and water is twoorders of magnitude faster than the timescale of isotopeexchange first calculated by Hoering and Kennedy (1957).However, our calculation only provides a lower limit on thetimescale for exchange. This is because our extrapolation ofthe oxygen isotope ratio of the experimental solution assumesthe maximum possible oxygen isotope exchange between sul-fate and water within the duration of the experiment (>1 y),when, in fact, there is little to no variation in δ18OSO4 of thesolutions.
The calculated half-life for experiment 1 is lower than thatof experiment 2, for each value of the oxygen isotope frac-tionation between sulfate and water that is assumed. Giventhe uncertainties involved in this extrapolation, it is not pos-sible to determine whether this difference is significant, how-ever it is possible that the reaction rates for experiment 1 and2 should differ slightly, because in each experiment sulfate iscomplexed by a different cation (magnesium and calcium,respectively).
( ) = −⎛
⎝⎜
⎞
⎠⎟ln 0.5 R
4X+Y
4XYt 1
2
Rennie and Tuchyn Isotope exchange between sulfate & water
171
Fig. 4. Bjerrum plot of the relative proportion of sulfate as H2SO4 (solid line), HSO4– (dashed line) and SO4
2- (dot/dash line) for 25°C (black lines) and200°C (grey lines). Dissociation constants at 200°C are taken from Ohmoto and Lasaga, (1982). The shaded grey region is the pH range investigated inthis study.
This result confirms that oxygen isotope analyses of sulfateare not re-set by exchange with water at room temperaturewhen being processed in the laboratory over short to interme-diate timescales, even at a pH of 1. This lends additional con-fidence to the robust nature of reported δ18OSO4 for a variety ofnatural systems.
Comments and recommendationsThe oxygen isotope composition of aqueous sulfate in
medium to low pH (1-5) solutions at room temperature doesnot show isotope exchange with water on short to mediumtimescales. This rules out abiotic sulfate-oxygen exchange as asignificant contribution to observed sulfate-oxygen isotopetrends measured, when working at ambient temperatures andmoderate pH values.
This is consistent with timescales for exchange extrapo-lated from rates calculated at hydrothermal temperatures. Thelower bound of sulfate-oxygen isotope exchange at pH 1.3 at25°C is ~ 2 × 105 hours. This lower bound estimate on the τ1/2is likely to be much smaller than the true half-life for sulfate-oxygen isotope exchange, because of the relatively shorttimescale of this experiment. This result engenders confidencein low temperature laboratory procedures when analyzingδ18OSO4 in low pH solutions.
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Submitted 28 October 2013Revised 6 February 2014Accepted 13 March 2014
Rennie and Tuchyn Isotope exchange between sulfate & water
