Acidification of Europa's Subsurface Ocean as a Consequence of Oxidant Delivery

Abstract

Oxidants are formed at the surface of Europa and may be delivered to the subsurface ocean, possibly in great quantities. Whether these substances would be available for biological metabolism is uncertain, because they may react with sulfides and other compounds to generate sulfuric and other acids. If this process has been active on Europa for much of its age, then not only would it rob the ocean of life-supporting oxidants but the subsurface ocean could have a pH of ∼2.6, which is so acidic as to present an environmental challenge for life, unless organisms consume or sequester the oxidants fast enough to ameliorate the acidification. Key Words: Europa—Ocean chemistry—Oxidants—Sulfate—Peroxide—Oxygen—Sulfuric acid. Astrobiology 12, 151–159.

1. Introduction

S trong oxidants, compounds capable of receiving electrons from other compounds, are rare in the Solar System due to the abundance of chemical reductants such as H and C, which rapidly react with oxidants to form oxides that include water and carbon dioxide. Whether a compound is an oxidant is dependent on the electron availability of the local environment, and most oxidants are formed by non-equilibrium processes. Some Solar System oxidants are likely produced by high-energy events that occur in atmospheres, for instance, the production of trace O₂ on Mars (McElroy et al., 1977). However, for the three large Solar System bodies that have thus far been found to be rich in strong oxidants, Earth, Europa, and Ganymede, production occurs at the surface. The source of oxidants on Earth is oxygenic photosynthesis, whereas for Europa and Ganymede it is irradiation of the icy crusts by high-energy particles.

Europa's oxidants include a thin molecular oxygen (O₂) atmosphere (Hall et al., 1995), hydrogen peroxide (H₂O₂) and O₂ as constituents of the icy crust (Carlson et al., 1999), and various oxidized sulfur species ( \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{\;\;2 - }$$ \end{document} compounds) on the surface (McCord et al., 1998: Carlson et al., 1999). Both O₂ and H₂O₂ probably formed by the radiolysis of ice at Europa's surface (Carlson et al., 2009; Paranicas et al., 2009) and are strong oxidants in most environments. Sulfates, as sulfate salts (McCord et al., 1998, 2010) and sulfuric acid (Carlson et al., 1999, 2002), are distinctive components of Europa's surface. Also present are weak oxidants such as SO₂, S, and CO₂. Possible sources of sulfur compounds are substances delivered from Io or endogenic material, the latter of which is suggested by association with surface features that may represent connections to the ocean. Although it is unclear whether these sulfates are actually capable of oxidizing other compounds on Europa's surface or subsurface, they contribute to environmental conditions more oxidizing than those common elsewhere in the Solar System.

Surficial oxidants are likely transported to the subsurface ocean, possibly in substantial quantities. There they encounter reductants and modify the ocean chemistry. Here, we investigate those effects and show how reaction of the oxidants with sulfide may acidify Europa's ocean, with potentially significant consequences for biological systems in the ocean.

2. Fluxes of Oxidants and Reductants to Europa's Subsurface Ocean

The rate of delivery of oxidants to Europa's ocean depends on the net rate of radiolytic production near the surface and on geological processes in the crust that may then deliver the oxidants down through the ice to the ocean below. Hand et al. (2007), who built on approaches by Chyba and Phillips (2001) and Cooper et al. (2001), estimated the amount of oxygen and H₂O₂ that accumulates as radiolytic chemistry proceeds and impact gardening protects the products by burial within the top few meters of ice. The dominant oxidant is O₂, with 3–10% H₂O₂ (Carlson et al., 1999). Hand et al. (2007) then assumed, in effect, that after the quantities increased over a certain characteristic period of time, the oxidants would be delivered to the ocean, after which the near surface would begin to be reloaded for the next delivery. Hand et al. suggested that a plausible delivery timescale is the apparent age of the surface, ∼50 million years, a figure based on constraints from the small number of craters on Europa (Zahnle et al., 2003). However, the process of delivery and, hence, the context of this timescale was not defined. With this assumption, they calculated that an average rate of ∼4×10⁹ mol/yr of O₂ and H₂O₂ would be delivered into the ocean.

Hand et al. noted that shorter delivery intervals would result in a greater average delivery rate; longer intervals would yield slower delivery into the ocean. They also noted that reaction with endogenic reductants would consume the oxidants, a process we consider in detail in this paper. These endogenic reductants are generated by fluid-rock interactions and include CH₄, H₂, H₂S, and Fe²⁺. With production rates roughly estimated as 1.5×10⁸, 3.6×10⁸, 1.0×10⁹, and 9×10⁷ mol/yr, respectively (McCollom, 1999; Hand et al., 2007), Hand et al. calculated that a substantial fraction of oxidants would remain in the ocean even after the reductants are consumed. These calculations were expanded by Hand et al. (2009) to include “oxidized,” “reduced,” and “biological” ocean compositions, as well as the effect of exogenous organics.

The oxidant delivery process assumed by Hand et al. (2007) may be an oversimplification for various reasons. For example, the timescale for erasing craters may not be the same as the oxidant delivery interval. Moreover, the geological processes on Europa that modify the surface probably do it on a continual basis, whereas the scenario by Hand et al. (2007) implicitly involves a gradual surface buildup of oxidants punctuated by infrequent sudden delivery to the ocean, which seems unlikely. Without the substantial rates of delivery assumed by Hand et al., oxidants that reach the ocean would be quickly consumed by the endogenic reductants as well as by accretionary organics (Zolotov and Shock, 2004). Moreover, Zolotov and Shock suggested that O₂ and H₂O₂ would escape from the ocean during melt-through events (Greenberg et al., 1999) that expose the ocean to the surface.

More recently, Greenberg (2010) considered the various geological processes that might be operating to bury near-surface oxidants and work them down to the ocean, including ridge formation and surface burial. That analysis suggested that the oxidant-rich layer near the top of the ice crust gradually thickened and occupied an ever-greater fraction of the crust until by ∼2 Gyr it occupied the entire crust. Up to that point, little oxygen reached the ocean. Thereafter, O₂ and H₂O₂ (mostly O₂) were delivered at a rate of ∼1–3×10¹¹ mol/yr, assuming the same radiolytic production rate of Hand et al. (2007). While the actual delivery rate remains uncertain, the plausibility of such high rates motivates consideration of the consequences for the chemistry of the ocean and for any life there.

3. Modeling the Ocean Chemistry

Oxidants that enter the subsurface ocean would be expected to affect its pH. Zolotov and Kargel (2009) showed that, with only a small amount of oxidant delivery, the pH would be high. But with the substantial delivery flux that may be possible, oxidants that react with the endogenic reductants, especially sulfide or organic compounds, could acidify the ocean.

Two reactions of special note include the oxidation of methane \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm CH}_4 + 2{ \rm O}_2 \rightarrow { \rm H}_2{ \rm CO}_3 + { \rm H}_2{ \rm O} \tag{1}\end{align*} \end{document}

and the oxidation of hydrogen sulfide \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm H}_2{ \rm S} + 2{ \rm O}_2 \rightarrow { \rm SO}_4^{\;\;2 - } + 2{ \rm H}^ + \tag{2}\end{align*} \end{document}

The latter reaction is known to be important in acidification of terrestrial geological environments. Experiments by Millero et al. (1987) and by Jennings et al. (2000) showed that the reaction of O₂ and H₂O₂ with H₂S and sulfide minerals pyrite or pyrrhotite decreases the pH of solutions to values between 2 and 3, and that these reactions occur rapidly [t _1/2∼10–100 h for sub-mM concentrations of H₂O₂ or O₂ and H₂S (Satterfield et al., 1954; Zhang and Millero, 1993)]. In those experiments, approximately two H⁺ ions were formed for every sulfur oxidized.

Here, we evaluate the generation of acid by reaction of oxidants with a wide range of plausibly relevant reductants, and we quantify the accompanying acidification of Europa's subsurface ocean by calculating the equilibrium chemistry. Computations were carried out with the program HSC (version 7.0, Outokompu Research Oy)¹. This code uses the GIBBS energy solver (White et al., 1958) to determine equilibrium concentrations and has been used previously to constrain sulfur chemistry in the Solar System (Pasek et al., 2005). The behavior of aqueous species in water is approximated by using the HSC Chem AQUA module that applies the Davies model (extended Debye-Hückel), the semi-empirical Pitzer model (with binary interactions only), and Harvie's modification of the Pitzer model (binary and ternary parameters). The code allows for the injection or removal of species and computes the resulting changes in solution chemistry with respect to time. Reaction kinetics are ignored in this study, because the reactions of major note (oxidation of reduced S species) are rapid and thus quickly reach equilibrium even at 273 K (Millero et al., 1987; Zhang and Millero, 1993; Jennings et al., 2000).

In our first model system, the oxidant flux into the ocean reacts with accumulating endogenic reductants dissolved in water, but we assume negligible chemical reaction with the underlying rock itself. Similarly, Kargel et al. (2000) considered a scenario in which the interaction with rock was minimal. As discussed in Section 5.1 below, it is quite plausible that physical conditions minimize the chemical interface at the ocean-rock boundary. Even if the rock is able to react freely with ocean water, this first type of model system might apply in the upper layer of the ocean (Section 5.1).

In this system, the initial abundances of H and O were set equal to the mass of Europa's ocean, which we set to 3×10²¹ kg and corresponds to a 100 km ocean depth (Anderson et al., 1997). The salinity was set equivalent to the salinity of “Europan Seawater” in McCollom (1999) for Na, K, Cl, and Mg. In this system, the dominant anion is Cl^-. The abundances of C, S, Fe, H, and O were then modified to account for the flux of substances that enter the ocean. First, reductants from the interior (H₂, CH₄, H₂S, and Fe²⁺) were added according to rates from McCollom (1999) and Hand et al. (2007). Then, after 2 billion years, oxidants from the crust (O₂ and H₂O₂) were added, according to the timing and rate calculated by Greenberg (2010), which were based on the radiolytic production model of Hand et al. (2007) and on an assessment of the dominant geological processes. We also tested cases with much slower delivery of oxidants into the ocean, but always commencing 2 billion years after the start of the endogenic reductant flow. These calculations were performed at 0°C and 120 bar (equivalent to 10 km of ice load) and solved for H⁺ concentration with respect to water. Changing the pressure of the system to larger pressures (1–2 kbar) was not found to change the results appreciably. Similarly, the solution chemistry is insensitive to the volume of the ocean (e.g., reducing the ocean depth from 100 to 50 km decreased the pH at 4.5 Gyr by only 0.04 units).

Our second type of model system includes chemical reaction with the rock on the seafloor. The rock would be expected to neutralize the acid produced. Our computational experiments quantified this effect for various quantities of reactive rock and showed the chemistry of the ocean after such reactions. We considered cases with basaltic and with ultramafic rock, using compositions given by Turekian and Wedepohl (1961). For each rock type, cases were examined with reactive rock volumes set equal to the upper 1 m, 1 km, and 25 km of sub-oceanic rock, for a total of six model systems. The interaction depth of 25 km corresponds to the maximum depth of porosity in the rock, according to Vance and Goodman (2009). For these experiments, we exposed to the rock seawater in which new compounds had first been generated from the reaction of oxidants with reductants. In fact, as shown in Section 4, for the rock-free cases, the pH reached a near steady-state very soon after the oxidants began to arrive in the already reductant-rich ocean. For our cases with the reactive rock component, we started with the ocean composition in this steady-state obtained from the rock-free cases after 4.5 billion years (i.e., 2.5 billion years after the oxidant flux begins entering the ocean). This choice of initial ocean composition is conservative in the sense that it maximizes the role of the rock exposure, because the total quantity of oxidants is limited to the amount delivered over the last 2.5 billion years, as opposed to 4.5 billion years of oxidant flux. Note too that the final ocean composition and pH at 4.5 Gyr would be the same if the rock were introduced earlier, because the total elemental budget at 4.5 Gyr would be unchanged. The calculations were performed at 2 kbar (expected under 150 km of water) for a range of temperatures from 0°C to 250°C.

For all the cases considered here, once the abundances of the elements were supplied, the HSC program then calculated the distribution of elements among the various chemical species (listed below) by solving the minimum thermodynamic energy state coupled to mass balance calculations. With the code, we tracked the quantities of the following compounds, which represent contributions from basaltic or ultramafic silicates, as well as from a suite of sulfate minerals that may form during oxidation of sulfides and iron sulfide minerals, as well as iron oxides, phyllosilicates, and carbonates: Al₂O₃, Al₂SiO₅ (andalusite), Al₂Si₂O₅(OH)₄ (kaolinite), CaAl₂Si₂O₈, CaCO₃, CaMgSi₂O₆, CaSO₄×2H₂O, Fe, Fe(OH)₂, Fe(OH)₃, Fe_0.877S, Fe_0.945O, Fe_0.947O, Fe₂(SO₄)₃, Fe₂O₃, Fe₂O₃×3H₂O, Fe₂O₃·H₂O, Fe₃O₄, FeCO₃, FeO, FeS, FeS₂, FeSiO₃, FeSO₄·4H₂O, FeSO₄·7H₂O, FeSO₄×H₂O, KAlSi₃O₈, MgCO₃, MgO, Mg(OH)₂, MgSO₄×H₂O, MgSO₄×2H₂O, MgSO₄×4H₂O, MgSO₄×5H₂O, MgSO₄×6H₂O, MgSO₄×7H₂O, MgSiO₃, Mg₃Si₄O₁₀(OH)₂ (talc), Mg₇Si₈O₁₆H₂ (anthophyllite), Mg₃Si₂O₅(OH)₄ (serpentine), Mg₅Al₂Si₃O₁₀(OH)₈ (chlorite), Mg₄Si₆O₂₁H₁₂ (sepiolite), NaAlSi₃O₈, NaCl, NaAl₂(AlSi₃O₁₀)(OH)₂ (paragonite), Ca₂FeAl₂Si₃-O₁₂OH (epidote), Ca₂Mg₅Si₈O₁₆H (tremolite), CaAl₂Si₃O₁₀-(OH)₂ (prehnite), Mg₇Si₈O₂₂(OH)₂ (cummingtonite), Mg₃Al₂Si₃O₁₂ (pyrope), Fe₃Al₂Si₃O₁₂ (almandine), Ca₃Al₂Si₃O₁₂ (grossular), Ca₃Fe₂Si₃O₁₂ (andradite), S, and SiO₂ (quartz).

We also considered the following aqueous species in order to model the complex chemistry of iron in aqueous solution, mineral dissolution, and acid-base chemistry that may result from oxidation of endogenic compounds: Al³⁺, CH₄, CO, CO₂, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm CO}_3^{2 - }$$ \end{document} , Ca²⁺, Cl⁻, Fe²⁺, Fe³⁺, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm Fe} ( { \rm OH} ) _2^ +$$ \end{document} , Fe(OH)₂, Fe(OH)₃, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm Fe} ( { \rm OH} ) _3^ -$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm Fe} ( { \rm OH} ) _4^ -$$ \end{document} , Fe₂(CO₃)₃, FeCO₃, FeO⁺, FeO, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm FeO}_2^ -$$ \end{document} , FeOH²⁺, FeOH⁺, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm FeSO}_4^ +$$ \end{document} , H⁺, H₂, H₂CO₃, H₂O, H₂O₂, H₂S, H₂SO₃, H₂SO₄, H₄SiO₄, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HCO}_3^ -$$ \end{document} , HFeO₂, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HFeO}_2^ -$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HO}_2^ -$$ \end{document} , HS⁻, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HS}_2^ -$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_3^ -$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_4^ -$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_5^ -$$ \end{document} , K⁺, Mg²⁺, Na⁺, O⁻, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm O}_2^{\;\;2 - }$$ \end{document} , O₂, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm O}_2^ -$$ \end{document} , O₃, OH, OH⁻, S²⁻, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm S}_2^{\;\;2 - }$$ \end{document} , SO₂, \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_3^{\;\;2 - }$$ \end{document} , SO₃, and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{\;\;2 - }$$ \end{document} .

4. Results

4.1. Cases without direct interaction with rock

In the absence of reactive rock, if the quantity of oxidants delivered to Europa's ocean becomes stoichiometrically equivalent to or exceeds the quantity of reductants, then the ocean becomes strongly acidic. Given the rate of delivery of reductants assumed in our calculations [based on McCollom (1999)], the early europan ocean would have been chemically reduced. Only when the oxidants begin to flow after the 2-billion-year delay [based on the timing and flux estimate of Greenberg (2010)] would the ocean begin to become oxidized. We found that the pH of the ocean drops below 3 in only 30 million years once the oxidants start to arrive. The pH then decreases slowly to 2.6 after another 2.5 billion years (i.e., at 4.5 Gyr from the start).

As shown in Fig. 1, similar low pH values would be reached even with much lower oxidant fluxes. In all these cases, most of the drop in pH occurred shortly after the delivery of oxidants begins (Fig. 2). The subsequent decrease was slow, because no sulfide remained after the initial drop; it was all consumed to form sulfate and acid. Oxidants would still be left over despite reaction with the reductants, so the ocean would remain strongly oxidizing with free O₂ available for other reactions, including metabolic processes. As long as the quantity of oxidants in this scenario is greater than the quantity of reductants, the final pH is fairly independent of how much greater it is; the pH will always drop to near 2.6.

FIG. 1.

pH as a function of O₂ flux and H₂O₂ flux, after the flux of reductants has continued for 4.5 billion years. These values are fairly constant from shortly after the oxidants begin to enter the ocean, which is taken for these calculations to commence 2 billion years after the reductant flux begins. Except for very small amounts of both H₂O₂ and O₂, pH values are generally significantly less than 3. For the maximum flux rates from Greenberg (2010), shown by the black dot, the resulting pH is 2.6.

FIG. 2.

The change in pH of Europa's ocean over time. The pH drops quickly once oxidant delivery begins at 2 Gyr. Values at 4.5 Gyr are also shown in Fig. 1. Three cases are shown here: 100%, 10%, and 1% of the flux estimated by Greenberg (2010). In the latter case (dashed line), the changes in slope after the commencement of oxygenation are a result of buffering by carbonate and other species.

Because the low pH results from the complete oxidation of sulfide, changes to this value are possible only with changes to the fluxes of reductants. The key reductant that leads to the acidification of Europa's ocean is sulfide as H₂S. If the flux of H₂S has changed significantly over time, or if the flux is different from the estimates by McCollom (1999) and by Hand et al. (2007), then the resulting pH of Europa's ocean could be affected. For example, an increase in the flux of H₂S by a factor of 100 to 10¹¹ mol/yr yields a pH as low as 1.1. Changing the CH₄ flux would also change the resulting pH slightly as a result of formation of H₂CO₃. However, H₂S oxidation is the most important reaction in acidification, unless the total H₂S flux is low, in which case the oxidation of CH₄ can be important. The dependence of the final pH on the H₂S flux is shown in Fig. 3.

FIG. 3.

The relationship between the flux of H₂S and the resulting pH. Here, the flux of oxidants is set equal to Greenberg's (2010) maximum estimates, but the result would be similar for slower oxidant delivery as long as the rate were adequate to consume the endogenic reductants.

In all the cases described above, iron speciates as hematite (Fe₂O₃); carbon is a mixture of dissolved CO₂ gas and H₂CO₃; sulfur is principally in sulfate (10:1 \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{\;\;2 - } :{ \rm HSO}_4^ -$$ \end{document} ); and Mg, Ca, K, and Na are dissolved ions.

4.2. Reaction with rock

The pH values resulting for the models that include rock interactions are summarized in Table 1. If the volume of interacting rock is equivalent to the upper 1 m, the rock has little effect on acidification, independent of composition or temperature; the pH becomes 2.6, just as for the case with no rock. However, if rock equivalent to a thickness of a kilometer or more is involved, then dissolution of olivine can neutralize the acid: \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm Mg}_2{ \rm SiO}_4 + 4{ \rm H}^ + \rightarrow 2{ \rm Mg}^{2 + } + { \rm H}_4{ \rm SiO}_4 \tag{3}\end{align*} \end{document}

Table 1.

Values of Oceanic pH for 12 Cases with Reactive Rock: Two Rock Types, Two Temperatures, and Three Assumed Volumes of Reactive Rock

		Depth of reactive rock layer
Composition	Temperature (°C)	1 m	1 km	25 km
Ultramafic	0	2.6	9.96	10
	250	2.6	4.8	5
Basaltic	0	2.6	3.1	4.6
	250	2.6	3.4	3.4

Quantities of reductants are based on 4.5 billion years of delivery at a constant rate; quantities of oxidants are based on 2.5 billion years of delivery at a constant rate (after the 2-billion-year delay).

For example, with ultramafic rock composition and a low temperature (0°C), the pH can become very high if kilometers of rock participate, although the effect is less pronounced at higher temperatures. If the rock is basaltic, then the degree of buffering of acid is decreased because basaltic rock is richer in Al and Si, which have poorer potential for neutralization of H₂SO₄ than Mg, Na, K, and Ca oxides. Moreover, in the basalt model the ocean is already close to saturation in Mg and Ca, which further limits neutralization by preventing Reaction 3 from reaching completion.

In these calculations, basalt rock alters to prehnite, sepiolite, and paragonite, whereas ultramafic rocks alter to prehnite and brucite. All iron is present as hematite in both systems. No sulfate minerals are produced in these calculations.

5. Discussion

If the flux of oxidants from the crust exceeds the flux of reductants from the interior (as seems plausible), the pH of Europa's subsurface ocean is generally decreased significantly. The decrease in pH is primarily due to the formation of acid during the oxidation of sulfide and is directly related to the flux of sulfide in the oceans. The acidification would have occurred soon after the oxidants began to enter the ocean, even if the sulfide (and other endogenic reductants) had been accumulating for 2 billion years beforehand. Even with an oxidant delivery rate much lower than that estimated by Greenberg (2010), the pH would have decreased to near its low steady-state value in less than a 100 million years (Fig. 2). An exception to these conclusions would be if the ocean were in chemical equilibrium with a large volume of ultramafic rock, which could neutralize the acid. We next discuss why that condition seems unlikely to be relevant for Europa. We then discuss the implications of our model as to the chemical composition of the ocean and for the possibility of life there.

5.1. Acid neutralization by rock

Neutralization of the acidic europan ocean by rock would primarily occur through dissolution of silicates to release Na, K, Ca, and Mg. Neutralization of the ocean will occur only if the rock at the bottom can provide these elements in sufficient quantity. Kargel et al. (2000) proposed a scenario in which the ocean bottom consists of non-reactive sulfates; hence, this system would not neutralize acid produced by reaction of sulfides with oxidants (though whether such chemistry would arise in this ocean is another question). Our results indicate that basalts are similarly ineffective, unless a huge volume, representing a layer tens of kilometers thick, is intimately reactive with the ocean. The basalts are relatively ineffective because they contain substantial amounts of SiO₂, which does not neutralize acid. Ultramafic rocks could be more effective, but the rocks must be reactive to a depth of ∼1 km, and temperatures should be less than 200°C to keep neutralization efficient.

Although ocean water may penetrate kilometers deep into the rock layer (Vance and Goodman, 2009), it is unlikely that more than a small fraction of the rock in that layer is able to interact with the ocean. For a substantial portion of the porous layer of rock to be chemically exposed to seawater, the rock would need to be porous on a very fine scale. For example, if the porous layer extended 25 km into the rock, it is unlikely that more than 1 in 25,000 atoms in that layer (equivalent to our 1-meter-thick-rock case, which was inadequate for neutralization) would be exposed to the ocean intimately enough to be in chemical equilibrium. In principle, more atoms could be exposed, but only with very small capillaries to carry the water, in which case the flow could hardly be adequate to keep the rock in equilibrium with the ocean. Moreover, it seems unlikely that the rock involved would be predominantly ultramafic, which (according to Table 1) further reduces the likelihood that the ocean has been neutralized.

Even so, it is conceivable that enough rock could be engaged to neutralize some ocean water just above the rock surface. More likely, vertical transport would keep the ocean water mixed by convection or thermal plumes over hot spots, in which case the neutralizing effect of the rock would be diluted in the large volume of the ocean. As we have shown, the entire ocean could only be neutralized if a huge volume of the rock were in intimate chemical equilibrium with the ocean.

Thus, an important issue is the rate and extent of vertical mixing. With vertical mixing, it is unlikely that enough rock can participate to neutralize the acidic ocean water. Such neutralization would require interaction with a huge amount of rock material, even if it were ultramafic. On the other hand, if vertical mixing is slow enough, or if it does not extend all the way up to the ice or down to the rock, the upper ocean would be acidic (as calculated in Section 4.1) at the same time as the deepest part has a higher pH (as calculated in Section 4.2).

The transport timescale due to chemical diffusion of components through the depth of the ocean is ∼Z ²/k, where Z is the depth of the europan ocean (from the water-ice to the water rock boundary) and k is the vertical diffusion coefficient. For a depth of 100 km, and a vertical diffusion coefficient of 10⁻⁵ m²/s, set equivalent to Earth's (Ledwell et al., 1993), the mixing timescale is hence of the order of ∼10⁷ years. This timescale suggests that chemical diffusion alone would be slow enough to allow a significant difference in pH between the upper ocean and the bottom of the ocean. However, vertical transport of components by heat plumes or convection in the ocean may mix the ocean on shorter timescales. Acid neutralization by rock will only be significant if the lower portion of the ocean is not involved in such mixing, so that the neutralizing species are not overly diluted or, if ultramafic rock is involved, in very large quantities.

5.2. Implications for ocean composition

In the “primordial” jovian nebular region, sulfur likely accreted into the jovian moons primarily as FeS and H₂S (Pasek et al., 2005). Sulfates have been observed on Europa's surface and are attributed to radiolysis that involved substances delivered from Io (Carlson et al., 2002, 2005, 2009) and to direct delivery from the subsurface ocean (Zolotov and Shock, 2001; McCord et al., 2010). If the sulfates have an oceanic origin, they may have been a leachate of the chondritic accretionary material of Europa (Kargel et al., 2000; Fanale et al., 2001) or, more likely, products of reactions of sulfides with oxidants (e.g., McKinnon and Zolensky, 2003).

Estimates of the oceanic composition of Europa have been based on extracts of carbonaceous chondritic meteorites. These estimates suggest a salty ocean initially comprising oxidized components, including sulfate, which is a major anion in carbonaceous chondrite extracts (Fanale et al., 2001). However, the sulfate extracted from carbonaceous chondrites may have been formed by weathering on Earth's surface in the years following their fall. Sulfate veins bear O isotope characteristics indicative of terrestrial oxygen, as does some of the water-soluble sulfate (Gounelle and Zolensky, 2001; Airieau et al., 2005). Sulfates are not found in more “pristine” chondrites, such as those of Tagish Lake (McKinnon and Zolensky, 2003). In most meteorites, sulfur is in sulfide minerals such as troilite, FeS, and pyrrhotite, Fe_1-xS. Additionally, sulfur chemistry in the nebula near the location of Jupiter accretion should have been dominated by reducing species such as H₂S (Grossman, 1972; Pasek et al., 2005). This suggests that sulfur around Jupiter was initially in a reduced state and present as sulfides. Moreover, sulfides are the dominant constituent from “black smoker” deep sea vents on Earth (Von Damm, 1990) and are predicted to dominate hydrothermal vent fluids on Europa (McCollom, 1999).

Our analysis suggests that sulfate on Europa may be an oxidation product of sulfide in Europa's subsurface ocean, which is in agreement with Hand et al. (2007). So long as the flux of oxidants is adequate, as seems plausible, then all sulfide will be consumed rapidly (on the timescale of mixing in the subsurface ocean) to form sulfate. If a flux of 10⁹ mol/yr of endogenic sulfide enters Europa's ocean (as estimated by McCollom, 1999 and Hand et al., 2007), then, over the age of Europa, approximately 5×10¹⁸ mol of sulfate will have formed from this reaction. Consequently, approximately 10¹⁹ mol of H⁺ would also form during this oxidation process, yielding the pH of the europan ocean of ∼2.6. At a pH of 2.6, the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{\;\;2 - }$$ \end{document} to \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_4^ -$$ \end{document} ratio is approximately 10:1 (or less with higher H₂S fluxes); hence, sulfate salts from the subsurface ocean should be a mixture of compounds with \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{\;\;2 - }$$ \end{document} and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_4^ -$$ \end{document} anions. If these salts reach the surface of Europa, for example through cracks or melt-through sites in the ice crust, then the radiolysis of a \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_4^ -$$ \end{document} salt to H₂SO₄ would be expected (Carlson et al., 2005). The best fits to Europa's hydrate appear to be mixtures of Na sulfates, Mg sulfates, and H sulfates (Carlson et al., 2005; Orlando et al., 2005), though a precise identification is not presently possible. Both sulfuric acid hydrate and sulfate salts are plausible products on Europa's surface via reaction of O₂ with sulfide in the subsurface ocean, even without a contribution from Io.

The possibility of an acidic ocean was previously raised by Kargel et al. (2000) as one of many possible scenarios for Europa's subsurface ocean system. The acidification in this system was attributed to reaction of SO₂—likely delivered by Io to Europa's surface—with water to form H₂SO₄, and required a non-reacting rock wall surface. Other scenarios suggest little acidity for Europa's subsurface ocean, especially if the oxidant flux to the ocean is limited, and the sub-oceanic rock has a composition equivalent to carbonaceous chondrite material (e.g., Zolotov and Kargel, 2009). Many models of Europa's subsurface ocean predict saturation with MgSO₄ (McCord et al., 1999; Kargel et al., 2000; Hand and Chyba, 2007), based in part on carbonaceous chondrite extracts. Hydrothermal reactions of acid in the ocean with rock may also dissolve Mg silicates, which could be placed on the surface as Mg-sulfates as Mg- and sulfate-rich water freezes. However, our model does not predict saturation of an MgSO₄ mineral in the ocean, as there is too little sulfur present from hydrothermal sources to saturate the ocean. If the amount of sulfur is changed substantially, as indicated in Fig. 2, then the water may become increasingly concentrated in sulfate, though saturation requires about 10³ times the current estimated flux of sulfur.

5.3. Conditions for life

The implications of an oxygenated ocean have been discussed by Hand et al. (2007) and by Greenberg (2010). The rates of delivery of oxidants into the ocean could be adequate to support significant biota, with concentrations adequate to support large eukaryotes, even macrofauna. However, as shown here, the oxygenation of the ocean would likely be accompanied by acidification. Eukaryotes, especially macrofauna reliant on biomineralization for structure, fare poorly in acidic water. Hence, if the acidification occurred, life in Europa's subsurface ocean would likely consist primarily of organisms analogous to one-celled bacteria- or archaea-like cells, unless more complex life-forms evolved specifically for this environment.

If the flux of oxidants to Europa's subsurface ocean was delayed by 1–3 billion years in accord with the analysis by Greenberg (2010), then prebiotic chemistry may have had ample time to generate organisms with genetic systems and protective structures (e.g., cells). An active anaerobic ecosystem could have developed prior to an increase in oxidant flux. Then, as the oxidant flux commenced, organisms would have had structures in place to protect them from damage due to oxidation, which would have given them a chance to evolve high-efficiency oxygen-consuming metabolisms (as happened on Earth as oxygen built up in the atmosphere and ocean). However, as the acidification process discussed here proceeded, the pH could have become so low that the continued existence of life, especially of any macroorganisms, would be problematic.

On the other hand, if organisms developed oxygen- or acid-consuming metabolisms or chemically changed the environment more quickly than the acidification occurred, then they might have minimized the deleterious effects. For example, just as terrestrial microorganisms can cause mineral weathering (Uroz et al., 2009), a similar effect by europan organisms could enhance acid neutralization by the seafloor rock. Also, if organisms could have metabolized the oxygen quickly enough as it entered the ocean, they might have delayed or prevented the acidification, but only if enough oxygen atoms remained sequestered within the bodies or remains of the organisms.

The ecosystem would need to evolve quickly to meet this challenge. According to Fig. 2, even if the oxidant flux were 10% of Greenberg's estimate, oxygen metabolisms and acid tolerance would need to evolve in ∼50 million years. It may be that at first only organisms close to the ocean bottom would be sufficiently protected from the acidic, oxic environment by the neutralizing effect of the rock. The ecosystem might then gradually evolve so that organisms could exploit the oxic environment and stop the acidification.

5.4. Biominerals

Formation of biominerals may be an important step in the diversification of macrocellular life (e.g., Ward and Brownlee, 2000). The carbonate minerals that form shells of mollusks and other invertebrates on Earth are highly soluble in water with a pH of 2.6 [up to concentrations in excess of 1 mol/L water (hereafter M) of Ca²⁺ or carbonate]. Similarly, calcium phosphate minerals, such as those that form vertebrate skeletons, are also highly soluble (up to concentrations in excess of 0.1 M of phosphate). However, if acidification proceeded in the europan ocean, then production of two of the dominant biominerals (apatite and calcite/aragonite) used by modern terrestrial life would not be feasible. Alternative biominerals are possible. For example, silica is an alternative used principally by terrestrial algal forms, and siliceous biominerals are not affected by acidic solutions. However, they can only form in water saturated in Si (approximately 1 mM concentration of Si as silicic acid). Given the importance of phosphorus to life (e.g., Pasek, 2008), an alternative to calcium phosphate biominerals might be iron phosphate biominerals, as iron phosphate is poorly soluble in acidic solutions (phosphate solubility of less than 10⁻⁹ M); hence, it is significantly more stable than calcium phosphates under these conditions.

6. Conclusion

Acidification of Europa's ocean likely commenced as soon as oxidants began to be delivered to the ocean, especially in the upper regions of the ocean. This effect may have been prevented if organisms evolved a process to utilize and consume oxidants and the ecosystem and rock-buffering modified the acidity over a relatively short timescale. Otherwise, the oxidant flux could have quickly overwhelmed the accumulated reductants and generated significant acid during the process. Life at the bottom of the ocean would have had a better chance to be protected from this acidification than life higher up in the ocean, closer to the water-ice boundary.

Unless europan ocean life evolved an effective biomineralization process and the ability to consume and sequester arriving oxidants, any organisms would have to tolerate a high quantity of oxidants and low pH. In that case, a surviving ecosystem in Europa's ocean might be analogous to microbial communities that live in acid mine drainage (e.g., Edwards et al., 2000). Thus, studies of organisms at Río Tinto may be especially pertinent to putative europan life (e.g., Fernandez-Remolar et al., 2005; Amils et al., 2007; Davila et al., 2008). These communities are dominated by acidophiles that oxidize iron and sulfide as sources of metabolic energy. Indeed, in these environments, Fe³⁺ can be a more efficient oxidant than even O₂; hence, these microbes may promote the active acidification of their own environment. If the dominant reductant in Europa's ocean is indeed sulfide, then such may have been the fate of Europa's ecosystem as well. It appears that only if life were able to evolve quickly enough as the oxidants arrived in the ocean could it have prevented development of such a hostile environment or come to live with it.

Footnotes

Acknowledgments

The authors thank Virginia Pasek for helpful edits and Tom McCollom for helpful discussion. Insightful reviews from Chris McKay and other, anonymous reviewers were especially helpful. M.A.P. and R.G. were supported by grants from NASA's Exobiology and Evolutionary Biology program (NNX10AT30G) and Planetary Protection program, respectively.

Disclosure Statement

The authors declare no competing interests.

1

The HSC chemical code stands for enthalpy, entropy, and heat capacity. More details on the code can be found at www.hsc-chemistry.net.

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