Effects of H 2 O, ph, and oxidation state on the stability of Fe minerals on Mars

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1 JOURNAL OF GEOPHYSICAL RESEARCH, VOL. 110,, doi: /2005je002482, 2005 Effects of H 2 O, ph, and oxidation state on the stability of Fe minerals on Mars P. L. King 1 and H. Y. McSween Jr. Department of Earth and Planetary Sciences, University of Tennessee, Knoxville, Tennessee, USA Received 2 May 2005; revised 7 August 2005; accepted 31 August 2005; published 21 December [1] On Mars, there is evidence that solutions may have been present on the surface episodically in the past. These solutions were derived by weathering minerals such as olivine and sulfides found in mafic-ultramafic rocks, and possibly sulfates and meteorite fragments. Upon removal of water and/or supersaturation of solutions, secondary minerals formed. We present a theoretical model for the formation of the solutions and the subsequent precipitation of Fe-bearing phases. The first Fe-bearing phases to weather via oxidation are sulfides (troilite or pyrrhotite) which produce secondary Fe (hydr)oxides or FeS 2 (pyrite or marcasite) and Fe 2+ sulfates such as melanterite. Melanterite may be replaced by Fe sulfates that have decreasing H 2 O/(3Fe Fe 2+ ), increasing oxidation, and increasing bulk OH/(OH + 2SO 4 ). At Meridiani Planum the presence of jarosite indicates that the solutions were oxidized with ph < 4.5. The solutions were likely Fe-Mg-(Ca)-SO 4 -(Cl)-rich and precipitated Fe (hydr)oxides, Fe phosphates, Fe sulfates with low OH/(OH + SO 4 ), Ca-Mg sulfates, and possible halides, along with Si-rich phases. The minerals indicating an acid environment were preserved by either removing components such as salts or water from the system (e.g., in dust storms or episodic aqueous events); isolating the mineral surfaces from fluids; using disequilibrium processes; or a combination. At other localities, jarosite is below detection limits, and schwertmannite may be present. If schwertmannite is present, the solution had nearneutral ph (4 < ph < 10) and Fe (hydr)oxides would have crystallized rapidly, leaving Mg-Na-(Ca)-SO 4 -Cl-rich solutions that likely precipitated Ca phosphates, Ca-Mg-Na sulfates, Fe sulfates with moderate-high OH/(OH + SO 4 ), halides, Si-rich phases, and possibly Fe-Mg-Ca carbonates. Citation: King, P. L., and H. Y. McSween Jr. (2005), Effects of H 2 O, ph, and oxidation state on the stability of Fe minerals on Mars, J. Geophys. Res., 110,, doi: /2005je Introduction [2] Currently, Mars is a relatively dry planet: water occurs in the atmosphere and ice caps, and is stored in the near surface or subsurface in solids and as transient films along grain boundaries [e.g., Arvidson et al., 2005; Bishop et al., 2002; Carr, 1996; Christensen, 2003; Feldman et al., 2002]. In contrast, in the past, aqueous solutions existed episodically on the Martian surface [e.g., Baker, 2001]; however, the age, lifetime, volume, formation mechanism, and composition of these solutions are ambiguous [Christensen, 2003; Jakosky and Phillips, 2001; Lee and McKay, 2003; Phillips et al., 2001]. To better resolve these ambiguities, it is necessary to determine the geologic and chemical conditions that formed Martian solutions. Did solutions form under acid fog or near-neutral ph leaching conditions? Were solutions oxidized? Can we constrain the quantity of solution or water present? Do the same solution processes operate at all 1 Visiting from Department of Earth Sciences, University of Western Ontario, London, Ontario, Canada. Copyright 2005 by the American Geophysical Union /05/2005JE locations on Mars? These questions can be addressed by examining secondary minerals on the Martian surface, particularly Fe-O-H-S minerals because their stability is sensitive to the chemistry of the precursor bulk solution, ph, oxidation state, and water abundance. [3] Relative to Earth, Mars has undergone limited weathering, as indicated by geochemical trends [Bell et al., 2000; McLennan, 2000; McSween and Keil, 2000] and the occurrence of olivine on the surface [Hoefen et al., 2003; Mustard et al., 2005]. Recent workers concur that the bulk solution compositions on Mars result from interaction of solutions with primary mafic-ultramafic crust [e.g., Catling, 1999; King et al., 2004; Tosca et al., 2004]. Thus the Martian surface can be treated as a mixture of primary mafic-ultramafic rocks and secondary minerals [e.g., Banin et al., 1992; Clark, 1993; Catling, 1999; King et al., 2004; McLennan, 2000; McSween and Harvey, 1998; McSween and Keil, 2000; Tosca et al., 2004]. The secondary minerals contain a record of the solutes contained in any preexisting solutions because they precipitate when they are saturated in the solution. The secondary minerals also contain a record of oxidation of mantle-derived minerals in the Martian atmosphere where the partial pressure of oxygen (PO 2 )is relatively higher. 1of15

2 [4] Secondary minerals are found in Martian dust and soil, and form cements, veins and coatings on rocks and in the Martian meteorites [e.g., Arvidson et al., 2005; Bridges et al., 2001; Squyres et al., 2004a, 2004b, 2004c]. The soils and secondary minerals are enriched in Mg, Na, S and Cl, with a strong correlation between Mg and S, which can be explained by concentrations of sulfate and halide salts [e.g., Clark and Van Hart, 1981]. Data from the recent Mars missions show that the secondary minerals include nanophase Fe 3+ -O-H phases, Ca-Mg-Fe sulfates and halides, Si-rich phase(s), and H-bearing phase(s) [Arvidson et al., 2005; Bell et al., 2004; Bibring et al., 2005; Christensen et al., 2004; Gendrin et al., 2005; Klingelhöfer et al., 2004; Langevin et al., 2005; Morris et al., 2004; Squyres et al., 2004a, 2004b, 2004c]. There is some evidence for weathering to clays, but the extent is unknown [Hamilton et al., 2003; McSween and Keil, 2000; Michalski et al., 2005; Wyatt and McSween, 2002]. In this paper, we concentrate on the Fe and S history of the Martian solutions because other work summarizes the acquisition and fate of the other solutes [e.g., King et al., 2004; Tosca et al., 2004]. [5] Three major models have been proposed to form solutions on Mars and involve altering mafic-ultramafic rocks with (1) hydrothermal fluids [Newsom, 1980; Newsom et al., 1999]; (2) low-temperature solutions [Bishop and Murad, 1996; Bishop et al., 2002; Burns, 1987, 1988, 1993; Burns and Fisher, 1990; Catling, 1999; McSween and Harvey, 1998; Warren, 1998]; or (3) magmatic acid fog [e.g., Banin et al., 1997; Clark and Van Hart, 1981; Morris et al., 1996; Settle, 1979; Tosca et al., 2004]. We refer to the first two alteration types as leaching models where solutions at a range of temperatures and volumes leach and dissolve the Martian surface and/or meteorite fragments to produce Mg-Na-Ca-SO 4 -Cl-rich solutions [King et al., 2004]. In contrast, in the acid fog model, the solutes are derived from a magmatic S-HCl vapor that interacts with H 2 O to form an acidic solution (ph < 2). The solution alters primary minerals with the highest relative dissolution rate to produce Mg-Fe-(Ca)-SO 4 -Cl-rich solutions [Hurowitz et al., 2005; Tosca et al., 2004, 2005]. Early versions of the acid fog model allowed the solutions to evolve to high ph conditions [e.g., Settle, 1979]. In subsequent discussion, the acid fog model refers to the current version of the model where ph remains low [Hurowitz et al., 2005; Tosca et al., 2004, 2005]; although those workers allow for some increases in ph, they emphasize that low ph (ph < 4) conditions are likely maintained. [6] There are significant differences between the leaching model and the acid fog model, even though the acid fog model also includes a component of leaching. First, in the leaching model the initial fluid/solution phase contains both cations and anions derived from silicate-sulfide rocks and crustal fluids, whereas in the acid fog model, the initial vapor phase contains only dissolved species derived from magmatic gases: HCl and SO 2, or perhaps H 2 S[Zolotov, 2003]. Second, the leaching model does not require high temperatures (although they are not ruled out) and as a consequence leaching can occur in a range of settings. In contrast, the acid fog model requires high temperatures adjacent to shallow level magma chambers or volcanoes. Third, in the leaching model the initial ph is unconstrained, but it is assumed that the solution ph may be raised via reactions like ðmg; FeÞ 2 SiO 4 olivine þ 4H þ aq ðr1þ ¼ 2Mg; ð Fe Þ2þ aq þ2h 2O þ SiO 2aq : At near-neutral ph, Fe 2+ is oxidized and Fe 3+ phases precipitate at high rates even at relatively low PO 2 depleting the solution in Fe [Burns, 1993; Marion et al., 2003]. Such solutions will become Mg-Na-Ca-SO 4 -Cl-rich [King et al., 2004, section 3.4] due to precipitation of Febearing phases, dissolution of phases like olivine, pyroxene and basaltic glass, and leaching of minerals such as feldspars. On the contrary, in the acid fog model as it is currently applied, the ph stays low so that Fe and Al remain in solution [Hurowitz et al., 2005]. The solutions are proposed to remain acidic because they have a high buffering capacity and reaction with basic materials (e.g., + silicate minerals) does not consume sufficient H aq to increase ph substantially [e.g., Tosca et al., 2005]. Such acidic solutions tend to result in higher weathering rates, producing solutions that reflect the composition of the mineral with the highest dissolution rate [Hurowitz et al., 2005]. For instance, the soils from the Viking, Mars Pathfinder and MER Gusev Plains rocks all indicate that their chemistry was largely derived from olivine, suggesting Mg-Fe-(Ca)-SO 4 -Cl-rich solutions [Hurowitz et al., 2005; Tosca et al., 2004, 2005]. Fourth, the leaching model assumes that the water:rock ratio is sufficiently high for ph to be raised (e.g., reaction (1)). In the acid fog model, it is necessary to invoke high water:rock ratios (>100) at low ph followed by low water:rock ratios (<10) at higher ph to preserve the mineral assemblages [Elwood Madden et al., 2004]. Alternately, in the low ph acid fog model it is necessary to maintain low water:rock ratios (10) [Hurowitz et al., 2005], or episodic events with varying water:rock (porosity) and solution compositions [Tosca et al., 2005]. We do not consider the water:rock ratio further because it has been modeled extensively by previous authors. 2. Background 2.1. Sources of Fe [7] The primary Fe-bearing phases on Mars have been evaluated previously to determine the sources of Fe and S in solution [e.g., O Connor, 1968; Gooding, 1978; Burns and Fisher, 1990; Burns, 1993; Tosca et al., 2004]. In general, Fe-bearing primary phases weather in the following sequence: oxides < silicates including olivine, pyroxene and glass < sulfides < natural metallic materials [e.g., Krauskopf and Bird, 1995], where the oxides in particular weather most readily at low ph [Tosca et al., 2004]. We recognize that this sequence is a generalization, and refer the reader to Burns [1993] for a discussion of the effects of mineral composition and grain size, temperature, Eh, ph, and ionic strength of the solution; those variables are not considered further. [8] Dissolution of multiple phases likely contributes cations and anions to the solution based on the bulk compositions of the Martian soils (soils = rock + oxides + salts) and analog experiments [e.g., Tosca et al., 2004]. 2of15

3 However, olivine is proposed to be the dominant mineral controlling the fluid composition on Mars [Tosca et al., 2004; Hurowitz et al., 2005]; thus we examine it further. [9] The dissolution rate of olivine is on the order of 10 8 to mol.m 2.s 1 and depends on ph, with rates lowest at near-neutral ph and increasing as ph is lowered or raised under atmospheric conditions (PCO 2 < atm [Wogelius and Walther, 1991]). At ph > 6, olivine dissolution rates are inversely correlated with PCO 2 ;for instance, at ph = 12 and PCO 2 = atm, dissolution rates decrease about an order of magnitude [Wogelius and Walther, 1991]. This behavior could be due to CO 3 2 surface complexes [Wogelius and Walther, 1991], but infrared CO 3 2 bands were measured on fresh olivine and attributed to fluid inclusions [Pokrovsky and Schott, 2000]. Because observations of the fluid inclusions were not reported, future research on the response of olivine dissolution to changing PCO 2 (and similarly PO 2 [see Burns, 1993]) will aid our interpretations of weathering on Mars. [10] In addition to the silicates and oxides, metals from meteorites could provide some Fe to Martian solutions because there could be at least g.m 2 meteoritic material on the Martian surface [Bland and Smith, 2000]. Finally, sulfides from both Martian rocks and meteorites may provide a readily available primary mineral source for both Fe and S species in solution [Burns, 1993; Marion et al., 2003] Sources of S [11] Sulfides present in the Martian meteorites include pyrrhotite, rare chalcopyrite, pentlandite, and troilite-pentlandite-chalcopyrite intergrowths with some alteration to marcasite [McSween and Treiman, 1998; Rochette et al., 2001]. Pyrite is not abundant on Mars, but it is a good analog for the other sulfides and many alteration models for Mars have been based on pyrite alteration (section 4.2). Even at negligible water contents and low PO 2 (like Mars), pyrite surfaces form Fe 2+ -O surface species or nanophase Fe species [Nesbitt et al., 1998]. Pyrite s dissolution rate in solution is to mol.m 2.s 1 at ph 6 7 in Earth s atmosphere at standard conditions [Moses and Herman, 1991]. On Mars, the absolute rates of all reactions are likely lower due to lower temperatures, lower PO 2, and higher ionic strength of the solutions [Burns, 1993], but sulfides still have higher dissolution rates than silicates and oxides. [12] Sulfur could also originate from a volcanic or hydrothermal fluid/gas phase. The S in such a phase may be derived from degassing of a basaltic melt (S in melt! Sin fluid/gas), but this is unlikely to provide large quantities of S because the maximum S solubility in silicate melts is relatively low, even at high pressures [e.g., O Neill and Mavrogenes, 2002; Wallace and Carmichael, 1992]. For this reason, the S in a volcanic or hydrothermal fluid/gas is more likely to be derived from a fluid/gas interacting with sulfides or immiscible sulfide liquids. Thus the processes that produce S-rich solutions could be the same in both the acid fog and leaching models; however, as we note above, other features distinguish the two processes. [13] In evaluating the S source, we can examine the extreme case where S was derived from sulfides in a closed system. The observed SO 3 ( wt.% absolute) in the soils [Clark et al., 1982; Foley et al., 2003; Gellert et al., 2004; Rieder et al., 2004] can be derived from vol.% FeS in the rock; however, the volume of sulfides is much lower in the Martian meteorites [McSween and Treiman, 1998]. This mass balance problem may be resolved in several ways. First, it is more likely that any solution percolates along grain boundaries in an open system and therefore a greater volume of sulfides can contribute to S enrichments at the Martian surface. Second, it is likely that some magmatic sulfur, not derived from altering sulfides, is incorporated in solutions. Third, it is possible that the intrinsic S content of Mars crust is high [Clark and Baird, 1979] relative to the volume of sulfides in the Martian meteorites. For example, S could be stored in secondary sulfates like those observed in the Martian meteorites [Bridges et al., 2001] (see section 6.1). In summary, primary sulfides, even at relatively low abundances, could provide significant amounts of the long term S budget to Martian solutions, but other magmatic and secondary sulfate sources are also possible Secondary Fe (Hydr)oxides [14] On Mars, secondary Fe-bearing minerals include poorly characterized Fe (hydr)oxides where we use that term to include Fe hydroxides, Fe oxides, Fe oxyhydroxides or a combination [Bell et al., 2000, 2004; Klingelhöfer et al., 2004; Morris et al., 2004]. Many of these minerals are likely nanophase and may include ferrihydrite, hematite, magnetite, maghemite, goethite, akaganéite, feroxyhyte, and lepidocrocite (Table 1). The large number of predicted phases reflects the difficulty of characterizing nanophase Fe (hydr)oxides [Bell, 1996; Dyar et al., 2005a; Dyar and Schaefer, 2003; Morris et al., 2004]. Furthermore, the Fe (hydr)oxides tend to be converted to one another by dehydration (loss of H 2 O), dehydroxylation (loss of OH), thermal transformation, and/or oxidation [Cornell and Schwertmann, 1996]. Amorphous, metastable Fe (hydr)oxide phases are kinetically favored at low temperatures [Russell, 1979] like those on Mars, but thermodynamic models are easier to apply to crystalline phases. [15] In this paper, we take the simplifying assumption that the Fe (hydr)oxides may be treated as a group [e.g., Marion et al., 2003] because their thermodynamic properties grade into one another, particularly if crystallinity and grain size effects are taken into account [Burns, 1993; Cornell and Schwertmann, 1996]. These minerals generally age to hematite because it is the most thermodynamically stable phase in the Martian environment [Gooding, 1978; Morris et al., 1989], although goethite may be present in the polar regions or in the subsurface [Zolotov et al., 1997] Secondary Fe Salts [16] Iron sulfates are also present on Mars [Klingelhöfer et al., 2004; Morris et al., 2004] and are also commonly fine-grained, poorly crystalline phases with compositions that are difficult to characterize [Dyar et al., 2005a; Dyar and Schaefer, 2003; Lane et al., 2004, 2005]. Therefore a large set of Fe sulfates are possible on Mars with early suggestions including botryogen (Mg-bearing hohmannite), römerite, copiapite, and jarosite group minerals (e.g., hydronium-jarosite) [Burns, 1987] (see Table 1 for mineral 3of15

4 formulae). Additional Fe sulfates have been inferred using theoretical, experimental and analytical methods and include schwertmannite, bilinite, Fe 2+ -(hydrous) sulfates (e.g., melanterite), szomolnokite, other hydrous Fe sulfates (halotrichite and ferricopiapite), and/or jarosite group minerals [Bell, 1996; Bell et al., 1994; Bishop and Murad, 1996; Elwood Madden et al., 2004; Gendrin et al., 2005; Klingelhöfer et al., 2004; Lane et al., 2004; Marion et al., 2003; Morris et al., 1996, 2004; Tosca et al., 2004, 2005]. [17] Fe-bearing carbonates are found in Martian meteorites [Bridges et al., 2001] and Fe chlorides have been modeled [Burns, 1993; Marion et al., 2003]. We first explore the crystallization of Fe (hydr)oxides and Fe sulfates and return to the Fe chloride and Fe carbonate phases later because there is no clear evidence that the latter two phases are abundant on Mars surface. 3. Models of Fe-S-Bearing Solutions, Hematite, and Fe 2+ Sulfates on Mars [18] To model the possible reactions in the Fe-O-H-S system, we use thermodynamic data at standard temperature and pressure (T = 278 K, P = 1 bar, Table 2) because lower temperature data are not available in many cases. The lower temperatures and lower fo 2 values on Mars will cause the phase boundaries move slightly [Burns, 1993; Gooding, 1978], although the general phase topology is likely to be correct, on the basis of comparisons with the FREZCHEM model [Marion et al., 2003]. [19] Because liquid H 2 O is only transient on the surface of Mars, to model Fe-S-bearing solutions we first consider the effects of oxidation (gas-solid or solution-solid reactions) and then consider the effects of adding water (solution-solid reactions). However, we do not intend to imply that alteration necessarily followed that sequence. Instead, Martian solutions were likely produced episodically or perhaps even produced on the molecular scale [Bishop et al., 2002]. [20] As stated above, primary sulfides and magmatic S- gases and secondary sulfates provide good sources for the S on Mars and may contribute to a low-temperature solution (H 2 S aq, HS aq, SO 2 4 aq, or HSO 4 aq ) or high-temperature acid fog (SO 2 or H 2 S vapor). If sulfide is the main S source, we can model the initial breakdown as a reaction driven by oxidation, that is, transfer of electrons following Fe 2+ =Fe 3+ +e and S 2 =S 4+ +6e =S 6+ +8e. The Fe 2+ and Fe 3+ may be in solution or minerals. The S 2 is initially in sulfide, H 2 S aq,h 2 S g,orhs aq. The S 4+ could be in SO 2g and S 6+ may be in the form of SO 2 4aq, HSO 4aq or sulfates. Commonly these types of reactions are written as if they are driven by addition of O 2 or photons, or loss of H 2. Here, we examine the reactions mediated by O 2 as a tool to examine oxygen fugacity (fo 2 ) in solutions on Mars as recorded by the secondary minerals, although we emphasize that we are not ruling out photodissociation of H 2 OorCO 2, H 2 -loss reactions, or reactions with hydrogen peroxide [see Schaefer, 1996]. [21] We can bracket the possible fo 2 conditions on the Martian surface using the fo 2 of the near-surface rocks (minimum surface value) and the fo 2 of the atmosphere (maximum surface fo 2 is 10 5 bars). The near-surface Martian meteorites formed at a maximum of 1.5 log units Table 1. Minerals in the Fe-O-S-H System That May Exist on the Surface of Mars Abbreviation Name Formula Sulfides cpy chalcopyrite CuFeS 2 mar marcasite FeS 2 pent pentlandite (Fe, Ni) 9 S 8 po pyrrhotite Fe 1-x S py pyrite FeS 2 tro troilite FeS Fe (Hydr)oxides ak akaganéite b-feo(oh) fero feroxyhyte d 0 -FeO(OH) fer ferrihydrite Fe(OH) 3 goe goethite a-feo(oh) hem hematite Fe 2 O 3 lep lepidocrocite g-feo(oh) mag magnetite Fe 2+,Fe 3+ 2 O 4 mhem maghemite Fe 2 O 3 CuFeS 2 Sulfate Salts amar amarantite Fe 3+ (SO 4 )(OH).3H 2 O bil bilinite Fe 2+ Fe 3+ 2 (SO 4 ) 4.22H 2 O bot botryogen MgFe((SO 4 ) 2 (OH).7H 2 O but butlerite Fe 3+ (SO 4 )(OH).2H 2 O cop copiapite Fe 2+ Fe 3+ 4 (SO 4 ) 6 (OH) 2.20H 2 O coq coquimbite Fe 3+ 2 (SO 4 ) 3.9H 2 O fcop ferrocopiapite Fe 3+ 5 O(SO 4 ) 6 (OH).20H 2 O fhex ferrohexahydrite Fe 2+ SO 4.6H 2 O fib fibroferrite Fe 3+ (SO 4 )(OH).5H 2 O hal halotrichite FeAl 2 (SO 4 )4.22H 2 O h-jar hydronium jarosite H 3 OFe 3+ 3 (SO 4 ) 2 (OH) 6 hoh hohmannite Fe 3+ 2 (SO 4 ) 2 (OH) 2.7H 2 O jar jarosite KFe 3+ 3 (SO 4 ) 2 (OH) 6 kor kornelite Fe 3+ 2 (SO 4 ) 3.7H 2 O lau lausenite Fe 3+ 2 (SO 4 ) 3.6H 2 O mel melanterite Fe 2+ (SO 4 ).7H 2 O mer mercallite K(HSO 4 ) metv metavoltine (K,Na) 5 Fe 3+ 3 (SO 4 ) 6 (OH) 2.8H 2 O mhoh metahohmannite Fe 3+ 2 (SO 4 ) 2 (OH) 2.3H 2 O n-jar natrojarosite NaFe 3+ 3 (SO 4 ) 2 (OH) 6 que quenstedtite Fe 3+ 2 (SO 4 ) 3.10H 2 O rhom rhomboclase H 3 OFe 3+ (SO 4 ) 2.3H 2 O rom römerite Fe 2+ Fe 3+ 2 (SO 4 ) 4.14H 2 O roz rozenite Fe 2+ (SO 4 ).4H 2 O sh1/8 schwertmannite1/8 Fe 8 O 8 (SO 4 )(OH) 6 sh3/16 schwertmannite3/16 Fe 8 O 8 (SO 4 ) 1.5 (OH) 5 sidn sideronatrite Na 2 Fe 3+ 3 (SO 4 ) 2 (OH).3H 2 O sidt siderotil Fe 2+ (SO 4 ).5H 2 O szo szomolnokite Fe 2+ (SO 4 ).H 2 O vol voltaite K 2 Fe 2+ 5 Fe 4 (SO 4 ) 12 18H 2 O above the quartz-fayalite-magnetite buffer (QFM +1.5, Martian meteorite MIL03346 [Dyar et al., 2005b]). That value is equivalent to an fo 2 of bars (using equations from O Neill [1987]), or 9 log units below the hematite-magnetite buffer (HM-9). At those conditions, sulfides such as troilite (FeS) from meteoritic material and pyrrhotite (Fe 1-x S) from the Martian igneous rocks may be present. Thermodynamic data for troilite are available [Robie et al., 1978], but data for pyrrhotite vary therefore we use troilite as a proxy for pyrrhotite (following Jerz and Rimstidt [2003]). At such low fo 2 (10 78 bars), any Fe that is not hosted by sulfides will be found in silicates (e.g., olivine, pyroxene, etc.), Fe 2+ -rich oxides (e.g., magnetite, chromite, ilmenite), aqueous Fe 2+ species (Fe 2+ aq ), or as surface and nanophase species on preexisting Fe 2+ oxides and sulfides [e.g., Nesbitt et al., 1998]. 4of15

5 Table 2. Reactions and Equilibrium Constants Used to Construct Figures 1, 4a, and 4b at 25 C and 10 5 Pa Phase Reaction logk Figure Ref. a A Tro-Mel FeS + 2O 2 +7H 2 O = FeSO 4.7H 2 O JR B Tro-FeSO 4 FeS + 2O 2 = FeSO 4aq JR C Tro-Py FeS + H 2 SO 4 = FeS = 2 O 2 +H 2 O JR D Mel-FeSO 4 FeSO 4.7H 2 O = FeSO 4aq +7H 2 O JR E FeS 2 -Mel FeS = 2 O 2 +8H 2 O = FeSO 4.7H 2 O+H 2 SO JR F Py-FeSO 4 FeS = 2 O 2 +H 2 O = FeSO 4aq +H 2 SO JR G Fer-Fe 2+ 4Fe(OH) 3s +8H aq = 4Fe aq +O H 2 O l NM H Jar-Fer KFe 3 (SO 4 ) 2 (OH) 6s +3H 2 O l = 3Fe(OH) 3s +K aq + 2SO + 4aq +3H aq 18 4 NM I Jar-Fe 2+ KFe 3 (SO 4 ) 2 (OH) 6s +3e +6H aq = 3Fe aq + 2SO + 4aq +6H 2 O l +6K aq 30 4 NM J Fer-Fe 3+ 4Fe(OH) 3s +3H aq =Fe aq +3H 2 O l M K Sh-Fe 3+ FeO(SO 4 ) q (OH) 1-2q s + (3-2q)H aq =Fe aq + qso 4aq + (2-2q)H 2 Ol M L Sh-Fer FeO(SO 4 ) q (OH) 1-2q sh + (1+2q)H 2 O = Fe(OH) 3fer + qso qH + b 4 M M Sh1/8-Fer Fe 8 O 8 (SO 4 )(OH) 6s + 10H 2 O l = 8Fe(OH) 3s +SO + 4aq +2H aq B N Sh3/16-Fer Fe 8 O 8 (SO 4 ) 1.5 (OH) 5s H 2 O l = 8Fe(OH) 3s +SO + 4aq +2H aq B O Water 2H 2 O l =O 2g +4H aq +4e NM a NM, Nordstrom and Munoz [1986]; JR, Jerz and Rimstidt [2003]; M, Majzlan et al. [2004]; B, Bigham et al. [1996] equations recast as a function of Fe(OH) 3. logk values for the reactions were calculated by rearranging the equations where appropriate. b ph = {(qlog[so 2 4 ]) logk equation K + logk equation J (1+2q)logaH 2 O}/2q. [22] When mafic-ultramafic rocks are exposed to the Martian atmosphere that has a fo 2 of 10 5 bars (HM + 64), oxidation likely occurs. One of the first reactions is the conversion of Fe 2+ to Fe 3+ with transfer of one electron (above). For instance, Fe 2+ in sulfides may produce magnetite (Fe(II), Fe(III) 2 O 4 ) and/or Fe 2+ -bearing surface species on the sulfide (section 2.2). At fo bars, magnetite converts to hematite at the hematite-magnetite (HM) buffer (Figure 1). Alternately, if the reaction occurs in an acidic solution (ph < 2), the reaction is ðr2þ Fe 2þ mineral þ1 = 4 O 2 þ H þ aq ¼ Fe3þ aq þ1 = 2 H 2 O: 3+ As written, reaction (2) involves minor hydrolysis of Fe aq and the reaction consumes H + aq [Nordstrom et al., 2000], similar to reaction K in Table 2. But, at more basic conditions (ph > 2), the Fe 3+ aq forms Fe (hydr)oxides and acid is produced, for example, [2003]). At low water concentrations (e.g., relative humidity 50%, log ah 2 O 0.3), if fo 2 is increased troilite reacts to form Fe 2+ sulfates such as melanterite (Fe 2+ (SO 4 ).7H 2 O) for a range of ah 2 SO 4 conditions (Figure 1; reaction A in Table 2). At high water concentrations (e.g., relative humidity 100%, log ah 2 O 0.0), troilite converts to aqueous Fe species if fo 2 is increased over a range of ah 2 SO 4 conditions (Figure 1; reaction B in Table 2). If ah 2 SO 4 is increased, then troilite reacts to form FeS 2 minerals such as pyrite or marcasite (Figure 1; reaction C in Table 2). We will refer to pyrite subsequently, but marcasite has similar thermodynamic properties and is therefore implied [Robie et al., 1978]. In the presence of water and O 2, pyrite forms H 2 SO 4 aq and aqueous Fe species or Fe phases (e.g., reactions E and F in Table 2), ðr3þ Fe 3þ aq þ 3H 2O ¼ FeðOHÞ 3 þ 3H þ aq ; which is the reverse of reaction J in Table 2. These variations in ph as a function of the stability of Fe 3+ aq versus Fe (hydr)oxides make it complicated to model the Fe-O-S system; therefore, in this section we use H 2 SO 4 aq as the major S-bearing aqueous species (following Jerz and Rimstidt [2004]). The other advantage to this approach is that H 2 SO 4aq provides a measure of the bulk acidity of the solution because it represents the quantitative capacity of the solution to react with a strong base. However, we note that H 2 SO 4 aq dissociates into a range of aqueous species (H + aq,h 2 S aq,hs aq,so 2 4aq and HSO 4aq ), and in particular affects both aso 2 4 and ph via ðr4þ H 2 SO 4aq ¼ HSO 4aq þ Hþ aq ¼ SO2 4aq þ 2Hþ aq : Later we examine the Fe-O-S-H system as a function of SO 4 2 and ph for convenience. [23] At oxygen fugacity (fo 2 ) conditions above the HM buffer, troilite may be altered to different phases depending on the activity of water (ah 2 O), activity of sulfuric acid (ah 2 SO 4 ), and fo 2 (discussed in detail by Jerz and Rimstidt Figure 1. Diagram showing the stability fields for troilite (pyrrhotite), melanterite, FeS 2 (pyrite or marcasite), and Fe species in solution as a function of wetting or relative humidity, logfo 2 (log units above the hematite-magnetite buffer) and log[h 2 SO 4 ] (see reaction (4)). This diagram was constructed using data listed in Table 2 from Jerz and Rimstidt [2003]. 5of15

6 Figure 2. (a) Diagram showing the stoichiometries of the Fe sulfate and Fe (hydr)oxide minerals (modified after Jerz and Rimstidt [2003]). The mineral abbreviations and formulae are given in Table 1. The general trend observed in acid mine drainage situations is shown as path a, and this may be a possible pathway for mineral formation on Mars. An alternate path b that commences with decreasing H 2 O/ (3Fe Fe 2+ ) is also shown. (b) Fe 2 O 3 -H 2 O-SO 3 diagram (modified after Merwin and Posnjak [1937]) showing the stability of different Fe 3+ sulfates, goethite, and solution. The mineral abbreviations and formulae are given in Table 1. The trajectory c shows hypothetical hydration of coquimbite. The trajectory d shows the hypothetical dehydration of a solution with slightly lower Fe 2 O 3 /SO 3 to form rhomboclase. where the oxidation conditions control the Fe 2+ /Fe 3+.We note that recent work indicates that there may in fact be a continuum in the surface properties of troilite, pyrrhotite and pyrite [Thomas et al., 2003], but here we have considered these minerals as pure thermodynamic end-members. 4. Models of Fe Sulfates and Fe (Hydr)oxides 4.1. Variations in Bulk Mineralogy [24] If troilite or pyrrhotite is altered at low to moderate ah 2 SO 4 via leaching or acid fog processes then melanterite may form (reaction A in Table 2; Figure 1) [also Tosca et al., 2005]. Once melanterite has formed, then on Earth the solution will precipitate Fe sulfate and Fe oxide minerals in a sequence close to path a in Figure 2a. In other words, the bulk mineral assemblage in equilibrium with the solution tends to follow a sequence of decreasing 3Fe 3+ / (3Fe Fe 2+ ) (oxidation); decreasing H 2 O/(3Fe Fe 2+ ); and increasing OH/(OH + 2SO 4 )[Jerz and Rimstidt, 2003]. Commonly, melanterite crystallizes first, followed by more oxidized minerals in the römerite group (römeritebilinite), copiapite-group, followed by minerals with decreasing H 2 O/(3Fe Fe 2+ ) and increasing OH/(OH + 2SO 4 ): fibroferrite-group (fibroferrite, amarantite, butlerite) and hohmannite-group (hohmannite, metahohmannite), followed by jarosite-group minerals, schwertmannite and finally Fe (hydr)oxides such as goethite or ferrihydrite (Figure 2a; see Table 1 for mineral abbreviations and formulae). Similar sequences of minerals are observed in acid mine drainage sites [Jerz and Rimstidt, 2003; Hammarstrom et al., 2005] and in addition OH-free sulfates such as coquimbite and volaite are common [Jambor et al., 2000]. [25] It is possible that in the low-ph 2 O atmosphere of Mars the Fe sulfates precipitate from solution in a sequence where the first bulk mineral assemblage has less H 2 O/ (3Fe Fe 2+ ) and is followed by mineral assemblages that are progressively more oxidized and richer in OH/(OH + 2SO 4 ) (path b in Figure 2a). In this case, any initial melanterite would be replaced by mineral assemblages that are progressively depleted in H 2 O/(3Fe Fe 2+ ) such as ferrohexahydrite, sidterotil, rozenite and szomolnokite. Szomolnokite could then be followed by other more oxidized sulfate assemblages such as copiapite (if there is a slight hydration), metahohmannite, jarosite-group minerals (if H 3 O +,K +,orna + are available), and schwertmannite (Figure 2a). Eventually Fe (hydr)oxide phases (e.g., hematite or magnetite) precipitate from solution, bringing to an end the trend within the bulk mineral assemblage of oxidation, decreasing H 2 O/(3Fe Fe 2+ ) and increasing OH/(OH + 2SO 4 ) (path b in Figure 2a). If the system was closed during the precipitation sequence from melanterite to Fe (hydr)oxides, then the solution would have a higher H 2 O/(3Fe 3+ +Fe 2+ ), lower 3Fe 3+ /(3Fe 3+ +Fe 2+ ), and lower OH/(OH + 2SO 4 ). [26] While it would be preferable to model paths a and b in their entirety, this is not possible because the necessary thermodynamic data do not exist. Tosca et al [2005] provided elegant models of Fe-sulfate formation on Mars, but similarly their models were limited by the available thermodynamic data. Furthermore, laboratory experiments on Fe sulfates are difficult to interpret because Fe-S-O phases are difficult to synthesize and analyze therefore reaction stoichiometry may not be well constrained between data sets [e.g., Majzlan et al., 2004]. Also, paragenetic sequences for sulfate minerals with high 3Fe 3+ /(3Fe 3+ + Fe 2+ ) are difficult to define due to the influence of kinetic factors and local variations in chemistry in nature (e.g., variable Mg/Fe, Fe 2 O 3 /SO 3,H 2 O or ph) [e.g., Jambor et al., 2000; Jamieson et al., 2005]. For example, Mg-bearing copiapite is found in the vicinity of hydronium-jarosite at Iron Mountain, California, USA, although those phases are not predicted thermodynamically [Jamieson et al., 2005]. In this section we only considered mineral assemblages that form from a solution, but we note that oxidation and variations in H 2 O/(3Fe Fe 2+ ) also may occur between the solid sulfate and a gas (e.g., in the atmosphere) Oxidation and ph Effects [27] If troilite or pyrrhotite were altered at high ah 2 SO 4 via leaching or acid fog processes, then FeS 2 (pyrite or marcasite) will form (Figure 1). For example, marcasite has been observed as alteration rims on pyrrhotite in some Martian meteorites indicating that the reaction occurs in Martian rocks. To our knowledge, studies have not been attempted to determine if the marcasite formed on Mars or on Earth, but it is less soluble than salts found in Martian meteorites that have also been identified on Mars (e.g., gypsum, Mg sulfates; above). 6of15

7 [28] If FeS 2 is further oxidized then sulfuric acid is produced along with melanterite (reaction E in Table 2) or aqueous Fe species (e.g., FeSO 4 aq; reaction F in Table 2), or Fe (hydr)oxides (reactions (2) and (3)). Some pyrite oxidation reactions are exothermic and terrestrial solutions have been measured at 47 C [Nordstrom et al., 2000] although the latter situation may be mediated by biologic processes and high PO 2 on Earth. Nonetheless, on both Earth and Mars, the generation of sulfuric acid and heat may accelerate dissolution reactions and drive evaporation of any coexisting solutions. [29] The phase relations for pyrite stability have been examined for Mars by many authors using Eh-pH diagrams and we summarize their results in Figures 3a 3c (see caption for how to convert Eh to PO 2 ). At near-neutral to basic conditions and low-moderate Eh, FeS 2 (pyrite) reacts to form HS, as Eh increases SO 2 4aq then a Fe (hydr)oxide form (Figures 3a and 3b; similar to reactions (2) and (3) above). The Fe (hydr)oxide may be ferrihydrite (Figure 3b) or goethite dependent on the bulk composition, coexisting phases, and the thermodynamic data used [e.g., Bigham et al., 1996; Burns, 1988] and magnetite, although proposed (Figure 3b), is unlikely (above). [30] If conditions are neutral-acidic, then at low Eh pyrite reacts to form H 2 S aq, at slightly higher Eh SO 2 4 forms (Figure 3c), and at slightly higher Eh Fe 2+ aq or other Fe species (e.g., FeSO 4, FeHCO 3 ) form. As Eh is raised further then either a Fe (hydr)oxide or a hydrous Fe sulfate like schwertmannite (Figures 3b and 3c) or melanterite (reaction E in Table 2) may form. At extremely acidic conditions, pyrite reacts at low Eh to form H 2 S aq and at 3+ higher Eh to form HSO 4 aq, then jarosite (KFe 3 (SO 4 ) 2 (OH) 6 ) or other Fe 3+ phases (Figures 3a 3c, dependent on the availability of other cations). [31] The Eh versus ph diagrams (Figures 3a and 3b) illustrate very clearly the stability fields for Fe sulfides and their weathering products (similar to Figure 1). They are also useful in showing that FeS 2 is stable in Martian igneous rocks (QFM + 1.5; Figure 3c) and at the HM buffer, but that any FeS 2 is unstable when exposed to the Martian atmosphere (Figure 3c). Yet, these diagrams have limited use because they only indicate the effect of varying PO 2 for ah 2 O = 1. Also, the Fe sulfate (jarosite or schwertmannite) ferrihydrite and Fe 2+ aq ferrihydrite boundaries are poorly defined (Figure 3c). For example, the Fe sulfate Fe (hydr)oxide boundary varies from ph 2 toph7. [32] This variability is due to effects of bulk composition, in particular aso 2 4 and afe total. Because we have poor constraints on the ionic strength of the solutions and water:- rock ratio, we make the simplifying assumption that the concentrations measured in the Martian soils are equivalent to activities (aso 2 4 =[SO 2 4 ] and afe total =[Fe total ]). We note that this assumption is not strictly correct in a thermodynamic sense because activity coefficients are probably less than one resulting in activities that are lower than concentrations. Nonetheless, this approach places maximum constraints on the behavior of the Martian solutions and has previously been used to model secondary minerals on the Martian surface [e.g., Catling, 1999; Marion et al., 2003]. We also assume that [SO 2 4 ] is a useful measure of the S in solution, despite the fact that there are the other possible species (e.g., H 2 S, HS, and HSO 2 4 ). We now examine the effects of SO 2 4,Fe total,ah 2 O and ph on Fe sulfate, Fe (hydr)oxide and Fe 2+ aq stability, because the boundaries between those phases appear to be most problematic. This approach also allows us to consider changes that occur along the trend of increasing OH/(OH + SO 4 ) (Figure 2a) SO 2 4,Fe total, and ah 2 O Effects [33] To examine phase stabilities in the (K)-Fe-O-S system, we have constructed logso 2 4 versus ph diagrams (Figures 4a 4c). The diagrams were constructed using the reactions and equilibrium constants given in Table 2 and assuming activity of potassium, ak + =[K + ] = 0.01 and [Fe total ] = 0.3 based on the concentration of those elements in the Mars Pathfinder and MER soil analyses [Foley et al., 2003; Gellert et al., 2004; Rieder et al., 2004]. We assume that that the jarosite is K-bearing, but this is not critical to the general model and other forms of jarosite may be present (H 3 O +,Na +, etc.). In Figure 4a, we show the phase boundaries as a function of log fo 2 and also for [Fe total ]= 0.01, because as Fe-bearing phases are removed from 2+ in solution, the Fe total decreases. We only consider Fe aq this diagram because our primary interest is the behavior of that phase boundary, but we note that other species Figure 3. Comparison of different Eh versus ph models for Fe phases on Mars and in acid mine drainage environments. Note that the phase boundaries depend on bulk composition, mineral species, and the thermodynamic data chosen. (a) Phase diagram from Burns [1988] (bold minerals and species, solid and dot-dash line, ak + =10 4, afe total =10 2, aso 2 4 =10 2 ) and Fairén et al. [2004] (underlined minerals and species, dotted line, ak + =10 4.2, afe total =10 4.3, aso 2 4 = ). (b) Phase diagram for acid mine drainage from Majzlan et al. [2004] (bold minerals and species, solid line, afe total =10 4, aso 2 4 = ) showing stabilities in the Fe-O-S-H system; note that jarosite would plot in the Fe aq field if K + were present. Phase diagram from Bigham et al. [1996] (normal text for minerals and species, dotted line, ak + = 10 4, afe 2+ =10 3.5, afe 3+ = to 2.3, aso 2 4 =10 2 ). Phase diagram from Elwood Madden et al. [2004] (underlined minerals and species, dashed line, ana + =10 3, aso 2 4 =10 2, afe total not given). Numbered dots indicate different water:rock ratios. If the water:rock was initially high (e.g., 100), then jarosite is stable; if the reaction path moves to higher ph as more basalt is weathered or H 2 O is removed (e.g., water:rock 10 to 0.1), then the solution moves out of the jarosite stability field and into the Fe 3+ mineral stability field [Elwood Madden et al., 2004]. (c) Summary diagram showing the range of the boundaries for Fe sulfates and Fe species in the literature along with fo 2 conditions relevant on the Martian surface (QFM + 1.5, HM, and the Martian atmosphere). The jarosite, schwertmannite, and Fe 2+ aq fields show a stability range due to different choices of the bulk composition in the literature. The H-S-O species in solution are also shown [from King et al. 2004]. Consistent with previous workers, we show these diagrams as Eh-pH diagrams. However, we note that these diagrams may be converted to PO 2 -ph diagrams using reaction O and the equilibrium constant given in Table 2, if ah 2 Ois known or assumed. Alternately, the reader may examine the phases relative to the fo 2 buffers illustrated. 7of15

8 may be important in detail (e.g., FeSO 4aq,Fe 3+ aq, melanterite, goethite, etc.). [34] The boundaries of the Fe 2+ aq field are dependent on logfo 2 and the total Fe concentration (Figure 4a). The Fe 2+ aq schwertmannite boundary is poorly constrained because goethite may be an intermediate phase [Bigham et al., 1996], and so we omitted that boundary for clarity, but it should behave in a similar manner to the Fe 2+ aq jarosite boundary. At relatively reducing conditions (logfo 2 = 65), consistent with the stability limit of troilite (Figure 1), the Fe 2+ aq field is relatively large. The Fe 2+ aq field is smaller for higher Fe total because the solid phases are stabilized. At high logfo 2 ( 10; HM + 59), the Fe 2+ aq field disappears and Fe sulfate minerals and ferrihydrite coexist. In this manner, the precipitation of Fe sulfates and ferrihydrite is strongly enhanced by oxidation. [35] The stability of the jarosite ferrihydrite and schwertmannite ferrihydrite boundary both depend on log[so 4 2 ] and ph (Figure 4b). Schwertmannite has a range of compositions defined by: FeO(SO 4 ) q (OH) 1-2q and here we use Sh1/8 for FeO(SO 4 ) 1/8 (OH) 3/4 and Sh3/16 for FeO(SO 4 ) 3/16 (OH) 5/8. The conditions for the schwertmannite ferrihydrite boundary vary greatly as a function of composition and the thermodynamic data used. For example, schwertmannite is stable to higher ph using thermodynamic data from Majzlan et al. [2004] relative to data from Bigham et al. [1996]. Both sets of data predict schwertmannite stability that is consistent with its occur- Figure 3 8of15

9 (point a in Figure 4b). Note that this value differs slightly from previous authors (Figures 3a 3c) because they used slightly different bulk compositions. Schwertmannite is stable at the same or higher ph, with Sh3/16 stable at ph < 4.8 or 7.5 (points a and b in Figure 4b); and, Sh1/8 stable at ph < 7.8 and 9.7 (points c and d in Figure 4b). The lower ph boundary for schwertmannite is poorly defined, but we estimate that it should be stable above jarosite stability at ph greater than 4 [e.g., Bigham et al., 1996]. The log[so 4 2 ] of the Martian soils is likely a lower limit for log[so 4 2 ] because the soils include nondissolved components. Therefore the maximum stability of the Fe sulfates is likely at higher ph. [37] If ah 2 O is decreased, then the jarosite phase relations shift to slightly higher ph, but the effect is not large (Figure 4c). In contrast, if ah 2 O is decreased the schwertmannite ferrihydrite reaction moves up to about one ph unit higher; stabilizing schwertmannite to higher ph (Figure 4b). Schwertmannite shows this behavior relative to jarosite because it contains a higher mole fraction of H. Figure 4. Graphs of log[so 2 4 ] versus ph for jarosite, ferrihydrite, and Fe 2+ aq based on reactions in Table 2. To construct the diagrams, we assumed ak + = 0.01 and a maximum Fe total of 0.3 based on data from Pathfinder and MER soils. (a) Diagram showing the effect of varying logfo 2 and afe total at fixed ah 2 O = 1. (b) Diagram showing the stability fields of jarosite ferrihydrite and schwertmannite ferrihydrite showing the effect of using data from different authors. Sh1/8 refers to FeO(SO 4 ) 1/8 (OH) 3/4 and Sh3/16 to FeO(SO 4 ) 3/16 (OH) 5/8. The S content of the Pathfinder and MER soils, [SO 2 4 ] = 0.08, is plotted on the diagram to give a constraint on the maximum ph for the different Fe sulfates as described in the text. (c) Diagram showing the effect of varying ah 2 O for the different reactions. Note that the schwertmannite stability curves are shifted to higher ph with decreasing ah 2 O. rence in acidic mining lakes on Earth where log[so 4 2 ]= 2 and ph < 7, [Regenspurg et al., 2004]. [36] If we examine the maximum stability of the Fe sulfates using the composition of the Martian soil (log[so 4 2 ]= 1), then jarosite is stable at ph < Models for Minerals Coexisting With Fe Sulfates and Fe (Hydr)oxides 5.1. Secondary Si-Rich Phases [38] A Si-rich phase occurs on Mars, but its specific identity is unknown with candidates including andesitic glass [Bandfield et al., 2000], clay minerals [McLennan, 2003; Wyatt and McSween, 2002], zeolites [Ruff, 2004], and silica-rich coatings [Kraft et al., 2003; Michalski et al., 2005]. The occurrence of this secondary phase on the Martian surface is consistent with release of SiO 2 into solution during weathering and subsequent precipitation. Silica is an extremely mobile anion during weathering [McLennan, 2003], and tends to precipitate when the solution becomes oversaturated in silica at low temperature or at ph < 9 [Krauskopf and Bird, 1995]. The final form of the secondary Si-rich phase could be incorporated in coatings, cements, overgrowths, veins and/or dust [Kraft et al., 2003; Michalski et al., 2005] and aqueous or hydrothermal processes may convert the material locally to poorly crystalline clays [McLennan, 2003; Wyatt and McSween, 2002] or zeolites [Ruff, 2004]. Such Si-rich coatings have been documented on Mars [Kraft et al., 2003; Michalski et al., 2005; McLennan, 2003] and will effectively slow alteration of the primary mafic minerals (e.g., olivine and pyroxene) Other Salts [39] In addition to the Fe-rich minerals in the Martian soils, more Mg- and Ca-rich secondary minerals may be present also. For example, phosphates, sulfates and halides will precipitate once they become saturated in solution, dependent on the bulk composition and extent of solution evolution as discussed by King et al. [2004]. Fe-rich phosphates are most likely in acidic conditions, whereas Ca-bearing phosphates are more likely in near-neutral conditions [McBride, 1994]. Fe chlorides are likely only a minor component (see Marion et al. [2003] for a more detailed discussion). [40] Carbonates may occur on the surface of Mars [Bandfield et al., 2003] and Fe-Mg-Ca carbonate phases 9of15

10 have been identified in the Martian meteorites [Bridges et al., 2001] and are thermodynamically stable on the Martian surface [Gooding, 1978]. However, the evidence for such phases on Mars surface is limited [Lane et al., 2004; Stockstill et al., 2005]. Carbonate may persist only for a short time on the surface, if at all, dependent on (1) salt crystallization sequences that may result in carbonates crystallizing at depth [King et al., 2004; Warren, 1998]; (2) whether or not carbonates undergo photodecomposition [Mukhin et al., 1996]; and (3) acidic solutions that may cause the carbonate to dissolve or degas [Burns and Fisher, 1993; Fairén et al., 2004; King et al., 2004]. 6. Discussion 6.1. Changes in Fe Sulfate Mineralogy Related to Episodic Events [41] Episodic remobilization and redeposition of secondary phases on Mars is inferred on the basis of large-scale geomorphology [e.g., Jakosky and Phillips, 2001] and small-scale veins and mineral-shaped molds [e.g., Arvidson et al., 2005; Squyres et al., 2004a]. The secondary phases remobilized may include hydrated salts that are metastable on the Martian surface [e.g., Gooding, 1978] and have been subjected to small deviations in water content, atmospheric pressure, temperature and/or radiative bombardment. Except for the latter case, these changes may be driven by volcanic/hydrothermal events, seasonal sublimation or atmospheric events, percolation of solution films along mineral surfaces by capillary action, or hydrologic processes including groundwater transport. A consequence of episodic aqueous events is that previously precipitated minerals may be dissolved/hydrolyzed and reprecipitate elsewhere, perhaps as different minerals in a manner similar to acid mine drainage sites [Bigham and Nordstrom, 2000]. [42] For example, if a copiapite-group mineral is exposed to water with a low OH/(OH + 2SO 4 ), it may dissolve and reprecipitate as bilinite. Such a pathway is the opposite direction to path a described in section 4.1, but late bilinite is observed in acid mine drainage settings [Jambor et al., 2000] and has been modeled for Mars [Tosca et al., 2005]. Similar pathways could be constructed for other minerals and waters. Alternately, slight changes in the bulk composition of the infiltrating solution may cause the mineralogy to change, in particular the Fe-sulfate minerals are very sensitive to the Fe 2 O 3 /SO 3. For instance, if coquimbite is exposed to water and forms a solution, the bulk composition will migrate along path c in Figure 2b. If the water is then lost via evaporation or freezing, the bulk composition should follow path c in reverse. Instead, if the system is no longer closed and the bulk composition changes, the reverse path will not be followed. Such a scenario could occur if a solution with a different precipitation and dissolution history enters the system, perhaps during an episodic event. For example, if Fe 2 O 3 /SO 3 decreases slightly, the return path will not follow path c and may instead follow path d where rhomboclase precipitates (Figure 2b). Precipitation of rhomboclase consumes H + aq, in effect storing H + in solid form: ðr5þ Fe 3þ aq þ 2SO2 4aq þ Hþ aq þ 4H 2O l ¼ H 3 OFe 3þ ðso 4 Þ 2 :3H 2 O: Conversely, if the rhomboclase dissolves then H + is released to solution. [43] Whereas rhomboclase precipitation increases the ph of a solution, other minerals tend to decrease ph dependent on their solubility and stability (Figure 2b). Solutions are acidified when minerals such as copiapite-group minerals precipitate; for example, ferrocopiapite precipitation releases H + aq : ðr6þ 4Fe 3þ aq þ Fe2þ aq þ 6SO2 4aq þ 22H 2O ¼ Fe 3þ 5OSO ð 4 Þ 6 ðohþ:20h 2 O þ 2H þ aq : Some minerals such as jarosite may either consume or release H + dependent on the conditions (compare reactions H and I in Table 2). [44] In summary, if minerals dissolve/hydrolyze and precipitate during multiple episodes, the solution composition will change locally. Mineral dissolution/hydrolysis and precipitation could involve H + -consuming (e.g., reactions (1), (2), (5), or G, I, K in Table 2) or H + -producing reactions (e.g., reactions (3), (6) or H, L, M, N in Table 2) that may affect the mineralogy. Similarly, variations in OH/ (OH + 2SO 4 )orfe 2 O 3 /SO 3 may change the final mineral assemblage Comparison of the Low ph Acid Fog and Leaching Models [45] As indicated above and in Table 3, the major distinction between the low ph acid fog and leaching models is that the former requires a low ph and Fe and Al remain in solution. Acidic solutions tend to result in higher weathering rates, producing Fe-Mg-(Ca)-SO 4 -(Cl) solutions. The minerals derived from the acid fog process include Fe sulfates including those with low OH/(OH + SO 4 ), Fe (hydr)oxides, Fe phosphates (section 5.2), some Ca-Mg-Na sulfates, and possible Ca-, Mg-, or Na + K-(Ca) halides [Tosca et al., 2004; King et al., 2004]. [46] In the leaching model, the solution composition may trend toward near-neutral ph resulting in Fe (hydr)oxide precipitation (Figures 4a and 4b), and causing the initial solution to become Fe-poor and Mg-Na-(Ca)-rich [King et al., 2004]. The minerals derived from the leaching process may include Fe sulfates with moderate-high OH/(OH + SO 4 ), Fe (hydr)oxides, Ca-Mg-Na sulfates, Ca phosphates, and possible Mg-Na + K halides and Fe-Mg-Ca carbonates [King et al., 2004]. [47] To summarize, the major divergence between the low ph acid fog model as it is currently applied and the leaching model is whether or not the solution ph is raised. Because primary minerals (e.g., olivine, pyroxene) are identified on the Martian surface, any solutions would have the opportunity to react via reaction (1) if the process occurs in a closed system. The mineral assemblages indicating an acid environment were preserved by processes such as (1) removal of components from the system, for example, in dust storms or in episodic aqueous events that vary water:rock ratios (for details, see Elwood Madden et al. [2004] and Tosca et al. [2005]); (2) mineral surfaces that are inaccessible because they are coated with Si-rich coatings; or (3) disequilibrium processes due to a range of processes including terminating a reaction before it has reached completion or kineticlimited processes rather than thermodynamically limited 10 of 15